Chemistry Education...

Baggie Chemistry

2009-03-28T22:36:00.000-07:00

Experiment with Chemical Reactions

By Anne Marie Helmenstine, Ph.D., About.com
Overview

An ordinary ziploc bag can unlock a world of interest in chemistry and in the reactions within and around us. In this project, safe materials are mixed to change colors and produce bubbles, heat, gas, and odor. Explore endothermic and exothermic chemical reactions and help students develop skills in observation, experimentation, and inference. These activities are targeted for students in grade 3, 4, and 5, although they may also be used for higher grade levels.

Objectives
The purpose is to generate student interest in chemistry. Students will observe, experiment, and learn to draw inferences.

Materials

These quantities are suitable for a group of 30 students to perform each activity 2-3 times:

* 5-6 plastic ziploc-style bags per lab group
* 5-6 clear plastic vials or test tubes (may be used instead of baggies)
* 1 gallon bromothymol blue indicator
* 10-ml graduated cylinders, one per lab group
* teaspoons, 1-2 per lab group
* 3 pounds calcium chloride (CaCl2, from chemical supply house or from store selling this type of 'road salt' or 'laundry aid')
* 1-1/2 pounds sodium bicarbonate (NaHCO3, baking soda)

Activities

Explain to the students that they will be performing chemical reactions, making observations about the results of these reactions, and then designing their own experiments to explain their observations and test hypotheses that they develop. It may be helpful to review the steps of the scientific method.

1. First, direct the students to spend 5-10 minutes exploring the lab materials using all of their senses except taste. Have them write down their observations regarding the way the chemicals look and smell and feel, etc.

2. Have the students explore what happens when the chemicals are mixed in baggies or test tubes. Demonstrate how to level a teaspoon and measure using a graduated cylinder so that students can record how much of a substance is used. For example, a student could mix a teaspoon of sodium bicarbonate with 10 ml of bromothymol blue solution. What happens? How does this compare with the results of mixing a teaspoon of calcium chloride with 10 ml of indicator? What if a teaspoon of each solid and the indicator are mixed? Students should record what they mixed, including quantities, the time involved to see a reaction (warn them that everything will happen very fast!), the color, temperature, odor, or bubbles involved... anything they can record. There should be observations such as:

* Gets hot
* Gets cold
* Turns yellow
* Turns green
* Turns blue
* Produces gas

3. Show students how these observations can be written down to describe rudimentary chemical reactions. For example, calcium chloride + bromothymol blue indicator --> heat. Have the students write out reactions for their mixtures.

4. Next, students can design experiments to test hypotheses they develop. What do they expect to happen when quantities are changed? What would happen if two components are mixed before a third is added? Ask them to use their imagination.

5. Discuss what happened and go over the meanings of the results.

Toxic Elements

2009-02-19T15:33:00.001-08:00

Have you ever wondered which elements are toxic? Everything is toxic if the dose is high enough, so I've compiled a short list of elements that have no nutritional value, even in trace amounts. Some of these elements accumulate in the body, so there is no truly safe exposure limit for those elements (e.g., lead, mercury). Barium and aluminum are examples of elements which can be excreted, at least to a certain extent. Most of these elements are metals. The man-made elements are radioactive and toxic whether they are metals or not.

Aluminum
Antimony
Arsenic (metalloid)
Barium
Beryllium
Cadmium
Hexavalent Chromium Cr⁶⁺ (Cr³⁺ is necessary in trace amounts for proper nutrition)
Lead
Mercury
Osmium
Thallium
Vanadium

radioactive metals
Polonium (metalloid)
Thorium
Radium
Uranium
Transuranium elements (e.g., polonium, americium)
Radioactive isotopes of metals that might not otherwise be highly toxic (e.g., cobalt-60, strontium-90)

From: about.com

Memorize the Elements

2009-02-19T15:33:00.000-08:00

If you take a chemistry class there is an excellent chance you will be required to memorize the names and order of the first few elements of the periodic table. Even if you don't have to memorize the elements for a grade, it is helpful to be able to recall that information rather than look it up every time you need it. Here is a mnemonic you can use to help make the memorization process easier. The symbols for the elements are associated with words that form a phrase. If you can remember the phrase and know the symbols for the elements then you can memorize the order of the elements.

Hi! - H
He - He
Lies - Li
Because - Be
Boys - B
Can - C
Not - N
Operate - O
Fireplaces - F

New - Ne
Nation - Na
Might - Mg
Also - Al
Sign - Si
Peace - P
Security - S
Clause - Cl

A - Ar
King - K
Can - Ca

From: about.com

Liquid Elements

2009-02-19T15:28:00.008-08:00

There are two elements that are liquid at the temperature technically designated 'room temperature' or 298 K (25° C) and a total of six elements that can be liquids at actual room temperatures and pressures.

Liquid at 25°C

Bromine
Mercury

Become Liquid 25°C-40°C

Francium
Cesium
Gallium
Rubidium

Bromine (symbol Br and atomic number 35) and mercury (symbol Hg and atomic number 80) are both liquids at room temperature. Bromine is a reddish-brown liquid, with a melting point of 265.9 K. Mercury is a toxic shiny silvery metal, with a melting point of 234.32 K.

Francium, cesium, gallium, and rubidium are four elements that melt at temperatures slightly higher than room temperature. Francium (symbol Fr and atomic number 87), a radioactive and reactive metal, melts around 300 K. Francium is the most electropositive of all the elements. Cesium (symbol Cs and atomic number 55), a soft metal that violently reacts with water, melts at 301.59 K. The low melting point and softness of francium and cesium are a consequence of the size of their atoms. In fact, cesium atoms are larger than those of any other element. Gallium (symbol Ga and atomic number 31), a grayish metal, melts at 303.3 K. Gallium can be melted by body temperature, as in a gloved hand. Rubidium (symbol Rb and atomic number 37) is soft, silvery-white reactive metal, with a melting point of 312.46 K. Rubidium spontaneously ignites to form rubidium oxide. Like cesium, rubidium reacts violently with water.

From: about.com

Which Elements Are Named After People?

2009-02-19T15:28:00.007-08:00

There are 13 elements named after people, although only 12 of the names are formally accepted by the International Union of Pure and Applied Chemistry (IUPAC).

* bohrium (Bh, 107) – Niels Bohr
* curium (Cm, 96) – Pierre and Marie Curie
* einsteinium (Es, 99) – Albert Einstein
* fermium (Fm, 100) – Enrico Fermi
* gallium (Ga, 31) – both named after Gallia (Latin for France) and its discoverer, Lecoq de Boisbaudran (le coq, the French word for 'rooster' translates to gallus in Latin)
* hahnium (105) – Otto Hahn (Dubnium, named for Dubna in Russia, is the IUPAC-accepted name for element 105)
* lawrencium (Lr, 103) – Ernest Lawrence
* meitnerium (Mt, 109) – Lise Meitner
* mendelevium (Md, 101) – Dmitri Mendeleev
* nobelium (No, 102) – Alfred Nobel
* roentgenium (Rg, 111) – Wilhelm Roentgen (formerly Ununumium)
* rutherfordium (Rf, 104) – Ernest Rutherford
* seaborgium (Sg, 106) – Glenn T. Seaborg

From: about.com

Which Elements Are Named for Places?

2009-02-19T15:28:00.006-08:00

This is an alphabetical list of element toponyms or elements named for places or regions. Ytterby in Sweden has given its name to four elements: Erbium, Terbium, Ytterbium and Yttrium.

* Americium – America, the Americas
* Berkelium – University of California at Berkeley
* Californium – State of California and University of California at Berkeley
* Copper - probably named for Cyprus
* Darmstadtium – Darmstadt, Germany
* Dubnium – Dubna, Russia
* Erbium – Ytterby, a town in Sweden
* Europium – Europe
* Francium – France
* Gallium – Gallia, Latin for France. Also named for Lecoq de Boisbaudran, the element's discoverer (Lecoq in Latin is gallus)
* Germanium – Germany
* Hafnium – Hafnia, Latin for Copenhagen
* Hassium – Hesse, Germany
* Holmium – Holmia, Latin for Stockholm
* Lutetium – Lutecia, ancient name for Paris
* Magnesium – Magnesia prefecture in Thessaly, Greece
* Polonium – Poland
* Rhenium – Rhenus, Latin for Rhine, a German province
* Ruthenium – Ruthenia, Latin for Russia
* Scandium – Scandia, Latin for Scandinavia
* Strontium – Strontian, a town in Scotland
* Terbium – Ytterby, Sweden
* Thulium – Thule, a mythical island in the far north (Scandinavia?)
* Ytterbium – Ytterby, Sweden
* Yttrium – Ytterby, Sweden

From: about.com

Elements in the Human Body

2009-02-19T15:28:00.005-08:00

99% of the mass of the human body is made up of only six elements: oxygen, carbon, hydrogen, nitrogen, calcium, and phosphorus. Every organic molecule contains carbon. Since 65-90% of each body cell consists of water (by weight), it isn't surprising that oxygen and hydrogen are major components of the body.

Here's is a look at the major elements in the body and what these elements do.

Images 1-11 of 11

Enter Gallery

Oxygen	Carbon	Hydrogen	Nitrogen
Calcium	Phosphorus	Potassium	Sodium
Chlorine	Magnesium	Sulfur

From: about.com

What Are the Elements in the Human Body?

2009-02-19T15:28:00.004-08:00

Most of the human body is made up of water, H₂O, with cells consisting of 65-90% water by weight. Therefore, it isn't surprising that most of a human body's mass is oxygen. Carbon, the basic unit for organic molecules, comes in second. 99% of the mass of the human body is made up of just six elements: oxygen, carbon, hydrogen, nitrogen, calcium, and phosphorus.

Oxygen (65%)
Carbon (18%)
Hydrogen (10%)
Nitrogen (3%)
Calcium (1.5%)
Phosphorus (1.0%)
Potassium (0.35%)
Sulfur (0.25%)
Sodium (0.15%)
Magnesium (0.05%)
Copper, Zinc, Selenium, Molybdenum, Fluorine, Chlorine, Iodine, Manganese, Cobalt, Iron (0.70%)
Lithium, Strontium, Aluminum, Silicon, Lead, Vanadium, Arsenic, Bromine (trace amounts)

Reference: H. A. Harper, V. W. Rodwell, P. A. Mayes, Review of Physiological Chemistry, 16th ed., Lange Medical Publications, Los Altos, California 1977.

What Is the Most Abundant Element?

2009-02-19T15:28:00.003-08:00

The most abundant element in the universe is hydrogen, which makes up about 3/4 of all matter! Helium makes up most of the remaining 25%. Oxygen is the third most abundant element in the universe. All of the other elements are relatively rare.

The chemical composition of the earth is quite a bit different from that of the universe. The most abundant element in the earth's crust is oxygen, making up 46.6% of the earth's mass. Silicon is the second most abundant element (27.7%), followed by aluminum (8.1%), iron (5.0%), calcium (3.6%), sodium (2.8%), potassium (2.6%). and magnesium (2.1%). These eight elements account for approximately 98.5% of the total mass of the earth's crust. Of course, the earth's crust is only the outer portion of the earth. Future research will tell us about the composition of the mantle and core.

Reference:
Element Distribution in the Earth's Crust
http://ww2.wpunj.edu/cos/envsci-geo/distrib_resource.htm

How Are Elements Named?

2009-02-19T15:28:00.002-08:00

Do you know which element is Azote, with the symbol Az? Element names are not the same in every country. Many countries have adopted the element names that have been agreed upon by the International Union of Pure and Applied Chemistry (IUPAC). According to the IUPAC, "elements can be named after a mythological concept, a mineral, a place or country, a property, or a scientist".

If you look at the Periodic Table, you'll see some of the higher numbered elements either have no names (only numbers like 118) or else their names are just another way of saying the number (e.g., Ununoctium). The discovery of these elements hasn't been sufficiently documented for the IUPAC to feel a name is justified yet, or else there is a dispute over who gets credit for the discovery (and the honor of selecting an official name).

From: about.com

How Are New Elements Discovered?

2009-02-19T15:28:00.001-08:00

Dmitri Mendeleev is credited with making the first periodic table that resembles the modern periodic table. His table ordered the elements by increasing atomic weight (we use atomic number today). He could see recurring trends, or periodicity, in the properties of the elements. His table could be used to predict the existence and characteristics of elements that hadn't been discovered.

When you look at the modern periodic table, you won't see gaps and spaces in the order of the elements. New elements aren't exactly discovered anymore. However, they can be made, using particle accelerators and nuclear reactions. A new element is made by adding a proton (or more than one) to a pre-existing element. This can be done by smashing protons into atoms or by colliding atoms with each other. The last few elements in the table will have numbers or names, depending on which table you use. All of the new elements are highly radioactive. It's difficult to prove that you have made a new element, because it decays so quickly.

From: about.com

Timeline of Element Discovery

2009-02-19T15:22:00.001-08:00

Timeline of Element Discovery
History of Chemistry

Here's a helpful table chronicling the discovery of the elements. The date is listed for when the element was first isolated. In many cases, the presence of a new element was suspected years or even thousands of years before it could be purified. Click on an element's name to see its entry in the Periodic Table and get facts for the element.

From: about.com

Gases - General Properties of Gases

2009-02-19T15:22:00.000-08:00

Gases - General Properties of Gases
Gas Facts and Equations

By Anne Marie Helmenstine, Ph.D., About.com

All pure substances display similar behavior in the gas phase. At 0° C and 1 atmosphere of pressure, one mole of every gas occupies about 22.4 liters of volume. Molar volumes of solids and liquids, on the other hand, vary greatly from one substance to another. In a gas at 1 atmosphere, the molecules are approximately 10 diameters apart. Unlike liquids or solids, gases occupy their containers uniformly and completely. Because molecules in a gas are far apart, it is easier to compress a gas than it is to compress a liquid. In general, doubling the pressure of a gas reduces its volume to about half of its previous value. Doubling the mass of gas in a closed container doubles its pressure. Increasing the temperature of a gas enclosed in a container increases its pressure.

Because different gases act similarly, it is possible to write a single equation relating volume, pressure, temperature, and quantity of gas. This Ideal Gas Law and the related Boyle's Law, Law of Charles and Gay-Lussac, and Dalton's Law are central to understanding the more complex behavior of real gases.

Ideal Gas Law:
PV = nRT

Boyle's Law:
PV = k₁

Law of Charles and Gay-Lussac:
V = k₂T

Dalton's Law:
P_tot = P_a + P_b

where:
P is pressure, P_tot is total pressure, P_a and P_b are component pressures
V is volume
n is number of moles
T is temperature
k₁ and k₂ are constants

From: about.com

Element Symbols

2009-02-19T15:21:00.001-08:00

Element Symbols
Abbreviations for the Chemical Elements

By Anne Marie Helmenstine, Ph.D., About.com

It's easier to navigate the periodic table and write chemical equations and formulae once you know the symbols for the elements. However, sometimes it's easy to confuse symbols of elements with similar names. Other elements have symbols that don't seem to relate to their names at all! For these elements, the symbol usually refers to an older element name that isn't used any more. Here's an alphabetical list of element symbols with the corresponding element name. Keep in mind that the names for the elements (and their symbols) may be different in languages other than English.

From: about.com

Ideal Gas Law - Worked Chemistry Problems

2009-02-19T15:21:00.000-08:00

By Anne Marie Helmenstine, Ph.D., About.com

Problem

A hydrogen gas thermometer is found to have a volume of 100.0 cm³ when placed in an ice-water bath at 0°C. When the same thermometer is immersed in boiling liquid chlorine, the volume of hydrogen at the same pressure is found to be 87.2 cm³. What is the temperature of the boiling point of chlorine?

Solution

For hydrogen, PV = nRT, where P is pressure, V is volume, n is number of moles, R is the gas constant, and T is temperature.

Initially:

P₁ = P, V₁ = 100 cm³, n₁ = n, T₁ = 0 + 273 = 273 K

PV₁ = nRT₁

Finally:

P₂ = P, V₂ = 87.2 cm³, n₂ = n, T₂ = ?

PV₂ = nRT₂

Note that P, n, and R are the same. Therefore, the equations may be rewritten:

P/nR = T₁/V₁ = T₂/V₂

and T₂ = V₂T₁/V₁

Plugging in the values we know:

T₂ = 87.2 cm³ x 273 K / 100.0 cm³

T₂ = 238 K

Answer

238 K (which could also be written as -35°C)

From: about.com

Why Do Batteries Discharge More Quickly in Cold Weather?

2009-02-05T22:35:00.000-08:00

Why Do Batteries Discharge More Quickly in Cold Weather?

By Anne Marie Helmenstine, Ph.D., About.com

Answer: The electric current generated by a battery is produced when a connection is made between its positive and negative terminals. When the terminals are connected, a chemical reaction is initiated that generates electrons to supply the current of the battery. Lowering the temperature causes chemical reactions to proceed more slowly, so if a battery is used at a low temperature then less current is produced than at a higher temperature. As the batteries run down they quickly reach the point where they cannot deliver enough current to keep up with the demand. If the battery is warmed up again it will operate normally.

One solution to this problem is to make certain batteries are warm just prior to use. Preheating batteries is not unusual for certain situations. If the battery is already warm and insulated, it may make sense to use the battery's own power to operate a heating coil. It is reasonable to have batteries warm for use, but the discharge curve for most batteries is more dependent on battery design and chemistry than on temperature. This means that if the current drawn by the equipment is low in relation to the power rating of the cell, then the effect of temperature may be negligible.

On the other hand, when a battery is not in use, it will slowly lose its charge as a result of leakage between the terminals. This chemical reaction is also temperature dependent, so unused batteries will lose their charge more slowly at cooler temperatures than at warmer temperatures. For example, certain rechargeable batteries may go flat in approximately two weeks at normal room temperature, but may last more than twice as long if refrigerated.

From: about.com

Measurement of Heat Flow & Enthalpy Change Coffee Cup and Bomb Calorimetry

2009-02-05T22:34:00.000-08:00

Measurement of Heat Flow & Enthalpy Change

Coffee Cup and Bomb Calorimetry
A calorimeter is a device used to measure the quantity of heat flow in a chemical reaction. Two of the most common types of calorimeters are the coffee cup calorimeter and the bomb calorimeter.

Coffee Cup Calorimeter

A coffee cup calorimeter is essentially a polystyrene (Styrofoam) cup with a lid. The cup is partially filled with a known volume of water and a thermometer is inserted through the lid of the cup so that its bulb is below the water surface. When a chemical reaction occurs in the coffee cup calorimeter, the heat of the reaction if absorbed by the water. The change in the water temperature is used to calculate the amount of heat that has been absorbed (used to make products, so water temperature decreases) or evolved (lost to the water, so its temperature increases) in the reaction.

Heat flow is calculated using the relation:

q = (specific heat) x m x Dt

where q is heat flow, m is mass in grams, and Dt is the change in temperature. The specific heat is the amount of heat required to raise the temperature of 1 gram of a substance 1 degree Celsius. The specific heat of water is 4.18 J/(g·&degC).

For example, consider a chemical reaction which occurs in 200 grams of water with an initial temperature of 25.0&degC. The reaction is allowed to proceed in the coffee cup calorimeter. As a result of the reaction, the temperature of the water changes to 31.0&degC. The heat flow is calculated:

q_water = 4.18 J/(g·&degC) x 200 g x (31.0&degC - 25.0&degC)

q_water = +5.0 x 10³ J

In other words, the products of the reaction evolved 5000 J of heat, which was lost to the water. The enthalpy change, DH, for the reaction is equal in magnitude but opposite in sign to the heat flow for the water:

DH_reaction = -(q_water)

Recall that for an exothermic reaction, DH <>water is positive. The water absorbs heat from the reaction and an increase in temperature is seen. For an endothermic reaction, DH > 0; q_water is negative. The water supplies heat for the reaction and a decrease in temperature is seen.

Bomb Calorimeter

A coffee cup calorimeter is great for measuring heat flow in a solution, but it can't be used for reactions which involve gases, since they would escape from the cup. The coffee cup calorimeter can't be used for high temperature reactions, either, since these would melt the cup. A bomb calorimeter is used to measure heat flows for gases and high temperature reactions.

A bomb calorimeter works in the same manner as a coffee cup calorimeter, with one big difference. In a coffee cup calorimeter, the reaction takes place in the water. In a bomb calorimeter, the reaction takes place in a sealed metal container, which is placed in the water in an insulated container. Heat flow from the reaction crosses the walls of the sealed container to the water. The temperature difference of the water is measured, just as it was for a coffee cup calorimeter. Analysis of the heat flow is a bit more complex than it was for the coffee cup calorimeter because the heat flow into the metal parts of the calorimeter must be taken into account:

q_reaction = - (q_water + q_bomb)

where q_water = 4.18 J/(g·&degC) x m_water x Dt

The bomb has a fixed mass and specific heat. The mass of the bomb multiplied by its specific heat is sometimes termed the calorimeter constant, denoted by the symbol C with units of joules per degree Celsius. The calorimeter constant is determined experimentally and will vary from one calorimeter to the next. The heat flow of the bomb is:

q_bomb = C x Dt

Once the calorimeter constant is known, calculating heat flow is a simple matter. The pressure within a bomb calorimeter often changes during a reaction, so the heat flow may not be equal in magnitude to the enthalpy change.

from: about.com

How To Create an Exothermic Chemical Reaction

2009-02-04T19:13:00.000-08:00

How To Create an Exothermic Chemical Reaction

By Anne Marie Helmenstine, Ph.D., About.com

Exothermic chemical reactions produce heat. In this reaction vinegar is used to remove the protective coating from steel wool, allowing it to rust. When the iron combines with oxygen, heat is released.
Difficulty: Average
Time Required: 15 minutes
Here's How:

1. Place the thermometer in the jar and close the lid. Allow about 5 minutes for the thermometer to record the temperature, then open the lid and read the thermometer.
2. Remove the thermometer from the jar (if you didn't already in Step 1).
3. Soak a piece of steel wool in vinegar for 1 minute.
4. Squeeze the excess vinegar out of the steel wool.
5. Wrap the wool aroung the thermometer and place the wool/thermometer in the jar, sealing the lid.
6. Allow 5 minutes, then read the temperature and compare it with the first reading.
7. Chemistry is Fun!

Tips:

1. Not only does the vinegar remove the protective coating on the steel wool, but once the coating is off its acidity aids in oxidation (rust) of the iron in the steel.
2. The thermal energy given off during this chemical reaction causes the mercury in the thermometer to expand and rise up the column of the thermometer tube.
3. In the rusting of iron, four atoms of solid iron react with three molecules of oxygen gas to form two molecules of solid rust (iron oxide).

What You Need:

* Thermometer
* Jar with Lid
* Steel Wool
* Vinegar

From: about.com

How To Create an Endothermic Chemical Reaction (Safe)

2009-02-04T18:43:00.003-08:00

How To Create an Endothermic Chemical Reaction (Safe)

By Anne Marie Helmenstine, Ph.D., About.com
Most endothermic reactions contain toxic chemicals, but this reaction is safe and easy. Use it as a demonstration or vary the amounts of citric acid and sodium bicarbonate to make an experiment.
Difficulty: Average
Time Required: Minutes
Here's How:

1. Pour the citric acid solution in a styrofoam coffee cup. Use a thermometer or other temperature probe to record the initial temperature.
2. Stir in the baking soda (sodium bicarbonate). Track the change in temperature as a function of time.
3. The reaction is: H3C6H5O7(aq) + 3 NaHCO3(s) --> 3 CO2(g) + 3 H2O(l) + NaC6H5O7(aq)
4. When you have completed your demonstration or experiment, simply wash the cup out in a sink. No toxic chemicals to mess with!

Tips:

1. Feel free to vary the concentration of the citric acid solution or the quantity of sodium bicarbonate.
2. An endothermic is a reaction that requires energy to proceed. The intake of energy may be observed as a decrease in temperature as the reaction proceeds. Once the reaction is complete, the temperature of the mixture will return to room temperature.

What You Need:

* 25 ml citric acid soln
* 15 g baking soda
* styrofoam cup
* thermometer
* stirring rod

From: about.com

Top Reasons Why Students Fail Chemistry

2009-02-04T18:43:00.002-08:00

Top Reasons Why Students Fail Chemistry

By Anne Marie Helmenstine, Ph.D., About.com

Are you taking a chemistry class? Are you worried you might not pass? Chemistry is a subject many students prefer to avoid, even if they have an interest in science, because of its reputation for lowering grade point averages. However, it isn't as bad as it seems, especially if you avoid these common mistakes.
1. Procrastinating
Never do today what you can put off until tomorrow, right? Wrong! The first few days in a chemistry class may be very easy and could lull you into a false sense of security. Don't put off doing homework or studying until halfway through the class. Mastering chemistry requires you to build concept upon concept. If you miss the basics, you'll get yourself into trouble. Pace yourself. Set aside a small segment of time each day for chemistry. It will help you to gain long-term mastery. Don't cram.

2. Insufficient Math Preparation
Don't go into chemistry until you understand the basics of algebra. Geometry helps, too. You will need to be able to perform unit conversions. Expect to work chemistry problems on a daily basis. Don't rely too much on a calculator. Chemistry and physics use math as an essential tool.
3. Not Getting or Reading the Text
Yes, there are classes in which the text is optional or completely useless. This isn't one of those classes. Get the text. Read it! Ditto for any required lab manuals. Even if the lectures are fantastic, you'll need the book for the homework assignments. A study guide may be of limited use, but the basic text is a must-have.
4. Psyching Yourself Out
I think I can, I think I can... you have to have a positive attitude toward chemistry. If you truly believe you will fail you may be setting yourself up for a self-fulfilling prophecy. If you have prepared yourself for the class, you have to believe that you can be successful. Also, it's easier to study a topic you like than one you hate. Don't hate chemistry. Make your peace with it and master it.
5. Not Doing Your Own Work
Study guides and books with worked answers in the back are great, right? Yes, but only if you use them for help and not as an easy way to get your homework done. Don't let a book or classmates do your work for you. They won't be available during the tests, which will count for a big portion of your grade.

From: about.com

Why You Should Get your Doctoral Degree Going for the Ph.D.

2009-02-04T18:43:00.001-08:00

If you are interested in a chemistry or another science career, there are multiple reasons why you should consider pursuing your doctorate, rather than stopping at a master's degree or a bachelor's degree:

More Money

Let's start with a compelling reason for higher education -- money. There is no guarantee that having a terminal degree will earn the big bucks (don't get into science for money), but there are several states and companies that compute salaries based on education. The education can count for several years of experience. In some situations, a Ph.D. has access to a pay scale not offered to persons without the terminal degree, no matter how much experience he or she has.

More Career Options

In the US, you can't teach college level courses without at least 18 graduate hours in the same field of study. However, Ph.D.s technically can teach college courses in any field. In academia, a Master's degree may provide a glass ceiling for advancement, especially to management positions. The terminal degree offers more research options, including some lab management positions not available otherwise, as well as post-doctoral positions.

Prestige

In addition to getting the 'Doctor' in front of your name, having a Ph.D. commands a certain level of respect, particularly in scientific and academic circles. There are individuals who feel a Ph.D. is pretentious (a Piled higher and Deeper degree), but with work experience too, even these folk usually concede a Ph.D. is an expert in his or her field.

More Affordable Education

If you are seeking a Master's degree, you will probably have to pay for it. On the other hand, teaching and research assistantships and tuition reimbursement usually are available for doctoral candidates. It would cost a school or research facility considerably more money to pay outright for such skilled labor. Don't feel you have to get a Master's degree before pursuing a Doctorate. Different schools have different requirements, but a Bachelor's degree is usually sufficient to get admitted into a Ph.D. program.

What Are Some Careers in Chemistry?

2009-02-04T18:43:00.000-08:00

What Are Some Careers in Chemistry?

Answer: The career options in chemistry are practically endless! However, your employment options depend on how far you have taken your education. A 2-year degree in chemistry won't get you very far. You could work in some labs washing glassware or assist at a school with lab preparation, but you wouldn't have much advancement potential and you could expect a high level of supervision. A college bachelor's degree in chemistry (B.A., B.S.) opens up more opportunities. A 4-year college degree can be used to gain admittance to advanced degree programs (e.g., graduate school, medical school, law school). With the bachelor's degree, you can get a bench job, which would allow you to run equipment and prepare chemicals. A bachelor's degree in chemistry or education (with a lot of chemistry) is necessary to teach at the K-12 level. A master's degree in chemistry, chemical engineering, or other field opens up far more options. A terminal degree, such as a Ph.D. or M.D., leaves the field wide open. In the United States you need at least 18 graduate credit hours to teach at the college level (preferable a Ph.D.). Most scientists who design and supervise their own research programs have terminal degrees. Chemistry is a part of biology and physics, plus, there are lots of categories of chemistry! Here's look at some of the career options related to chemistry:

Chemistry
Ethnobotany
Environmental Law
Patent Law
Technical Writing
Pharmaceuticals
Oceanography
Software Design
Space Exploration
Government Policy
Forensic Science
Biotechnology
Metallurgy
Ceramics Industry
Plastics Industry
Paper Industry
Medicine
Teaching
Engineering
Geochemistry
Agrochemistry
Military Systems

This list isn't remotely complete. You can work chemistry into any industrial, educational, scientific, or governmental field. Chemistry is a very versatile science. Mastery of chemistry is associated with excellent analytical and mathematical skills. Students of chemistry are able to solve problems and think things through. These skills are useful for any job!

From: about.com

Preface to Concept Development Studies in Chemistry

2009-02-03T18:37:00.004-08:00

Preface to Concept Development Studies in Chemistry

Module by: John S. Hutchinson

Summary: This is the preface to the Concept Development Studies in Chemistry, a series of modules for introducing chemical concepts in a General Chemistry course.

Why Concept Development Studies?

The body of knowledge called Science consists primarily of models and concepts, based on observations and deduced from careful reasoning. Viewed in this way, Science is a creative human endeavor. The models, concepts, and theories we use to describe nature are accomplishments equal in creativity to any artistic, musical, or literary work.

Unfortunately, textbooks in Chemistry traditionally present these models and concepts essentially as established facts, stripped of the clever experiments and logical analyses which give them their human essence. As a consequence, students are typically trained to memorize and apply these models, rather than to analyze and understand them. As a result, creative, analytical students are inclined to feel that they cannot "do" Chemistry, that they cannot understand the concepts, or that Chemistry is dull and uninteresting.

This collection of Concept Development Studies in Chemistry is presented to redirect the focus of learning. In each concept development study, a major chemical concept is developed and refined by analysis of experimental observations and careful reasoning. Each study begins with the definition of an initial Foundation of assumed knowledge, followed by a statement of questions which arise from the Foundation. Analysis of these questions is presented as a series of observations and logical deductions, followed by further questions. This detailed process is followed until the conceptual development of a model provides a reasonable answer to the stated questions.

Concept Development Studies in Chemistry is written with two benefits to the reader in mind. First, by analyzing each concept development through critical reasoning, you will gain a much deeper understanding of a significant concept. In addition to knowing how to work with a model, you will have both an understanding of why the model is believable and an appreciation of the essential beauty of the model. Second, the reasoning required to understand these concept development studies will enhance your development of critical, analytical thinking, a skill which is most important to success in Science. As a note, these studies are not intended to be historical developments, although the experiments presented are the ones which led to the concepts discussed. Only a small amount of historical information has been included for perspective.

How to Study the Concept Development Studies

You should study each concept development study, not by memorization, but by carefully thinking about the experiments and the logical development of the concepts and models. Each study is short, and is meant to be read slowly and meticulously. Each sentence contains substance to be studied and understood. You should, at each step in the analysis, challenge yourself as to whether you can reproduce the reasoning leading to the next conclusion. One good way to do this is to outline the concept development study, making sure you understand how each piece of the argument contributes to the development of a concept or model.

It is very important to understand that scientific models and theories are almost never "proven," unlike mathematical theorems. Rather, they are logically developed and deduced to provide simple explanations of observed phenomenon. As such, you will discover many times in these concept development studies when a conclusion is not logically required by an observation and a line of reasoning. Instead, we may arrive at a model which is the simplest explanation of a set of observations, even if it is not the only one. Scientists most commonly abide by the principle of Occam's razor, one statement of which might be that the explanation which requires the least assumptions is the best one.

One very important way to challenge your understanding is to study in a group in which you take turns explaining the development of the model. The ability to explain a concept is a much stronger indicator of your understanding than the ability to solve a problem using the concept. Use the questions at the end of the concept development studies to practice your skill at explaining technical arguments clearly and concisely.

Acknowledgments

My own thinking in writing Concept Development Studies in Chemistry has been strongly influenced by three books:
• The Historical Development of Chemical Concepts, by Roman Mierzecki; • The History of Chemistry, by John Hudson; • Chemical Principles, by Richard Dickerson, Harry Gray, and Gilbert Haight.

I am deeply appreciative of the contributions of Joanna Fair, Karen Aiani Stevens, Kevin Ausman, and Karin Wright in reviewing and criticizing early drafts of the manuscript for this text. I am also indebted to Susan Wiediger, not just for her technical expertise and her knowledge of the educational literature, but also for her commitment to the concept behind this book and this approach to teaching. I appreciate the hard work of Jeffrey Silverman to convert these documents for use in the Connexions Project at Rice University.

Concept Development Studies in Chemistry would not have been written were it not for the encouragement of my wife Paula, who reminded me at the most difficult of times that writing it was the right thing to do. I will be forever grateful.

JSHfrom: http://cnx.org/content/m12594/latest/

The Ideal Gas Law

2009-02-03T18:37:00.003-08:00

The Ideal Gas Law

Module by: John S. Hutchinso

Foundation

We assume as our starting point the atomic
molecular theory. That is, we assume that all matter is composed of
discrete particles. The elements consist of identical atoms, and
compounds consist of identical molecules, which are particles
containing small whole number ratios of atoms. We also assume that
we have determined a complete set of relative atomic weights,
allowing us to determine the molecular formula for any
compound.

Goals

The individual molecules of different
compounds have characteristic properties, such as mass, structure,
geometry, bond lengths, bond angles, polarity, diamagnetism or
paramagnetism. We have not yet considered the properties of mass
quantities of matter, such as density, phase (solid, liquid or gas)
at room temperature, boiling and melting points, reactivity, and so
forth. These are properties which are not exhibited by individual
molecules. It makes no sense to ask what the boiling point of one
molecule is, nor does an individual molecule exist as a gas, solid,
or liquid. However, we do expect that these material or bulk
properties are related to the properties of the individual
molecules. Our ultimate goal is to relate the properties of the
atoms and molecules to the properties of the materials which they
comprise.

Achieving this goal will require considerable
analysis. In this Concept Development Study, we begin at a somewhat
more fundamental level, with our goal to know more about the nature
of gases, liquids and solids. We need to study the relationships
between the physical properties of materials, such as density and
temperature. We begin our study by examining these properties in
gases.

From: http://cnx.org/content/m12594/latest/

Energetics of Chemical Reactions

2009-02-03T18:37:00.002-08:00

Energetics of Chemical Reactions

Module by: John S. Hutchinson

The Foundation

We begin our study of the energetics of
chemical reactions with our understanding of mass relationships,
determined by the stoichiometry of balanced reactions and the
relative atomic masses of the elements. We will assume a conceptual
understanding of energy based on the physics of mechanics, and in
particular, we will assume the law of conservation of energy. In
developing a molecular understanding of the reaction energetics, we
will further assume our understanding of chemical bonding via
valence shell electron pair sharing and molecular orbital
theory.

Goals

The heat released or consumed in a chemical
reaction is typically amongst the most easily observed and most
readily appreciated consequences of the reaction. Many chemical
reactions are performed routinely specifically for the purpose of
utilizing the heat released by the reaction.

We are interested here in an understanding of
the energetics of chemical reactions. Specifically, we wish to know
what factors determine whether heat is absorbed or released during
a chemical reaction. With that knowledge, we seek to quantify and
predict the amount of heat anticipated in a chemical reaction. We
expect to find that the quantity of heat absorbed or released
during a reaction is related to the bonding of the molecules
involved in the reaction.

Prior to answering these questions, we must
first answer a few questions regarding the nature of heat. Despite
our common familiarity with heat (particularly in Houston), the
concept of heat is somewhat elusive to define. We recognize heat as
"whatever it is that makes things hot," but this definition is too
imprecise to permit measurement or any other conceptual progress.
Exactly how do we define and measure heat?Energetics of Chemical Reactions

Module by: John S. Hutchinson

The Foundation

We begin our study of the energetics of
chemical reactions with our understanding of mass relationships,
determined by the stoichiometry of balanced reactions and the
relative atomic masses of the elements. We will assume a conceptual
understanding of energy based on the physics of mechanics, and in
particular, we will assume the law of conservation of energy. In
developing a molecular understanding of the reaction energetics, we
will further assume our understanding of chemical bonding via
valence shell electron pair sharing and molecular orbital
theory.

Goals

The heat released or consumed in a chemical
reaction is typically amongst the most easily observed and most
readily appreciated consequences of the reaction. Many chemical
reactions are performed routinely specifically for the purpose of
utilizing the heat released by the reaction.

We are interested here in an understanding of
the energetics of chemical reactions. Specifically, we wish to know
what factors determine whether heat is absorbed or released during
a chemical reaction. With that knowledge, we seek to quantify and
predict the amount of heat anticipated in a chemical reaction. We
expect to find that the quantity of heat absorbed or released
during a reaction is related to the bonding of the molecules
involved in the reaction.

Prior to answering these questions, we must
first answer a few questions regarding the nature of heat. Despite
our common familiarity with heat (particularly in Houston), the
concept of heat is somewhat elusive to define. We recognize heat as
"whatever it is that makes things hot," but this definition is too
imprecise to permit measurement or any other conceptual progress.
Exactly how do we define and measure heat?From: http://cnx.org/content/m12594/latest/

2009-02-03T18:37:00.001-08:00

Chemical Bonding and Molecular Energy Levels

Module by: John S. Hutchinson

Summary: A development of the quantum mechanical concepts of bonding using valence bond and molecular orbital descriptions to account of bond strength and molecular ionization energies.

Foundation

Our basis for understanding chemical bonding and the structures of molecules is the electron orbital description of the structure and valence of atoms, as provided by quantum mechanics. We assume an understanding of the periodicity of the elements based on the nuclear structure of the atom and our deductions concerning valence based on electron orbitals.

Goals

Our model of valence describes a chemical bond as resulting from the sharing of a pair of electrons in the valence shell of the bonded atoms. This sharing allows each atom to complete an octet of electrons in its valence shell, at least in the sense that we count the shared electrons as belonging to both atoms. However, it is not clear that this electron counting picture has any basis in physical reality. What is meant, more precisely, by the sharing of the electron pair in a bond, and why does this result in the bonding of two atoms together? Indeed, what does it mean to say that two atoms are bound together? Furthermore, what is the significance of sharing a pair of electrons? Why aren’t chemical bonds formed by sharing one or three electrons, for example?

We seek to understand how the details of chemical bonding are related to the properties of the molecules formed, particularly in terms of the strengths of the bonds formed.

Observation 1: Bonding with a Single Electron

We began our analysis of the energies and motions of the electrons in atoms by observing the properties of the simplest atom, hydrogen, with a single electron. Similarly, to understand the energies and motions of electrons which lead to chemical bonding, we begin our observations with the simplest particle with a chemical bond, which is the
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {} molecular ion. Each hydrogen nucleus has a charge of +1. An
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {} molecular ion therefore has a single electron. It seems inconsistent with our notions of valence that a single electron, rather than an electron pair, can generate a chemical bond. However, these concepts have been based on observations on molecules, not molecular ions like
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {}. And it is indeed found that
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {} is a stable bound molecular ion.

What forces and motions hold the two hydrogen nuclei close together in the
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {} ion? It is worth keeping in mind that the two nuclei must repel one another, since they are both positively charged. In the absence of the electron, the two nuclei would accelerate away from one another, rather than remaining in close proximity. What is the role of the electron? Clearly, the electron is attracted to both nuclei at the same time, and, in turn, each nucleus is attracted to the electron. The effect of this is illustrated in Fig. 1. In Fig. 1a, the electron is “outside” of the two nuclei. In this position, the electron is primarily attracted to the nucleus on the left, to which it is closer. More importantly, the nucleus on the right feels a greater repulsion from the other nucleus than attraction to the electron, which is farther away. As a result, the nucleus on the right experiences a strong force driving it away from the hydrogen atom on the left. This arrangement does not generate chemical bonding, therefore. By contrast, in Fig. 1b, the electron is between the two nuclei. In this position, the electron is roughly equally attracted to the two nuclei, and very importantly, each nucleus feels an attractive force to the electron which is greater than the repulsive force generated by the other nucleus. Focusing on the electron’s energy, the proximity of the two nuclei provides it a doubly attractive environment with a very low potential energy. If we tried to pull one of the nuclei away, this would raise the potential energy of the electron, since it would lose attraction to that nucleus. Hence, to pull one nucleus away requires us to add energy to the molecular ion. This is what is meant by a chemical bond: the energy of the electrons is lower when the atoms are in close proximity than when the atoms are far part. This “holds” the nuclei close together, since we must do work (add energy) to take the nuclei apart.
Figure 1Figure 1 (graphics1.png)

Note that the chemical bond in Fig. 1b results from the electron’s position between the nuclei. On first thought, this appears to answer our question of what we mean by “sharing an electron pair” to form a chemical bond. An electron positioned between two nuclei is “shared” to the extent that its potential energy is lowered due to attraction to both nuclei simultaneously.

On second thought, though, this description must be inaccurate. We have learned our study of Energy Levels in Atoms that an electron must obey the uncertainty principle and that, as a consequence, the electron does not have a definite position, between the nuclei or otherwise. We can only hope to specify a probability for observing an electron in a particular location. This probability is, from quantum mechanics, provided by the wave function. What does this probability distribution look like for the
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {} molecular ion?

To answer this question, we begin by experimenting with a distribution that we know: the 1s electron orbital in a hydrogen atom. This we recall has the symmetry of a sphere, with equal probability in all directions away from the nucleus. To create an
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {} molecular ion from a hydrogen atom, we must add a bare second hydrogen nucleus (an
H+H+ size 12{H rSup { size 8{+{}} } } {}ion). Imagine bringing this nucleus closer to the hydrogen atom from a very great distance (see Fig. 2a). As the
H+H+ size 12{H rSup { size 8{+{}} } } {}ion approaches the neutral atom, both the hydrogen atom’s nucleus and electron respond to the electric potential generated by the positive charge. The electron is attracted and the hydrogen atom nucleus is repelled. As a result, the distribution of probability for the electron about the nucleus must become distorted, so that the electron has a greater probability of being near the
H+H+ size 12{H rSup { size 8{+{}} } } {} ion and the nucleus has a greater probability of being farther from the ion. This distortion, illustrated in Fig. 2b, is called “polarization”: the hydrogen atom has become like a “dipole”, with greater negative charge to one side and greater positive charge to the other.
Figure 2Figure 2 (graphics2.png)

This polarization must increase as the
H+H+ size 12{H rSup { size 8{+{}} } } {} ion approaches the hydrogen atom until, eventually, the electron orbital must be sufficiently distorted that there is equal probability for observing the electron in proximity to either hydrogen nucleus (see Fig. 2c). The electron probability distribution in Fig. 2c now describes the motion of the electron, not in a hydrogen atom, but in an
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {} molecular ion. As such, we refer to this distribution as a “molecular orbital.”

We note that the molecular orbital in Fig. 2c is more delocalized than the atomic orbital in Fig. 2a, and this is also important in producing the chemical bond. We recall from the discussion of Atomic Energy Levels that the energy of an electron in an orbital is determined, in part, by the compactness of the orbital. The more the orbital confines the motion of the electron, the higher is the kinetic energy of the electron, an effect we referred to as the “confinement energy.” Applying this concept to the orbitals in Fig. 2, we can conclude that the confinement energy is lowered when the electron is delocalized over two nuclei in a molecular orbital. This effect contributes significantly to the lowering of the energy of an electron resulting from sharing by two nuclei.

Recall that the electron orbitals in the hydrogen atom are described by a set of quantum numbers. One of these quantum numbers is related to the symmetry or shape of the atomic orbital and is generally depicted by a letter. Recall that an “s” orbital is spherical in shape, and a “p” orbital has two lobes aligned along one axis. Similarly, the molecular orbitals for the
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {} molecular ion are described by a set of numbers which give the symmetry (or shape) of the orbital. For our purposes, we need only one of these descriptors, based on the symmetry of the orbital along the bond: if the molecular orbital has the symmetry of a cylinder, we refer to it as a “
σσ size 12{σ} {} orbital.” The orbital in Fig. 2c satisfies this condition.

We conclude that chemical bonding results from an electron in a molecular orbital which has substantial probability for the electron to be between two nuclei. However, this example illustrates chemical bonding with a single electron. Our rules of valence indicate that bonding typically occurs with a pair of electrons, rather than a single electron. Furthermore, this model of bonding does not tell us how to handle molecules with many electrons (say,
F2F2 size 12{F rSub { size 8{2} } } {}) where most of the electrons do not participate in the bonding at all.

Observation 2: Bonding and Non-Bonding in Diatomic Molecules

We now consider molecules with more than one electron. These are illustrated most easily by diatomic molecules (molecules with only two atoms) formed by like atoms, beginning with the hydrogen molecule,
H2H2 size 12{H rSub { size 8{2} } } {}. The most direct experimental observation of a chemical bond is the amount of energy required to break it. This is called the bond energy, or somewhat less precisely, the bond strength. Experimentally, it is observed that the bond energy of the hydrogen molecule
H2H2 size 12{H rSub { size 8{2} } } {} is 458 kJ/mol. By contrast, the bond energy of the
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {} molecular ion is 269 kJ/mol. Therefore, the bond in
H2H2 size 12{H rSub { size 8{2} } } {} is stronger than the bond in
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {}. Thus, the pair of shared electrons in
H2H2 size 12{H rSub { size 8{2} } } {} generates a stronger attractive force than does the single electron in
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {}.

Before deducing an explanation of this in terms of electron orbitals, we first recall the valence shell electron pair description of the bonding in
H2H2 size 12{H rSub { size 8{2} } } {}. Each hydrogen atom has a single electron. By sharing these two electrons, each hydrogen atom can fill its valence shell, attaining the electron configuration of helium.

How does this translate into the electron orbital picture of electron sharing that we have just described for the
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {} molecular ion? There are two ways to deduce the answer to this question, and, since they are both useful, we will work through them both. The first way is to imagine that we form an
H2H2 size 12{H rSub { size 8{2} } } {} molecule by starting with an
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {} molecular ion and adding an electron to it. As a simple approximation, we might imagine that the first electron’s probability distribution (its orbital) is not affected by the addition of the second electron. The second electron must have a probability distribution describing its location in the molecule as well. We recall that, in atoms, it is possible to put two electrons into a single electron orbital, provided that the two electrons have opposite values of the spin quantum number, ms. Therefore, we expect this to be true for molecules as well, and we place the added second electron in
H2H2 size 12{H rSub { size 8{2} } } {} into the same
σσ size 12{σ} {} orbital as the first. This results in two electrons in the region between the two nuclei, thus adding to the force of attraction of the two nuclei into the bond. This explains our observation that the bond energy of
H2H2 size 12{H rSub { size 8{2} } } {} is almost (although not quite) twice the bond energy of
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {}.

The second way to understand the electron orbital picture of
H2H2 size 12{H rSub { size 8{2} } } {} is to imagine that we form the molecule by starting with two separated hydrogen atoms. Each of these atoms has a single electron in a 1s orbital. As the two atoms approach one another, each electron orbital is polarized in the direction of the other atom. Once the atoms are close enough together, these two orbitals become superimposed. Now we must recall that these orbitals describe the wave-like motion of the electron, so that, when these two wave functions overlap, they must interfere, either constructively or destructively. In Fig. 3, we see the consequences of constructive and destructive interference. We can deduce that, in
H2H2 size 12{H rSub { size 8{2} } } {} the electron orbitals from the atoms must constructively interfere, because that would increase the electron probability in the region between the nuclei, resulting in bonding as before. Therefore, the
σσ size 12{σ} {} molecular orbital describing the two electrons in
H2H2 size 12{H rSub { size 8{2} } } {} can be understood as resulting from the constructive overlap of two atomic 1s electron orbitals.

We now add to our observations of diatomic molecules by noting that, of the diatomic molecules formed from like atoms of the first ten elements,
H2H2 size 12{H rSub { size 8{2} } } {},
Li2Li2 size 12{ ital "Li" rSub { size 8{2} } } {},
B2B2 size 12{B rSub { size 8{2} } } {},
C2C2 size 12{C rSub { size 8{2} } } {},
N2N2 size 12{N rSub { size 8{2} } } {},
O2O2 size 12{O rSub { size 8{2} } } {}, and
F2F2 size 12{F rSub { size 8{2} } } {} are stable molecules with chemical bonds, whereas
He2He2 size 12{ ital "He" rSub { size 8{2} } } {},
Be2Be2 size 12{ ital "Be" rSub { size 8{2} } } {}, and
Ne2Ne2 size 12{ ital "Ne" rSub { size 8{2} } } {} are not bound. In examining the electron configurations of the atoms of these elements, we discover a correspondence with which diatomic molecules are bound and which ones are not.
H,Li,B,N,andFH,Li,B,N,andF size 12{H, ital "Li",B,N,` ital "and"`F} {} all have odd numbers of electrons, so that at least one electron in each atom is unpaired. By contrast, He, Be, and Ne all have even numbers of electrons, none of which are unpaired. The other atoms, C and O both have an even number of electrons. However, as deduced in our understanding of the electron configurations in atoms, electrons will, when possible, distribute themselves into different orbitals of the same energy so as to reduce the effect of their mutual repulsion. Thus, in C and O, there are three 2p orbitals into which 2 and 4 electrons are placed, respectively. Therefore, in both atoms, there are two unpaired electrons. We conclude that bonds will form between atoms if and only if there are unpaired electrons in these atoms.

In
H2H2 size 12{H rSub { size 8{2} } } {}, the unpaired electrons from the separated atoms become paired in a molecular orbital formed from the overlap of the 1s atomic electron orbitals. In the case of a hydrogen atom, then, there are of course no paired electrons in the atom to worry about. In all other atoms, there certainly are paired electrons, regardless of whether there are or are not unpaired electrons. For example, in a lithium atom, there are two paired electrons in a 1s orbital and an unpaired electron in the 2s orbital. To form
Li2Li2 size 12{ ital "Li" rSub { size 8{2} } } {}, the unpaired electron from each atom can be placed into a molecular orbital formed from the overlap of the 2s atomic electron orbitals. However, what becomes of the two electrons paired in the 1s orbital in a Li atom during the bonding of
Li2Li2 size 12{ ital "Li" rSub { size 8{2} } } {}?

To answer this question, we examine
He2He2 size 12{ ital "He" rSub { size 8{2} } } {}, in which each atom begins with only the two 1s electrons. As we bring the two He atoms together from a large distance, these 1s orbitals should become polarized, as in the hydrogen atom. When the polarized 1s orbitals overlap, constructive interference will again result in a
σσ size 12{σ} {} molecular orbital, just as in
H2H2 size 12{H rSub { size 8{2} } } {}. Yet, we observe that
He2He2 size 12{ ital "He" rSub { size 8{2} } } {} is not a stable bound molecule. The problem which prevents bonding for
He2He2 size 12{ ital "He" rSub { size 8{2} } } {} arises from the Pauli Exclusion Principle: only two of the four electrons in
He2He2 size 12{ ital "He" rSub { size 8{2} } } {} can be placed into this
σσ size 12{σ} {} bonding molecular orbital. The other two must go into a different orbital with a different probability distribution. To deduce the form of this new orbital, we recall that the bonding orbital discussed so far arises from the constructive interference of the atomic orbitals, as shown in Fig. 3. We could, instead, have assumed destructive interference of these orbitals. Destructive interference of two waves eliminates amplitude in the region of overlap of the waves, also shown in Fig. 3. In the case of the atomic orbitals, this means that the molecular orbital formed from destructive interference decreases probability for the electron to be between in the nuclei. Therefore, it increases probability for the electron to be outside the nuclei, as in Fig. 1a. As discussed there, this arrangement for the electron does not result in bonding; instead, the nuclei repel each other and the atoms are forced apart. This orbital is thus called an anti-bonding orbital. This orbital also has the symmetry of a cylinder along the bond axis, so it is also a
σσ size 12{σ} {}orbital; to indicate that it is an anti-bonding orbital, we designate it with an asterisk,
σ*σ* size 12{σ rSup { size 8{1} } } {}
Figure 3Figure 3 (graphics3.png)

In
He2He2 size 12{ ital "He" rSub { size 8{2} } } {}, both the bonding and the anti-bonding orbitals must be used in order to accommodate four electrons. The two electrons in the bonding orbital lower the energy of the molecule, but the two electrons in the anti-bonding orbital raise it. Since two He atoms will not bind together, then the net effect must be that the anti-bonding orbital more than offsets the bonding orbital.

We have now deduced an explanation for why the paired electrons in an atom do not contribute to bonding. Both bonding and anti-bonding orbitals are always formed when two atomic orbitals overlap. When the electrons are already paired in the atomic orbitals, then there are too many electrons for the bonding molecular orbital. The extra electrons must go into the anti-bonding orbital, which raises the energy of the molecule, preventing the bond from forming.

Returning to the
Li2Li2 size 12{ ital "Li" rSub { size 8{2} } } {} example discussed above, we can develop a simple picture of the bonding. The two 1s electrons from each atom do not participate in the bonding, since the anti-bonding more than offsets the bonding. Thus, the paired “core” electrons remain in their atomic orbitals, unshared, and we can ignore them in describing the bond. The bond is formed due to overlap of the 2s orbitals and sharing of these electrons only. This is also consistent with our earlier view that the core electrons are closer to the nucleus, and thus unlikely to be shared by two atoms.

The model we have constructed seems to describe fairly well the bonding in the bound diatomic molecules listed above. For example, in a fluorine atom, the only unpaired electron is in a 2p orbital. Recall that a 2p orbital has two lobes, directed along one axis. If these lobes are assumed to lie along the axis between the two nuclei in
F2F2 size 12{F rSub { size 8{2} } } {}, then we can overlap them to form a bonding orbital. Placing the two unpaired electrons into this orbital then results in a single shared pair of electrons and a stable molecular bond.

Observation 3: Ionization energies of diatomic molecule

The energies of electrons in molecular orbitals can be observed directly by measuring the ionization energy. This is the energy required to remove an electron, in this case, from a molecule:

H

2

(
g
)
→

H

2

+

(
g

)
+

e

−

(
g
)

H

2

(
g
)
→

H

2

+

(
g

)
+

e

−

(
g
)

size 12{H rSub { size 8{2} } \( g \) rightarrow H rSub { size 8{2} } rSup { size 8{+{}} } \( g \) +e rSup { size 8{ - {}} } \( g \) } {}

The measured ionization energy of
H2H2 size 12{H rSub { size 8{2} } } {} is 1488 kJ/mol. This number is primarily important in comparison to the ionization energy of a hydrogen atom, which is 1312 kJ/mol. Therefore, it requires more energy to remove an electron from the hydrogen molecule than from the hydrogen atom, so we can conclude that the electron has a lower energy in the molecule. If we attempt to pull the atoms apart, we must raise the energy of the electron. Hence, energy is required to break the bond, so the molecule is bound.

We conclude that a bond is formed when the energy of the electrons in the molecule is lower than the energy of the electrons in the separated atoms. This conclusion seems consistent with our previous view of shared electrons in bonding molecular orbitals.

As a second example, we consider the nitrogen molecule,
N2N2 size 12{N rSub { size 8{2} } } {}. We find that the ionization energy of molecular nitrogen is 1503 kJ/mol, and that of atomic nitrogen is 1402 kJ/mol. Once again, we conclude that the energy of the electrons in molecular nitrogen is lower than that of the electrons in the separated atoms, so the molecule is bound.

As a third example, we consider fluorine,
F2F2 size 12{F rSub { size 8{2} } } {}. In this case, we find that the ionization energy of molecular fluorine is 1515 kJ/mol, which is smaller than the ionization energy of a fluorine atom, 1681 kJ/mol. This seems inconsistent with the bonding orbital concept we have developed above, which states that the electrons in the bond have a lower energy than in the separated atoms. If the electron being ionized has a higher energy in
F2F2 size 12{F rSub { size 8{2} } } {} than in F, why is
F2F2 size 12{F rSub { size 8{2} } } {} a stable molecule? Apparently, we need a more complete description of the molecular orbital concept of chemical bonding.

To proceed further, we compare bond energies in several molecules. Recall that the bond energy (or bond strength) is the energy required to separate the bonded atoms. We observe that the bond energy of
N2N2 size 12{N rSub { size 8{2} } } {} is 956 kJ/mol. This is very much larger than the bond energy of
H2H2 size 12{H rSub { size 8{2} } } {}, 458 kJ/mol, and of
F2F2 size 12{F rSub { size 8{2} } } {}, which is 160 kJ/mol. We can account for the unusually strong bond in nitrogen using both our valence shell electron pair sharing model and our electron orbital descriptions. A nitrogen atom has three unpaired electrons in its valence shell, because the three 2p electrons distribute themselves over the three 2p orbitals, each oriented along a different axis. Each of these unpaired electrons is available for sharing with a second nitrogen atom. The result, from valence shell electron pair sharing concepts, is that three pairs of electrons are shared between two nitrogen atoms, and we call the bond in
N2N2 size 12{N rSub { size 8{2} } } {} a “triple bond.” It is somewhat intuitive that the triple bond in
N2N2 size 12{N rSub { size 8{2} } } {} should be much stronger than the single bond in
H2H2 size 12{H rSub { size 8{2} } } {} or in
F2F2 size 12{F rSub { size 8{2} } } {}.

Now consider the molecular orbital description of bonding in
N2N2 size 12{N rSub { size 8{2} } } {}. Each of the three 2p atomic orbitals in each nitrogen atom must overlap to form a bonding molecular orbital, if we are to accommodate three electron pairs. Each 2p orbital is oriented along a single axis. One 2p orbital from each atom is oriented in the direction of the other atom, that is, along the bond axis. When these two atomic orbitals overlap, they form a molecular orbital which has the symmetry of a cylinder and which is therefore a
σσ size 12{σ} {} orbital. Of course, they also form a
σ*σ* size 12{σ rSup { size 8{1} } } {}orbital. The two electrons are then paired in the bonding orbital.
Figure 4Figure 4 (graphics4.png)

The other two 2p orbitals on each nitrogen atom are perpendicular to the bond axis. The constructive overlap between these orbitals from different atoms must therefore result in a molecular orbital somewhat different that what we have discussed before. As shown in Fig. 4, the molecular orbital which results now does not have the symmetry of a cylinder, and in fact, looks something more like a cylinder cut into two pieces. This we call a
ππ size 12{π} {} orbital. There are two such
ππ size 12{π} {} orbitals since there are two sets of p orbitals perpendicular to the bond axis. Figure 4 also shows that an anti-bonding orbital is formed from the destructive overlap of 2p orbitals, and this is called a
π*π* size 12{σ rSup { size 8{1} } } {} orbital. There are also two
π*π* size 12{σ rSup { size 8{1} } } {} orbitals formed from destructive overlap of 2p orbitals. In
N2N2 size 12{N rSub { size 8{2} } } {}, the three shared electron pairs are thus in a single
σσ size 12{σ} {} orbital and in two
ππ size 12{π} {} orbitals. Each of these orbitals is a bonding orbital, therefore all six electrons have their energy lowered in comparison to the separated atoms.

This is depicted in Fig. 5 in what is called a “molecular orbital energy diagram.” Each pair of atomic orbitals, one from each atom, is overlapped to form a bonding and an anti-bonding orbital. The three 2p orbitals from each atom form one
σσ size 12{σ} {} and
σ*σ* size 12{σ rSup { size 8{1} } } {} pair and two
ππ size 12{π} {} and
π*π* size 12{σ rSup { size 8{1} } } {} pairs. The lowering of the energies of the electrons in the
σσ size 12{σ} {}and
ππ size 12{π} {} orbitals is apparent. The ten n=2 electrons from the nitrogen atoms are then placed pairwise, in order of increasing energy, into these molecular orbitals. Note that, in agreement with the Pauli Exclusion Principle, each pair in a single orbital consists of one spin up and one spin down electron.
Figure 5Figure 5 (graphics5.png)

Recall now that we began the discussion of bonding in
N2N2 size 12{N rSub { size 8{2} } } {} because of the curious result that the ionization energy of an electron in
F2F2 size 12{F rSub { size 8{2} } } {} is less than that of an electron in an F atom. By comparing the molecular orbital energy level diagrams for
N2N2 size 12{N rSub { size 8{2} } } {} and
F2F2 size 12{F rSub { size 8{2} } } {} we are now prepared to answer this puzzle. There are five p electrons in each fluorine atom. These ten electrons must be distributed over the molecular orbitals whose energies are shown in Fig. 6. (Note that the ordering of the bonding 2p orbitals differ between
N2N2 size 12{N rSub { size 8{2} } } {} and
F2F2 size 12{F rSub { size 8{2} } } {}.) We place two electrons in the
σσ size 12{σ} {} orbital, four more in the two ππ size 12{π rSup { size 8{ * } } } {} orbitals, and four more in the two
π*π* size 12{σ rSup { size 8{1} } } {} orbitals. Overall, there are six electrons in bonding orbitals and four in anti-bonding orbitals. Since
F2F2 size 12{F rSub { size 8{2} } } {} is a stable molecule, we must conclude that the lowering of energy for the electrons in the bonding orbitals is greater than the raising of energy for the electrons in the antibonding orbitals. Overall, this distribution of electrons is, net, equivalent to having two electrons paired in a single bonding orbital.
Figure 6Figure 6 (graphics6.png)

This also explains why the ionization energy of
F2F2 size 12{F rSub { size 8{2} } } {} is less than that of an F atom. The electron with the highest energy requires the least energy to remove from the molecule or atom. The molecular orbital energy diagram in Fig. 6 clearly shows that the highest energy electrons in
F2F2 size 12{F rSub { size 8{2} } } {} are in anti-bonding orbitals. Therefore, one of these electrons is easier to remove than an electron in an atomic 2p orbital, because the energy of an anti-bonding orbital is higher than that of the atomic orbitals. (Recall that this is why an anti-bonding orbital is, indeed, anti-bonding.) Therefore, the ionization energy of molecular fluorine is less than that of atomic fluorine. This clearly demonstrates the physical reality and importance of the anti-bonding orbitals.

A particularly interesting case is the oxygen molecule,
O2O2 size 12{O rSub { size 8{2} } } {}. In completing the molecular orbital energy level diagram for oxygen, we discover that we must decide whether to pair the last two electrons in the same
2pπ*2pπ* size 12{2pπ rSup { size 8{ * } } } {}
orbital, or whether they should be separated into different
2pπ*2pπ* size 12{2pπ rSup { size 8{ * } } } {} orbitals. To determine which, we note that oxygen molecules are paramagnetic, meaning that they are strongly attracted to a magnetic field. To account for this paramagnetism, we recall that electron spin is a magnetic property. In most molecules, all electrons are paired, so for each “spin up” electron there is a “spin down” electron and their magnetic fields cancel out. When all electrons are paired, the molecule is diamagnetic meaning that it responds only weakly to a magnetic field.

If the electrons are not paired, they can adopt the same spin in the presence of a magnetic field. This accounts for the attraction of the paramagnetic molecule to the magnetic field. Therefore, for a molecule to be paramagnetic, it must have unpaired electrons. The correct molecular orbital energy level diagram for an
O2O2 size 12{O rSub { size 8{2} } } {} molecule is shown in Fig. 7.
Figure 7Figure 7 (graphics7.png)

In comparing these three diatomic molecules, we recall that
N2N2 size 12{N rSub { size 8{2} } } {} has the strongest bond, followed by
O2O2 size 12{O rSub { size 8{2} } } {} and
F2F2 size 12{F rSub { size 8{2} } } {}. We have previously accounted for this comparison with Lewis structures, showing that
N2N2 size 12{N rSub { size 8{2} } } {} is a triple bond,
O2O2 size 12{O rSub { size 8{2} } } {} is a double bond, and
F2F2 size 12{F rSub { size 8{2} } } {} is a single bond. The molecular orbital energy level diagrams in Figs. 5 to 7 cast a new light on this analysis. Note that, in each case, the number of bonding electrons in these molecules is eight. The difference in bonding is entirely due to the number of antibonding electrons: 2 for
N2N2 size 12{N rSub { size 8{2} } } {} , 4 for
O2O2 size 12{O rSub { size 8{2} } } {} , and six for
F2F2 size 12{F rSub { size 8{2} } } {}. Thus, the strength of a bond must be related to the relative numbers of bonding and antibonding electrons in the molecule. Therefore, we now define the bond order as

Bond

Order

=

1
2

# bonding electrons

−

# antibonding electrons

Bond

Order

=

1
2

# bonding electrons

−

# antibonding electrons

size 12{ ital "Bond"` ital "Order"= { {1} over {2} } left ( ital "bonding"` ital "electrons" - ital "antibonding"` ital "electrons" right )} {}

Note that, defined this way, the bond order for
N2N2 size 12{N rSub { size 8{2} } } {} is 3, for
O2O2 size 12{O rSub { size 8{2} } } {} is 2, and for
F2F2 size 12{F rSub { size 8{2} } } {} is 1, which agrees with our conclusions from Lewis structures. We conclude that we can predict the relative strengths of bonds by comparing bond orders.

Review and Discussion Questions

1. Why does an electron shared by two nuclei have a lower potential energy than an electron on a single atom? Why does an electron shared by two nuclei have a lower kinetic energy than an electron on a single atom? How does this sharing result in a stable molecule? How can this affect be measured experimentally?
2. Explain why the bond in an
H2H2 size 12{H rSub { size 8{2} } } {} molecule is almost twice as strong as the bond in the
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {} ion. Explain why the
H2H2 size 12{H rSub { size 8{2} } } {} bond is less than twice as strong as the
H2+H2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {} bond.
3. Be2Be2 size 12{H rSub { size 8{2} } } {} is not a stable molecule. What information can we determine from this observation about the energies of molecular orbitals?
4. Less energy is required to remove an electron from an
F2F2 size 12{F rSub { size 8{2} } } {} molecule than to remove an electron from an F atom. Therefore, the energy of that electron is higher in the molecule than in the atom. Explain why, nevertheless,
F2F2 size 12{F rSub { size 8{2} } } {} is a stable molecule, i.e., the energy of an
F2F2 size 12{F rSub { size 8{2} } } {} molecule is less than the energy of two F atoms.
5. Why do the orbitals of an atom "hybridize" when forming a bond?
6. Calculate the bond orders of the following molecules and predict which molecule in each pair has the stronger bond:
1. C2C2 size 12{C rSub { size 8{2} } } {} or
C2+C2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {}
2. B2B2 size 12{B rSub { size 8{2} } } {} or
B2+B2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {}
3. F2F2 size 12{F rSub { size 8{2} } } {} or
F2-F2- size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {}
4. O2O2 size 12{O rSub { size 8{2} } } {} or
O2+O2+ size 12{H rSub { size 8{2} } rSup { size 8{+{}} } } {}

7. Which of the following diatomic molecules are paramagnetic: CO,
Cl2Cl2 size 12{ ital "Cl" rSub { size 8{2} } } {}, NO,
N2N2 size 12{N rSub { size 8{2} } } {} ?
8. B2B2 size 12{B rSub { size 8{2} } } {} is observed to be paramagnetic. Using this information, draw an appropriate molecular orbital energy level diagram for
B2B2 size 12{B rSub { size 8{2} } } {}.From: http://cnx.org/content/m12594/latest/

Molecular Geometry and Electron Domain Theory

2009-02-03T18:37:00.000-08:00

Molecular Geometry and Electron Domain Theory

Module by: John S. Hutchinson

Foundation

We begin by assuming a
Lewis structure model for chemical bonding based on
valence shell electron pair sharing and the octet rule. We thus
assume the nuclear structure of the atom, and we further assume the
existence of a valence shell of electrons in each atom which
dominates the chemical behavior of that atom. A covalent chemical
bond is formed when the two bonded atoms share a pair of valence
shell electrons between them. In general, atoms of Groups IV
through VII bond so as to complete an octet of valence shell
electrons. A number of atoms, including C, N, O, P, and S, can form
double or triple bonds as needed to complete an octet. We know that
double bonds are generally stronger and have shorter lengths than
single bonds, and triple bonds are stronger and shorter than double
bonds.

Goals

We should expect that the properties of
molecules, and correspondingly the substances which they comprise,
should depend on the details of the structure and bonding in these
molecules. The relationship between bonding, structure, and
properties is comparatively simple in
diatomic molecules, which contain two atoms only,
e.g. HClHCl
or
O2O2.
A
polyatomic molecule contains more than two atoms. An
example of the complexities which arise with polyatomic molecules
is molecular geometry: how are the atoms in the molecule arranged
with respect to one another? In a diatomic molecule, only a single
molecular geometry is possible since the two atoms must lie on a
line. However, with a triatomic molecule (three atoms), there are
two possible geometries: the atoms may lie on a line, producing a
linear molecule, or not, producing a bent molecule. In molecules
with more than three atoms, there are many more possible
geometries. What geometries are actually observed? What determines
which geometry will be observed in a particular molecule? We seek a
model which allows us to understand the observed geometries of
molecules and thus to predict these geometries.

Once we have developed an understanding of the
relationship between molecular structure and chemical bonding, we
can attempt an understanding of the relationship of he structure
and bonding in a polyatomic molecule to the physical and chemical
properties we observe for those molecules.

Observation 1: Geometries of molecules

The geometry of a molecule includes a
description of the arrangements of the atoms in the molecule. At a
simple level, the molecular structure tells us which atoms are
bonded to which. At a more detailed level, the geometry includes
the lengths of all of these bonds, that is, the distances between
the atoms which are bonded together, and the angles between pairs
of bonds. For example, we find that in water,
H2OH2O,
the two hydrogens are bonded to the oxygen and each O-H bond length
is 95.72pm (where
1pm=10-12m1pm10-12m).
Furthermore,
H2OH2O
is a bent molecule, with the H-O-H angle equal to 104.5°.
(The measurement of these geometric properties is difficult,
involving the measurement of the frequencies at which the molecule
rotates in the gas phase. In molecules in crystalline form, the
geometry of the molecule is revealed by irradiating the crystal
with x-rays and analyzing the patterns formed as the x-rays
diffract off of the crystal.)

Not all triatomic molecules are bent, however.
As a common example,
CO2CO2
is a linear molecule. Larger polyatomics can have a variety of
shapes, as illustrated in Figure 1.
Ammonia,
NH3NH3,
is a pyramid-shaped molecule, with the hydrogens in an equilateral
triangle, the nitrogen above the plane of this triangle, and a
H-N-H angle equal to 107°. The geometry of
CH4CH4
is that of a tetrahedron, with all H-C-H angles equal to
109.5°. (See also Figure 2(a).)
Ethane,
C2H6C2H6,
has a geometry related to that of methane. The two carbons are
bonded together, and each is bonded to three hydrogens. Each H-C-H
angle is 109.5° and each H-C-C angle is
109.5°. By contrast, in ethene,
C2H4C2H4,
each H-C-H bond angle is 116.6° and each H-C-C bond angle is
121.7°. All six atoms of ethene lie in the same plane. Thus,
ethene and ethane have very different geometries, despite the
similarities in their molecular formulae.
Figure 1Molecular Structures

Molecular Structures (fig1.png)

We begin our analysis of these geometries by
noting that, in the molecules listed above which do
not contain double or triple bonds
(H2OH2O,
NH3NH3,
CH4CH4and
C2H6C2H6),
the bond angles are very similar, each equal to or very close to
the tetrahedral angle 109.5°. To account for the observed
angle, we begin with our valence shell electron pair sharing model,
and we note that, in the Lewis structures of these molecules, the
central atom in each bond angle of these molecules contains four pairs
of valence shell electrons. For methane and ethane, these four
electron pairs are all shared with adjacent bonded atoms, whereas
in

ammonia or water, one or two (respectively) of
the electron pairs are not shared with any other atom. These
unshared electron pairs are called
lone pairs . Notice that, in the two molecules with no
lone pairs, all bond angles are
exactly equal to the tetrahedral angle, whereas
the bond angles are only close in the molecules with lone
pairs

One way to understand this result is based on
the mutual repulsion of the negative charges on the valence shell
electrons. Although the two electrons in each bonding pair must
remain relatively close together in order to form the bond,
different pairs of electrons should arrange themselves in such a
way that the distances between the pairs are as large as possible.
Focusing for the moment on methane, the four pairs of electrons
must be equivalent to one another, since the four C-H bonds are
equivalent, so we can assume that the electron pairs are all the
same distance from the central carbon atom. How can we position
four electron pairs at a fixed distance from the central atom but
as far apart from one another as possible? A little reflection
reveals that this question is equivalent to asking how to place
four points on the surface of a sphere spread out from each other
as far apart as possible. A bit of experimentation reveals that
these four points must sit at the corners of a tetrahedron, an
equilateral triangular pyramid, as may be seen in Figure 2(b). If the carbon atom is at the
center of this tetrahedron and the four electron pairs at placed at
the corners, then the hydrogen atoms also form a tetrahedron about
the carbon. This is, as illustrated in Figure 2(a), the correct geometry of a methane
molecule. The angle formed by any two corners of a tetrahedron and
the central atom is 109.5°, exactly in agreement with the
observed angle in methane. This model also works well in predicting
the bond angles in ethane.
Figure 2Tetrahedral Structure of Methane
(a) The dotted lines illustrate that the hydrogens form a tetrahedron about the carbon atom.

Figure 2(a) (fig2a.png)

(b) The same tetrahedron is formed by placing four points on a sphere as far apart from one another as possible.

Figure 2(b) (fig2b.png)

We conclude that molecular geometry is
determined by minimizing the mutual repulsion of the valence shell
electron pairs. As such, this model of molecular geometry is often
referred to as the
valence shell electron pair repulsion (VSEPR) theory .
For reasons that will become clear, extension of this model implies
that a better name is the
Electron Domain (ED) Theory .

This model also accounts, at least
approximately, for the bond angles of
H2OH2O
and
NH3NH3.
These molecules are clearly not tetrahedral, like
CH4CH4,
since neither contains the requisite five atoms to form the
tetrahedron. However, each molecule does contain a central atom
surrounded by four pairs of valence shell electrons. We expect from
our Electron Domain model that those four pairs should be arrayed
in a tetrahedron, without regard to whether they are bonding or
lone-pair electrons. Then attaching the hydrogens (two for oxygen,
three for nitrogen) produces a prediction of bond angles of
109.5°, very close indeed to the observed angles of
104.5° in
H2OH2O
and 107° in
NH3NH3.

Note, however, that we do not describe the
geometries of
H2OH2O
and
NH3NH3
as "tetrahedral," since the
atoms of the molecules do not form
tetrahedrons, even if the valence shell electron pairs do. (It is
worth noting that these angles are not exactly equal to
109.5°, as in methane. These deviations will be discussed
later.)

We have developed the Electron Domain model to
this point only for geometries of molecules with four pairs of
valence shell electrons. However, there are a great variety of
molecules in which atoms from Period 3 and beyond can have more
than an octet of valence electrons. We consider two such molecules
illustrated in Figure 3.
Figure 3More Molecular Structures

More Molecular Structures (fig3.png)

First,
PCl5PCl5
is a stable gaseous compound in which the five chlorine atoms are
each bonded to the phosphorous atom. Experiments reveal that the
geometry of
PCl5PCl5
is that of a
trigonal bipyramid : three of the chlorine atoms form
an equilateral triangle with the P atom in the center, and the
other two chlorine atoms are on top of and below the P atom. Thus
there must be 10 valence shell electrons around the phosphorous
atom. Hence, phosphorous exhibits what is called an
expanded valence in
PCl5PCl5.
Applying our Electron Domain model, we expect the five valence
shell electron pairs to spread out optimally to minimize their
repulsions. The required geometry can again be found by trying to
place five points on the surface of a sphere with maximum distances
amongst these points. A little experimentation reveals that this
can be achieved by placing the five points to form a trigonal
bipyramid. Hence, Electron Domain theory accounts for the geometry
of
PCl5PCl5.

Second,
SF6SF6 is
a fairly unreactive gaseous compound in which all six fluorine
atoms are bonded to the central sulfur atom. Again, it is clear
that the octet rule is violated by the sulfur atom, which must
therefore have an expanded valence. The observed geometry of
SF6SF6,
as shown in Figure 3, is highly
symmetric: all bond lengths are identical and all bond angles are
90°. The F atoms form an
octahedron about the central S atom: four of the F
atoms form a square with the S atom at the center, and the other
two F atoms are above and below the S atom. To apply our Electron
Domain model to understand this geometry, we must place six points,
representing the six electron pairs about the central S atom, on
the surface of a sphere with maximum distances between the points.
The requisite geometry is found, in fact, to be that of an
octahedron, in agreement with the observed geometry.

As an example of a molecule with an atom with
less than an octet of valence shell electrons, we consider boron
trichloride,
BCl3BCl3.
The geometry of
BCl3BCl3 is
also given in Figure 3: it is
trigonal planar , with all four atoms lying in the same
plane, and all Cl-B-Cl bond angles equal to 120°. The three Cl
atoms form an equilateral triangle. The Boron atom has only three
pairs of valence shell electrons in
BCl3BCl3.
In applying Electron Domain theory to understand this geometry, we
must place three points on the surface of a sphere with maximum
distance between points. We find that the three points form an
equilateral triangle in a plane with the center of the sphere, so
Electron Domain is again in accord with the observed
geometry.

We conclude from these predictions and
observations that the Electron Domain model is a reasonably
accurate way to understand molecular geometries, even in molecules
which violate the octet rule.

Observation 2: Molecules with Double or Triple Bonds

In each of the molecules considered up to this
point, the electron pairs are either in single bonds or in lone
pairs. In current form, the Electron Domain model does
not account for the observed geometry of
C2H4C2H4,
in which each H-C-H bond angle is 116.6° and each H-C-C bond
angle is 121.7° and all six atoms lie in the same plane.
Each carbon atom in this molecule is surrounded by four pairs of
electrons, all of which are involved in bonding,
i.e. there are no lone pairs. However, the
arrangement of these electron pairs, and thus the bonded atoms,
about each carbon is not even approximately tetrahedral. Rather,
the H-C-H and H-C-C bond angles are much closer to 120°, the
angle which would be expected if
three electron pairs were separated in the
optimal arrangement, as just discussed for
BCl3BCl3.

This observed geometry can be understood by
re-examining the Lewis structure. Recall that, although there are
four electron pairs about each carbon atom, two of these pairs form
a double bond between the carbon atoms. It is tempting to assume
that these four electron pairs are forced apart to form a
tetrahedron as in previous molecules. However, if this were this
case, the two pairs involved in the double bond would be separated
by an angle of 109.5° which would make it impossible for
both pairs to be localized between the carbon atoms. To preserve
the double bond, we must assume that the two electron pairs in the
double bond remain in the same vicinity. Given this assumption,
separating the three
independent groups of electron pairs about a
carbon atom produces an expectation that all three pairs should lie
in the same plane as the carbon atom, separated by 120°
angles. This agrees very closely with the observed bond angles. We
conclude that the our model can be extended to understanding the
geometries of molecules with double (or triple) bonds by treating
the multiple bond as two electron pairs confined to a single
domain. It is for this reason that we refer to the
model as Electron Domain theory.

Applied in this form, Electron Domain theory
can help us understand the linear geometry of
CO2CO2.
Again, there are four electron pairs in the valence shell of the
carbon atom, but these are grouped into only two domains of two
electron pairs each, corresponding to the two C=O double bonds.
Minimizing the repulsion between these two domains forces the
oxygen atoms to directly opposite sides of the carbon, producing a
linear molecule. Similar reasoning using Electron Domain theory as
applied to triple bonds correctly predicts that acetylene,
HCCHHCCH,
is a linear molecule. If the electron pairs in the triple bond are
treated as a single domain, then each carbon atom has only two
domains each. Forcing these domains to opposite sides from one
another accurately predicts 180° H-C-C bond angles.

Observation 3: Distortions from Expected Geometries

It is interesting to note that some molecular
geometries
(CH4CH4,
CO2CO2,
HCCHHCCH)
are exactly predicted by the Electron Domain model, whereas in
other molecules, the model predictions are only approximately
correct. For examples, the observed angles in ammonia and water
each differ slightly from the tetrahedral angle. Here again, there
are four pairs of valence shell electrons about the central atoms.
As such, it is reasonable to conclude that the bond angles are
determined by the mutual repulsion of these electron pairs, and are
thus expected to be 109.5°, which is close but not
exact.

One clue as to a possible reason for the
discrepancy is that the bond angles in ammonia and water are both
less than 109.5°. Another is that both
ammonia and water molecules have lone pair electrons, whereas there
are no lone pairs in a methane molecule, for which the Electron
Domain prediction is exact. Moreover, the bond angle in water, with
two lone pairs, is less than the bond angles in ammonia, with a
single lone pair. We can straightforwardly conclude from these
observations that the lone pairs of electrons must produce a
greater repulsive effect than do the bonded pairs. Thus, in
ammonia, the three bonded pairs of electrons are forced together
slightly compared to those in methane, due to the greater repulsive
effect of the lone pair. Likewise, in water, the two bonded pairs
of electrons are even further forced together by the two lone pairs
of electrons.

This model accounts for the comparative bond
angles observed experimentally in these molecules. The valence
shell electron pairs repel one another, establishing the geometry
in which the energy of their interaction is minimized. Lone pair
electrons apparently generate a greater repulsion, thus slightly
reducing the angles between the bonded pairs of electrons. Although
this model accounts for the observed geometries, why should lone
pair electrons generate a greater repulsive effect? We must guess
at a qualitative answer to this question, since we have no
description at this point for where the valence shell electron
pairs actually are or what it means to share an electron pair. We
can assume, however, that a pair of electrons shared by two atoms
must be located somewhere between the two nuclei, otherwise our
concept of "sharing" is quite meaningless. Therefore, the powerful
tendency of the two electrons in the pair to repel one another must
be significantly offset by the localization of these electrons
between the two nuclei which share them. By contrast, a lone pair
of electrons need not be so localized, since there is no second
nucleus to draw them into the same vicinity. Thus more free to move
about the central atom, these lone pair electrons must have a more
significant repulsive effect on the other pairs of
electrons.

These ideas can be extended by more closely
examining the geometry of ethene,
C2H4C2H4
. Recall that each H-C-H bond angle is 116.6° and each H-C-C
bond angle is 121.7°, whereas the Electron Domain theory
prediction is for bond angles exactly equal to 120°. We can
understand why the H-C-H bond angle is slightly less than
120° by assuming that the two pairs of electrons in the C=C
double bond produce a greater repulsive effect than do either of
the single pairs of electrons in the C-H single bonds. The result
of this greater repulsion is a slight "pinching" of the H-C-H bond
angle to less than 120°.

The concept that lone pair electrons produce a
greater repulsive effect than do bonded pairs can be used to
understand other interesting molecular geometries. Sulfur
tetrafluoride,
SF4SF4,
is a particularly interesting example, shown in Figure 4.
Figure 4Molecular Structure of SF4

Molecular Structure of SF4 (fig4.png)

Note that two of the fluorines form close to a
straight line with the central sulfur atom, but the other two are
approximately perpendicular to the first two and at an angle of
101.5° to each other. Viewed sideways, this structure
looks something like a seesaw.

To account for this structure, we first
prepare a Lewis structure. We find that each fluorine atom is
singly bonded to the sulfur atom, and that there is a lone pair of
electrons on the sulfur. Thus, with five electron pairs around the
central atom, we expect the electrons to arrange themselves in a
trigonal bipyramid, similar to the arrangement in
PCl5PCl5
in Figure 3. In this case, however,
the fluorine atoms and the lone pair could be arranged in two
different ways with two different resultant molecular structures.
The lone pair can either go on the axis of the trigonal
bipyramid (i.e. “above” the sulfur) or on the
equator of the bipyramid (i.e. “beside” the sulfur).

The actual molecular structure in Figure 4 shows clearly that the lone pair
goes on the equatorial position. This can be understood if we
assume that the lone pair produces a greater repulsive effect than
do the bonded pairs. With this assumption, we can deduce that the
lone pair should be placed in the trigonal bipyramidal arrangement
as far as possible from the bonded pairs. The equatorial position
does a better job of this, since only two bonding pairs of
electrons are at approximately 90° angles from the
lone pair in this position. By contrast, a lone pair in the axial
position is approximately 90° away from three bonding
pairs. Therefore, our Electron Domain model assumptions are
consistent with the observed geometry of
SF4SF4.
Note that these assumptions also correctly predict the observed
distortions away from the 180° and
120° angles which would be predicted by a trigonal
bipyramidal arrangement of the five electron pairs.

Review and Discussion Questions
Exercise 1

Using a styrofoam or rubber ball, prove to
yourself that a tetrahedral arrangement provides the maximum
separation of four points on the surface of the ball. Repeat this
argument to find the expected arrangements for two, three, five,
and six points on the surface of the ball.
Exercise 2

Explain why arranging points on the surface of
a sphere can be considered equivalent to arranging electron pairs
about a central atom.
Exercise 3

The valence shell electron pairs about the
central atom in each of the molecules
H2OH2O,
NH3NH3,
and
CH4CH4 are
arranged approximately in a tetrahedron. However, only
CH4CH4 is
considered a tetrahedral molecule. Explain why these statements are
not inconsistent.
Exercise 4

Explain how a comparison of the geometries of
H2OH2O
and
CH4CH4 leads
to a conclusion that lone pair electrons produce a greater
repulsive effect than do bonded pairs of electrons. Give a physical
reason why this might be expected.
Exercise 5

Explain why the octet of electrons about each
carbon atom in ethene,
C2H4C2H4,
are not arranged even approximately in a tetrahedron.
Exercise 6

Assess the accuracy of the following reasoning
and conclusions:

A trigonal bipyramid forms when there are five
electron domains. If one ED is a lone pair, then the lone pair
takes an equatorial position and the molecule has a seesaw
geometry. If two EDs are lone pairs, we have to decide among the
following options: both axial, both equatorial, or one axial and
one equatorial. By placing both lone pairs in the axial positions,
the lone pairs are as far apart as possible, so the trigonal planar
structure is favored.

Exercise 7

Assess the accuracy of the following reasoning
and conclusions:

The Cl-X-Cl bond angles in the two molecules are
identical, because the bond angle is determined by the repulsion of
the two Cl atoms, which is identical in the two molecules.

Figure 5
Figure 5 (fig5.png)

From: http://cnx.org/content/m12594/latest/

Nuclear Chemistry

2009-02-02T19:14:00.000-08:00

Nuclear Chemistry
An Introduction

by Anthony Carpi, Ph.D.

Traditional chemical reactions occur as a result of the interaction between valence electrons around an atom's nucleus (see our Chemical Reactions module for more information). In 1896, Henri Becquerel expanded the field of chemistry to include nuclear changes when he discovered that uranium emitted radiation. Soon after Becquerel's discovery, Marie Sklodowska Curie began studying radioactivity and completed much of the pioneering work on nuclear changes. Curie found that radiation was proportional to the amount of radioactive element present, and she proposed that radiation was a property of atoms (as opposed to a chemical property of a compound). Marie Curie was the first woman to win a Nobel Prize and the first person to win two (the first, shared with her husband Pierre and Becquerel for discovering radioactivity; the second for discovering the radioactive elements radium and polonium).

Radiation and Nuclear Reactions
In 1902, Frederick Soddy proposed the theory that "radioactivity is the result of a natural change of an isotope of one element into an isotope of a different element." Nuclear reactions involve changes in particles in an atom's nucleus and thus cause a change in the atom itself. All elements heavier than bismuth (Bi) (and some lighter) exhibit natural radioactivity and thus can "decay" into lighter elements. Unlike normal chemical reactions that form molecules, nuclear reactions result in the transmutation of one element into a different isotope or a different element altogether (remember that the number of protons in an atom defines the element, so a change in protons results in a change in the atom). There are three common types of radiation and nuclear changes:

1. Alpha Radiation (α) is the emission of an alpha particle from an atom's nucleus. An α particle contains two protons and two neutrons (and is similar to a He nucleus: particle-alpha ). When an atom emits an a particle, the atom's atomic mass will decrease by four units (because two protons and two neutrons are lost) and the atomic number (z) will decrease by two units. The element is said to "transmute" into another element that is two z units smaller. An example of an a transmutation takes place when uranium decays into the element thorium (Th) by emitting an alpha particle, as depicted in the following equation:

(Note: in nuclear chemistry, element symbols are traditionally preceded by their atomic weight (upper left) and atomic number (lower left).
2. Beta Radiation (β) is the transmutation of a neutron into a proton and a electron (followed by the emission of the electron from the atom's nucleus: particle-beta ). When an atom emits a β particle, the atom's mass will not change (since there is no change in the total number of nuclear particles), however the atomic number will increase by one (because the neutron transmutated into an additional proton). An example of this is the decay of the isotope of carbon named carbon-14 into the element nitrogen:

3. Gamma Radiation (γ) involves the emission of electromagnetic energy (similar to light energy) from an atom's nucleus. No particles are emitted during gamma radiation, and thus gamma radiation does not itself cause the transmutation of atoms, however γ radiation is often emitted during, and simultaneous to, α or β radioactive decay. X-rays, emitted during the beta decay of cobalt-60, are a common example of gamma radiation.

Half-Life

Radioactive decay proceeds according to a principal called the half-life. The half-life (T½) is the amount of time necessary for one-half of the radioactive material to decay. For example, the radioactive element bismuth (210Bi) can undergo alpha decay to form the element thallium (206Tl) with a reaction half-life equal to five days. If we begin an experiment starting with 100 g of bismuth in a sealed lead container, after five days we will have 50 g of bismuth and 50 g of thallium in the jar. After another five days (ten from the starting point), one-half of the remaining bismuth will decay and we will be left with 25 g of bismuth and 75 g of thallium in the jar. As illustrated, the reaction proceeds in halfs, with half of whatever is left of the radioactive element decaying every half-life period.
decay-graph - Radioactive Decay of Bismuth-210 (T½ = 5 days)

Radioactive Decay of Bismuth-210 (T½ = 5 days)

The fraction of parent material that remains after radioactive decay can be calculated using the equation:
Fraction remaining = 1
2n (where n = # half-lives elapsed)

The amount of a radioactive material that remains after a given number of half-lives is therefore:
Amount remaining = Original amount * Fraction remaining

The decay reaction and T½ of a substance are specific to the isotope of the element undergoing radioactive decay. For example, Bi210 can undergo a decay to Tl206 with a T½ of five days. Bi215, by comparison, undergoes b decay to Po215 with a T½ of 7.6 minutes, and Bi208 undergoes yet another mode of radioactive decay (called electron capture) with a T½ of 368,000 years!

Stimulated Nuclear Reactions
While many elements undergo radioactive decay naturally, nuclear reactions can also be stimulated artificially. Although these reactions also occur naturally, we are most familiar with them as stimulated reactions. There are two such types of nuclear reactions:

1. Nuclear fission: reactions in which an atom's nucleus splits into smaller parts, releasing a large amount of energy in the process. Most commonly this is done by "firing" a neutron at the nucleus of an atom. The energy of the neutron "bullet" causes the target element to split into two (or more) elements that are lighter than the parent atom.
fission reaction - The Fission Reaction of Uranium-235

The Fission Reaction of Uranium-235

The Fission of U235

During the fission of U235, three neutrons are released in addition to the two daughter atoms. If these released neutrons collide with nearby U235 nuclei, they can stimulate the fission of these atoms and start a self-sustaining nuclear chain reaction. This chain reaction is the basis of nuclear power. As uranium atoms continue to split, a significant amount of energy is released from the reaction. The heat released during this reaction is harvested and used to generate electrical energy.

Two Types of Nuclear Chain Reactions

2. Nuclear fusion: reactions in which two or more elements "fuse" together to form one larger element, releasing energy in the process. A good example is the fusion of two "heavy" isotopes of hydrogen (deuterium: H2 and tritium: H3) into the element helium.
fusion reaction - Nuclear Fusion of Two Hydrogen Isotopes

Nuclear Fusion of Two Hydrogen Isotopes

Nuclear Fusion

Fusion reactions release tremendous amounts of energy and are commonly referred to as thermonuclear reactions. Although many people think of the sun as a large fireball, the sun (and all stars) are actually enormous fusion reactors. Stars are primarily gigantic balls of hydrogen gas under tremendous pressure due to gravitational forces. Hydrogen molecules are fused into helium and heavier elements inside of stars, releasing energy that we receive as light and heat.

External Resources

• Crucibles: The Story of Chemistry from Ancient Alchemy to Nuclear Fission

• Principles of American Nuclear Chemistry: A Novel

Matter: States of Matter

2009-02-02T19:12:00.000-08:00

Matter: States of Matter

by Anthony Carpi, Ph.D.
water-boiling As a young boy, I remember staring in wonder at a pot of boiling water. Searching for an explanation for the bubbles that formed, I believed for a time that the motion of the hot water drew air down into the pot, which then bubbled back to the surface. Little did I know that what was happening was even more magical than I imagined - the bubbles were not air, but actually water in the form of a gas.

The different states of matter have long confused people. The ancient Greeks were the first to identify three classes (what we now call states) of matter based on their observations of water. But these same Greeks, in particular the philosopher Thales (624 - 545 b.c.), incorrectly suggested that since water could exist as a solid, liquid, or even a gas under natural conditions, it must be the single principal element in the universe from which all other substances are made. We now know that water is not the fundamental substance of the universe; in fact, it is not even an element.

To understand the different states in which matter can exist, we need to understand something called the Kinetic Molecular Theory of Matter. Kinetic Molecular Theory has many parts, but we will introduce just a few here. One of the basic concepts of the theory states that atoms and molecules possess an energy of motion that we perceive as temperature. In other words, atoms and molecules are constantly moving, and we measure the energy of these movements as the temperature of the substance. The more energy a substance has, the more molecular movement there will be, and the higher the perceived temperature will be. An important point that follows this is that the amount of energy that atoms and molecules have (and thus the amount of movement) influences their interaction with each other. Unlike simple billiard balls, many atoms and molecules are attracted to each other as a result of various intermolecular forces such as hydrogen bonds, van der Waals forces, and others. Atoms and molecules that have relatively small amounts of energy (and movement) will interact strongly with each other, while those that have relatively high energy will interact only slightly, if even at all, with others.

How does this produce different states of matter? Atoms that have low energy interact strongly and tend to “lock” in place with respect to other atoms. Thus, collectively, these atoms form a hard substance, what we call a solid. Atoms that possess high energy will move past each other freely, flying about a room, and forming what we call a gas. As it turns out, there are several known states of matter; a few of them are detailed below.
ice-cubes Solids are formed when the attractive forces between individual molecules are greater than the energy causing them to move apart. Individual molecules are locked in position near each other, and cannot move past one another. The atoms or molecules of solids remain in motion. However, that motion is limited to vibrational energy; individual molecules stay fixed in place and vibrate next to each other. As the temperature of a solid is increased, the amount of vibration increases, but the solid retains its shape and volume because the molecules are locked in place relative to each other. To view an example of this, click on the animation below which shows the molecular structure of ice crystals.

Solid matter - ice

water-liquid Liquids are formed when the energy (usually in the form of heat) of a system is increased and the rigid structure of the solid state is broken down. In liquids, molecules can move past one another and bump into other molecules; however, they remain relatively close to each other like solids. Often in liquids, intermolecular forces (such as the hydrogen bonds shown in the animation below) pull molecules together and are quickly broken. As the temperature of a liquid is increased, the amount of movement of individual molecules increases. As a result, liquids can “flow” to take the shape of their container but they cannot be easily compressed because the molecules are already close together. Thus liquids have an undefined shape, but a defined volume. In the example animation below we see that liquid water is made up of molecules that can freely move past one another, yet remain relatively close in distance to each other.

Liquid matter - water

A simulation of molecular movement within liquid water.

(Flash required)
clouds Gases are formed when the energy in the system exceeds all of the attractive forces between molecules. Thus gas molecules have little interaction with each other beyond occasionally bumping into one another. In the gas state, molecules move quickly and are free to move in any direction, spreading out long distances. As the temperature of a gas increases, the amount of movement of individual molecules increases. Gases expand to fill their containers and have low density. Because individual molecules are widely separated and can move around easily in the gas state, gases can be compressed easily and they have an undefined shape.

Gaseous matter - steam

Solids, liquids, and gases are the most common states of matter that exist on our planet. If you would like to compare the three states to one another, click on the comparison animation below. Note the differences in molecular motion of water molecules in these three states.

Solid-Liquid-Gas Comparison

sun Plasmas are hot, ionized gases. Plasmas are formed under conditions of extremely high energy, so high, in fact, that molecules are ripped apart and only free atoms exist. More astounding, plasmas have so much energy that the outer electrons are actually ripped off of individual atoms, thus forming a gas of highly energetic, charged ions. Because the atoms in plasma exist as charged ions, plasmas behave differently than gases, thus representing a fourth state of matter. Plasmas can be commonly seen simply by looking upward; the high energy conditions that exist in stars such as our sun force individual atoms into the plasma state.

As we have seen, increasing energy leads to more molecular motion. Conversely, decreasing energy results in less molecular motion. As a result, one prediction of Kinetic Molecular Theory is that if we continue to decrease the energy (measured as temperature) of a substance, we will reach a point at which all molecular motion stops. The temperature at which molecular motion stops is called absolute zero and has been calculated to be -273.15 degrees Celsius. While scientists have cooled substances to temperatures close to absolute zero, they have never actually reached absolute zero. The difficulty with observing a substance at absolute zero is that to “see” the substance, light is needed, and light itself transfers energy to the substance, thus raising the temperature. Despite these challenges, scientists have recently observed a fifth state of matter that only exists at temperatures very close to absolute zero.

Bose-Einstein Condensates represent a fifth state of matter only seen for the first time in 1995. The state is named after Satyendra Nath Bose and Albert Einstein who predicted its existence in the 1920’s. B-E condensates are gaseous superfluids cooled to temperatures very near absolute zero. In this weird state, all the atoms of the condensate attain the same quantum-mechanical state and can flow past one another without friction. Even more strangely, B-E condensates can actually “trap” light, releasing it when the state breaks down.

Several other less common states of matter have also either been described or actually seen. Some of these states include liquid crystals, fermionic condensates, superfluids, supersolids and the aptly named strange matter. To read more about these phases, visit the Phase page of the Wikipedia linked to below in the Further Exploration section.
water Phase Transitions
The transformation of one state of matter into another state is called a phase transition. The more common phase transitions even have names; for example, the terms melting and freezing describe phase transitions between the solid and liquid state, and the terms evaporation and condensation describe transitions between the liquid and gas state. Phase transitions occur at very precise points, when the energy (measured as temperature) of a substance in a given state exceeds that allowed in the state. For example, liquid water can exist at a range of temperatures. Cold drinking water may be around 4ºC. Hot shower water has more energy and thus may be around 40ºC. However, at 100°C under normal conditions, water will begin to undergo a phase transition into the gas phase. At this point, energy introduced into the liquid will not go into increasing the temperature; it will be used to send molecules of water into the gas state. Thus, no matter how high the flame is on the stove, a pot of boiling water will remain at 100ºC until all of the water has undergone transition to the gas phase. The excess energy introduced by a high flame will accelerate the liquid-to-gas transition; it will not change the temperature. The heat curve below illustrates the corresponding changes in energy (shown in calories) and temperature of water as it undergoes a phase transition between the liquid and gas states.
graph2 - heat curve

As can be seen in the graph above, as we move from left to right, the temperature of liquid water increases as energy (heat) is introduced. At 100ºC, water begins to undergo a phase transition and the temperature remains constant even as energy is added (the flat part of the graph). The energy that is introduced during this period goes toward breaking intermolecular forces so that individual water molecules can “escape” into the gas state. Finally, once the transition is complete, if further energy is added to the system, the heat of the gaseous water, or steam, will increase.

This same process can be seen in reverse if we simply look at the graph above starting on the right side and moving left. As steam is cooled, the movement of gaseous water molecules and thus temperature will decrease. When the gas reaches 100ºC, more energy will be lost from the system as the attractive forces between molecules reform; however the temperature remains constant during the transition (the flat part of the graph). Finally, when condensation is complete, the temperature of the liquid will begin to fall as energy is withdrawn.

Phase transitions are an important part of the world around us. For example, the energy withdrawn when perspiration evaporates from the surface of your skin allows your body to correctly regulate its temperature during hot days. Phase transitions play an important part in geology, influencing mineral formation and possibly even earthquakes. And who can ignore the phase transition that occurs at about -3ºC, when cream, perhaps with a few strawberries or chocolate chunks, begins to form solid ice cream.

Now we understand what is happening in a pot of boiling water. The energy (heat) introduced at the bottom of the pot causes a localized phase transition of liquid water to the gaseous state. Because gases are less dense than liquids, these localized phase transitions form pockets (or bubbles) of gas, which rise to the surface of the pot and burst. But nature is often more magical than our imaginations. Despite all that we know about the states of matter and phase transitions, we still cannot predict where the individual bubbles will form in a pot of boiling water.

External Resources

• Absolute Zero and the Conquest of Cold

• Inventing Temperature: Measurement and Scientific Progress (Oxford Studies in the Philosophy of Science)

Matter

2009-02-02T19:09:00.000-08:00

Matter
Atoms from Democritus to Dalton

by Anthony Carpi, Ph.D.
elements of matter Early humans easily distinguished between materials that were used for making clothes, shaping into tools, or good to eat, and they developed a language of words to describe these things, such as “fur,” “stone,” or “rabbit.” However, these people did not have our current understanding of the substances that made up those objects. Empedocles, a Greek philosopher and scientist who lived on the south coast of Sicily between 492 b.c. and 432 b.c., proposed one of the first theories that attempted to describe the things around us. Empedocles argued that all matter was composed of four elements: fire, air, water, and earth. The ratio of these four elements affected the properties of the matter. Stone was thought to contain a high amount of earth, while a rabbit was thought to have a higher ratio of both water and fire, thus making it soft and giving it life. Empedocles’s theory was quite popular, but it had a number of problems. For example, no matter how many times you break a stone in half, the pieces never resemble any of the core elements of fire, air, water, or earth. Despite these problems, Empedocles’s theory was an important development in scientific thinking because it was among the first to suggest that some substances that looked like pure materials, like stone, were actually made up of a combination of different "elements."
chemical reaction - ancient

A few decades after Empedocles, Democritus, another Greek who lived from 460 b.c. to 370 b.c., developed a new theory of matter that attempted to overcome the problems of his predecessor. Democritus’s ideas were based on reasoning rather than science, and drew on the teachings of two Greek philosophers who came before him: Leucippus and Anaxagoras. Democritus knew that if you took a stone and cut it in half, each half had the same properties as the original stone. He reasoned that if you continued to cut the stone into smaller and smaller pieces, at some point you would reach a piece so tiny that it could no longer be divided. Democritus called these infinitesimally small pieces of matter atomos, meaning "indivisible." He suggested that atomos were eternal and could not be destroyed. Democritus theorized that atomos were specific to the material that they made up, meaning that the atomos of stone were unique to stone and different from the atomos of other materials, such as fur. This was a remarkable theory that attempted to explain the whole physical world in terms of a small number of ideas.

Ultimately, though, Aristotle and Plato, two of the best-known philosophers of Ancient Greece, rejected the theories of Democritus. Aristotle accepted the theory of Empedocles, adding his own (incorrect) idea that the four core elements could be transformed into one another. Because of Aristotle’s great influence, Democritus’s theory would have to wait almost 2,000 years before being rediscovered.

In the seventeenth and eighteenth centuries a.d., several key events helped revive the theory that matter was made of small, indivisible particles. In 1643, Evangelista Torricelli, an Italian mathematician and pupil of Galileo, showed that air had weight and was capable of pushing down on a column of liquid mercury (thus inventing the barometer). This was a startling finding. If air - this substance that we could not see, feel, or smell - had weight, it must be made of something physical. But how could something have a physical presence, yet not respond to human touch or sight? Daniel Bernoulli, a Swiss mathematician, proposed an answer. He developed a theory that air and other gases consist of tiny particles that are too small to be seen, and are loosely packed in an empty volume of space. The particles could not be felt because unlike a solid stone wall that does not move, the tiny particles move aside when a human hand or body moves through them. Bernoulli reasoned that if these particles were not in constant motion they would settle to the ground like dust particles; therefore he pictured air and other gases as loose collections of tiny billiard-ball-like particles that are continuously moving around and bouncing off one another.
cinnabar - 3D Many scientists were busy studying the natural world at this time. Shortly after Bernoulli proposed his theory, the Englishman Joseph Priestley began to experiment with red mercury calx in 1773. Mercury calx, a red solid stone, had been known and coveted for thousands of years because when it is heated, it appears to turn into mercury, a silver liquid metal. Priestley had observed that it does not just turn into mercury; it actually breaks down into two substances when it is heated, liquid mercury and a strange gas. Priestley carefully collected this gas in glass jars and studied it. After many long days and nights in the laboratory, Priestley said of the strange gas, “what surprised me more than I can well express was that a candle burned in this air with a remarkably vigorous flame.” Not only did flames burn strongly in this gas, but a mouse placed in a sealed container of this gas lived for a longer period of time than a mouse placed in a sealed container of ordinary air. Priestley’s discovery revealed that substances could combine together or break apart to form new substances with different properties. For example, a colorless, odorless gas could combine with mercury, a silver metal, to form mercury calx, a red mineral.

Priestley called the gas he discovered dephlogisticated air, but this name would not stick. In 1778, Antoine Lavoisier, a French scientist, conducted many experiments with dephlogisticated air and theorized that the gas made some substances acidic. He renamed Priestley’s gas oxygen, from the Greek words that loosely translate as "acid maker". While Lavoisier’s theory about oxygen and acids proved incorrect, his name stuck. Lavoisier knew from other scientists before him that acids react with some metals to release another strange and highly flammable gas called phlogiston. Lavoisier mixed the two gases, phlogiston and the newly renamed oxygen, in a closed glass container and inserted a match. He saw that phlogiston immediately burned in the presence of oxygen and afterwards he observed droplets of water on the glass container. After careful testing, Lavoisier realized that the water was formed by the reaction of phlogiston and oxygen, and so he renamed phlogiston hydrogen, from the Greek words for "water maker". Lavoisier also burned other substances such as phosphorus and sulfur in air, and showed that they combined with air to make new materials. These new materials weighed more than the original substances, and Lavoisier showed that the weight gained by the new materials was lost from the air in which the substances were burned. From these observations, Lavoisier established the Law of Conservation of Mass, which says that mass is not lost or gained during a chemical reaction.
priestleys apparatus - An eighteenth-century chemistry bench.

An eighteenth-century chemistry bench.

Priestley, Lavoisier, and others had laid the foundations of the field of chemistry. Their experiments showed that some substances could combine with others to form new materials; other substances could be broken apart to form simpler ones; and a few key “elements” could not be broken down any further. But what could explain this complex set of observations? John Dalton, an exceptional British teacher and scientist, put together the pieces and developed the first modern atomic theory in 1803. To learn more about Priestley's and Lavoisier's experiments and how they formed the basis of Dalton's theories, try the interactive experiment Dalton's Playhouse, linked to below.

Dalton's Playhouse

An interactive, virtual set of experiments that allow you to recreate classic experiments from the nineteenth century.

Dalton made it a regular habit to track and record the weather in his home town of Manchester, England. Through his observations of morning fog and other weather patterns, Dalton realized that water could exist as a gas that mixed with air and occupied the same space as air. Solids could not occupy the same space as each other; for example, ice could not mix with air. So what could allow water to sometimes behave as a solid and sometimes as a gas? Dalton realized that all matter must be composed of tiny particles. In the gas state, those particles floated freely around and could mix with other gases, as Bernoulli had proposed. But Dalton extended this idea to apply to all matter – gases, solids and liquids. Dalton first proposed part of his atomic theory in 1803 and later refined these concepts in his classic 1808 paper A New System of Chemical Philosophy (which you can access through
Dalton's theory had four main concepts:

1. All matter is composed of indivisible particles called atoms. Bernoulli, Dalton, and others pictured atoms as tiny billiard-ball-like particles in various states of motion. While this concept is useful to help us understand atoms, it is not correct as we will see in later modules on atomic theory linked to at the bottom of this module.
2. All atoms of a given element are identical; atoms of different elements have different properties. Dalton’s theory suggested that every single atom of an element such as oxygen is identical to every other oxygen atom; furthermore, atoms of different elements, such as oxygen and mercury, are different from each other. Dalton characterized elements according to their atomic weight; however, when isotopes of elements were discovered in the late 1800s this concept changed.
3. Chemical reactions involve the combination of atoms, not the destruction of atoms. Atoms are indestructible and unchangeable, so compounds, such as water and mercury calx, are formed when one atom chemically combines with other atoms. This was an extremely advanced concept for its time; while Dalton’s theory implied that atoms bonded together, it would be more than 100 years before scientists began to explain the concept of chemical bonding.
4. When elements react to form compounds, they react in defined, whole-number ratios. The experiments that Dalton and others performed showed that reactions are not random events; they proceed according to precise and well-defined formulas. This important concept in chemistry is discussed in more detail below.

Some of the details of Dalton’s atomic theory require more explanation.

Elements: As early as 1660, Robert Boyle recognized that the Greek definition of element (earth, fire, air, and water) was not correct. Boyle proposed a new definition of an element as a fundamental substance, and we now define elements as fundamental substances that cannot be broken down further by chemical means. Elements are the building blocks of the universe. They are pure substances that form the basis of all of the materials around us. Some elements can be seen in pure form, such as mercury in a thermometer; some we see mainly in chemical combination with others, such as oxygen and hydrogen in water. We now know of approximately 116 different elements. Each of the elements is given a name and a one- or two-letter abbreviation. Often this abbreviation is simply the first letter of the element; for example, hydrogen is abbreviated as H, and oxygen as O. Sometimes an element is given a two-letter abbreviation; for example, helium is He. When writing the abbreviation for an element, the first letter is always capitalized and the second letter (if there is one) is always lowercase.

Atoms: A single unit of an element is called an atom. The atom is the most basic unit of the matter that makes up everything in the world around us. Each atom retains all of the chemical and physical properties of its parent element. At the end of the nineteenth century, scientists would show that atoms were actually made up of smaller, "subatomic" pieces, which smashed the billiard-ball concept of the atom (see our Atomic Theory I: The Early Days module).
water molecule-with hooks Compounds: Most of the materials we come into contact with are compounds, substances formed by the chemical combination of two or more atoms of the elements. A single “particle” of a compound is called a molecule. Dalton incorrectly imagined that atoms “hooked” together to form molecules. However, Dalton correctly realized that compounds have precise formulas. Water, for example, is always made up of two parts hydrogen and one part oxygen. The chemical formula of a compound is written by listing the symbols of the elements together, without any spaces between them. If a molecule contains more than one atom of an element, a number is subscripted after the symbol to show the number of atoms of that element in the molecule. Thus the formula for water is H2O, never HO or H2O2.

The idea that compounds have defined chemical formulas was first proposed in the late 1700s by the French chemist Joseph Proust. Proust performed a number of experiments and observed that no matter how he caused different elements to react with oxygen, they always reacted in defined proportions. For example, two parts of hydrogen always reacts with one part oxygen when forming water; one part mercury always reacts with one part oxygen when forming mercury calx. Dalton used Proust’s Law of Definite Proportions in developing his atomic theory.
balloon - definite proportions

The law also applies to multiples of the fundamental proportion, for example:
balloon - multiple proportions

In both of these examples, the ratio of hydrogen to oxygen to water is 2 to 1 to 1. When reactants are present in excess of the fundamental proportions, some reactants will remain unchanged after the chemical reaction has occurred.
balloon - excess reactant

The story of the development of modern atomic theory is one in which scientists built upon the work of others to produce a more accurate explanation of the world around them. This process is common in science, and even incorrect theories can contribute to important scientific discoveries. Dalton, Priestley, and others laid the foundation of atomic theory, and many of their hypotheses are still useful. However, in the decades after their work, other scientists would show that atoms are not solid billiard balls, but complex systems of particles. Thus they would smash apart a bit of Dalton’s atomic theory in an effort to build a more complete view of the world around us.

External Resources

• Memoir of John Dalton, and History of the Atomic Theory up to His Time

• The Development of Modern Chemistry

• Other Recommended Products

Fats and Proteins

2009-02-02T19:08:00.000-08:00

Fats and Proteins

by Anthony Carpi, Ph.D.

In addition to the carbohydrates, fats and proteins are the other two macronutrients required by the human body (see our Carbohydrates module).

Fats
Fats are a subgroup of compounds known as lipids that are found in the body and have the general property of being hydrophobic (meaning they are insoluble in water). Fats are also known as triglycerides, molecules made from the combination of one molecule of glycerol with three fatty acids, as depicted below:

The main purpose of fats in the body is to serve as a storage system and reserve supply of energy. During periods of low food consumption, fat reserves in the body can be mobilized and broken down to release energy. Fats serve as an insulation material to allow body heat to be conserved and fats line and protect delicate internal organs from physical damage. Fats in the diet can be converted to other lipids that serve as the main structural material in the membranes surrounding our cells. Fats are also used in the manufacture of some steroids and hormones that help regulate proper growth and maintenance of tissue in the body.

Fats can be classified as either saturated or unsaturated depending on the structure of the long carbon-carbon chains in the fatty acids (the R's in the diagram above). Fats that contain no double bonds in their fatty acid chains are referred to as saturated fats. These fats tend to be solid at room temperature, such as butter or animal fat. The consumption of saturated fats carries some health risks in that they have been linked to arteriosclerosis (hardening of the arteries) and heart disease. Unsaturated fats contain some number of double bonds in their structure. These fats are generally liquid at room temperature (fats that are liquid at room temperature are referred to as oils). Unsaturated fats can be either polyunsaturated (many double bonds) or monounsaturated fats (one or few double bonds). Recent research suggests that the healthiest of the fats in the human diet are the monounsaturated fats, such as olive oil and canola oil, because they appear to be beneficial in the fight against heart disease.

Proteins
Proteins are polymers of amino acids. Though there are hundreds of thousands of different proteins that exist in nature, they are all made up of different combinations of amino acids. Proteins are large molecules that may consist of hundreds, or even thousands, of amino acids. Amino acids all have the general structure:
aminoacid - General Structure of an Amino Acid

General Structure of an Amino Acid

The R in the diagram represents a functional group that varies depending on the specific amino acid in question. For example, R can be simply an H atom, as in the amino acid glycine, or a more complex organic group. When two amino acids bond together, the two ends of nearby amino acids (shown in red) are released and the carbon (called a carboxyl) end of one amino acid bonds to the nitrogen end of the adjacent one forming a peptide bond, as illustrated below:
peptide bond - A Peptide Bond

A Peptide Bond

When many amino acids bond together to create long chains, the structure is called a protein (it is also called a polypeptide because it contains many peptide bonds). Proteins serve two broad purposes in the human body. Structural proteins form most of the solid material in the human body. For example, the structural proteins keratin and collagen are the main component of your hair, muscles, tendons and skin. Functional proteins help carry out activities and functions in the human body. For example, hemoglobin is a functional protein that occurs in the red blood cells and helps to transport oxygen in the body. Myosin is a protein that occurs in muscle tissue and is responsible for the ability of muscles to contract. Insulin is a functional protein that helps regulate the storage of the sugar glucose in the human body. A subclass of the functional proteins is the group of polypeptides referred to as enzymes. Enzymes help to carry out specific chemical reactions in the body. For example, amylase is an enzyme that occurs both in human saliva and in the intestines that helps to break apart the glucose-glucose bonds in the carbohydrate starch, thus allowing your body to absorb the glucose and use it for energy.

There are an estimated 100,000 different proteins in the human body alone, and each of them is made up of a combination of different combinations of only twenty amino acids. Each protein has a different structure and performs a different function in the body. When we eat protein-containing foods (such as meat, fish, beans, eggs, cheese, etc.) the polypeptide chains are generally broken down in the digestive tract and the individual amino acids are absorbed into our bodies. These amino acids are then recombined into proteins specific to each individual person in a process called protein synthesis.

In order to carry out these very precise jobs in the body, each individual protein has to be unique and specific to the job in question. Three aspects of a protein's structure are specific to the job the protein does in the body. The first aspect of a protein's structure is called the primary structure (1°). The primary structure of a protein is the sequence of amino acids in the protein. The number of amino acids in a protein can vary from the hundreds to the thousands, and the sequence in which those twenty different amino acids just mentioned occur (obviously one amino acid can occur in a protein many times) is specific to the individual protein, just as the sequence of numbers in your phone number is specific to your phone. The secondary structure (2°) of a protein is defined by the way the long strands of amino acids coil about themselves. Just as a phone cord wraps around itself to form a coil, a protein will also wrap around itself, and the degree and tightness of the coil is specific to the protein in question. Once a protein is coiled, the protein will begin to fold onto itself (similar to the way a phone cord tangles around itself); this folding is specific to the protein's function and is called the protein's tertiary structure (3°).
protein - structure representation - Representation of the 1° structure (amino acid sequence - illustrated with different colors), 2° structure (coiling), and 3° structure (folding) of a protein.

Representation of the 1° structure (amino acid sequence - illustrated with different colors), 2° structure (coiling), and 3° structure (folding) of a protein.

External Resources

• Chemistry Connections: The Chemical Basis of Everyday Phenomena, Second Edition (Complementary Science)

• Nature's Robots: A History of Proteins (Oxford Paperbacks)

• Other Recommended Products

Chemical Reactions

2009-02-02T19:06:00.000-08:00

Chemical Reactions

by Anthony Carpi, Ph.D.

The reaction of two or more elements together results in the formation of a chemical bond between atoms and the formation of a chemical compound (see our Chemical Bonding module). But why do chemicals react together? The reason has to do with the participating atoms' electron configurations (see our The Periodic Table of Elements module).

In the late 1890s, the Scottish chemist Sir William Ramsay discovered the elements helium, neon, argon, krypton, and xenon. These elements, along with radon, were placed in group VIIIA of the periodic table and nicknamed inert (or noble) gases because of their tendency not to react with other elements (see our Periodic Table page). The tendency of the noble gases to not react with other elements has to do with their electron configurations. All of the noble gases have full valence shells; this configuration is a stable configuration and one that other elements try to achieve by reacting together. In other words, the reason atoms react with each other is to reach a state in which their valence shell is filled.

Let's look at the reaction of sodium with chlorine. In their atomic states, sodium has one valence electron and chlorine has seven.
sodium-configuration with e chlorine
Sodium Chlorine

Chlorine, with seven valence electrons, needs one additional electron to complete its valence shell with eight electrons. Sodium is a little bit trickier. At first it appears that sodium needs seven additional electrons to complete its valence shell. But this would give sodium a -7 electrical charge and make it highly imbalanced in terms of the number of electrons (negative charges) relative to the number of protons (positive charges). As it turns out, it is much easier for sodium to give up its one valence electron and become a +1 ion. In doing so, the sodium atom empties its third electron shell and now the outermost shell that contains electrons, its second shell, is filled - agreeing with our earlier statement that atoms react because they are trying to fill their valence shell.
sodium chloride - Sodium Chloride

Sodium Chloride

This trait, the tendency to lose electrons when entering into chemical reactions, is common to all metals. The number of electrons metal atoms will lose (and the charge they will take on) is equal to the number of electrons in the atom's valence shell. For all of the elements in group A of the periodic table, the number of valence electrons is equal to the group number (see our Periodic Table page).

Nonmetals, by comparison, tend to gain electrons (or share them) to complete their valence shells. For all of the nonmetals, except hydrogen and helium, their valence shell is complete with eight electrons. Therefore, nonmetals gain electrons corresponding to the formula = 8 - (group #). Chlorine, in group 7, will gain 8 - 7 = 1 electron and form a -1 ion.

Hydrogen and helium only have electrons in their first electron shell. The capacity of this shell is two. Thus helium, with two electrons, already has a full valence shell and falls into the group of elements that tend not to react with others, the noble gases. Hydrogen, with one valence electron, will gain one electron when forming a negative ion. However, hydrogen and the elements on the periodic table labeled metalloids, can actually form either positive or negative ions corresponding to the number of valence electrons they have. Thus hydrogen will form a +1 ion when it loses its one electron and a -1 ion when it gains one electron.

Reaction Energy
All chemical reactions are accompanied by a change in energy. Some reactions release energy to their surroundings (usually in the form of heat) and are called exothermic. For example, sodium and chlorine react so violently that flames can be seen as the exothermic reaction gives off heat. On the other hand, some reactions need to absorb heat from their surroundings to proceed. These reactions are called endothermic. A good example of an endothermic reaction is that which takes place inside of an instant '"cold pack." Commercial cold packs usually consist of two compounds - urea and ammonium chloride in separate containers within a plastic bag. When the bag is bent and the inside containers are broken, the two compounds mix together and begin to react. Because the reaction is endothermic, it absorbs heat from the surrounding environment and the bag gets cold.

Reactions that proceed immediately when two substances are mixed together (such as the reaction of sodium with chlorine or urea with ammonium chloride) are called spontaneous reactions. Not all reactions proceed spontaneously. For example, think of a match. When you strike a match you are causing a reaction between the chemicals in the match head and oxygen in the air. The match won't light spontaneously, though. You first need to input energy, which is called the activation energy of the reaction. In the case of the match, you supply activation energy in the form of heat by striking the match on the matchbook; after the activation energy is absorbed and the reaction begins, the reaction continues until you either extinguish the flame or you run out of material to react.

External Resources

• The Development of Modern Chemistry

• Chemical Reactions (Reading Essentials in Science - Physical Science)

• The Joy of Chemistry: The Amazing Science of Familiar Things

• Other Recommended Products

Chemical Equations

2009-02-02T19:03:00.000-08:00

Chemical Equations

by Anthony Carpi, Ph.D.

Chemical reactions happen all around us: when we light a match, start a car, eat dinner, or walk the dog. A chemical reaction is the process by which substances bond together (or break bonds) and, in doing so, either release or consume energy (see our Chemical Reactions module). A chemical equation is the shorthand that scientists use to describe a chemical reaction. Let's take the reaction of hydrogen with oxygen to form water as an example. If we had a container of hydrogen gas and burned this in the presence of oxygen, the two gases would react together, releasing energy, to form water. To write the chemical equation for this reaction, we would place the substances reacting (the reactants) on the left side of an equation with an arrow pointing to the substances being formed on the right side of the equation (the products). Given this information, one might guess that the equation for this reaction is written:

H + O arrow H2O

The plus sign on the left side of the equation means that hydrogen (H) and oxygen (O) are reacting. Unfortunately, there are two problems with this chemical equation. First, because atoms like to have full valence shells, single H or O atoms are rare. In nature, both hydrogen and oxygen are found as diatomic molecules, H2 and O2, respectively (in forming diatomic molecules the atoms share electrons and complete their valence shells). Hydrogen gas, therefore, consists of H2 molecules; oxygen gas consists of O2. Correcting our equation we get:

H2 + O2 arrow H2O

But we still have one problem. As written, this equation tells us that one hydrogen molecule (with two H atoms) reacts with one oxygen molecule (two O atoms) to form one water molecule (with two H atoms and one O atom). In other words, we seem to have lost one O atom along the way! To write a chemical equation correctly, the number of atoms on the left side of a chemical equation has to be precisely balanced with the atoms on the right side of the equation. How does this happen? In actuality, the O atom that we "lost" reacts with a second molecule of hydrogen to form a second molecule of water. During the reaction, the H-H and O-O bonds break and H-O bonds form in the water molecules, as seen in the simulation below.

The formation of water

Concept simulation - Reenacts the reaction of hydrogen and oxygen in formation of water.

The balanced equation is therefore written:

2H2 + O2 arrow 2H2O

In writing chemical equations, the number in front of the molecule's symbol (called a coefficient) indicates the number of molecules participating in the reaction. If no coefficient appears in front of a molecule, we interpret this as meaning one.

In order to write a correct chemical equation, we must balance all of the atoms on the left side of the reaction with the atoms on the right side. Let's look at another example. If you use a gas stove to cook your dinner, chances are that your stove burns natural gas, which is primarily methane. Methane (CH4) is a molecule that contains four hydrogen atoms bonded to one carbon atom. When you light the stove, you are supplying the activation energy to start the reaction of methane with oxygen in the air. During this reaction, chemical bonds break and re-form and the products that are produced are carbon dioxide and water vapor (and, of course, light and heat that you see as the flame). The unbalanced chemical equation would be written:

CH4(methane) + O2(oxygen) arrow CO2(carbon dioxide) + H2O(water)

Look at the reaction atom by atom. On the left side of the equation we find one carbon atom, and one on the right.
C H4 + O2 arrow C O2 + H2 O

Next we move to hydrogen: There are four hydrogen atoms on the left side of the equation, but only two on the right.
C H4 + O2 arrow C O2 + H2 O

Therefore, we must balance the H atoms by adding the coefficient "2" in front of the water molecule (you can only change coefficients in a chemical equation, not subscripts). Adding this coefficient we get:
C H4 + O2 arrow C O2 + 2H2 O

What this equation now says is that two molecules of water are produced for every one molecule of methane consumed. Moving on to the oxygen atoms, we find two on the left side of the equation, but a total of four on the right side (two from the CO2 molecule and one from each of two water molecules H2O).

To balance the chemical equation we must add the coefficient "2" in front of the oxygen molecule on the left side of the equation, showing that two oxygen molecules are consumed for every one methane molecule that burns.
CH4 + 2O2 arrow CO2 + 2H2O

Dalton's law of definite proportions holds true for all chemical reactions (see our Matter module). In essence, this law states that a chemical reaction always proceeds according to the ratio defined by the balanced chemical equation. Thus, you can interpret the balanced methane equation above as reading, "one part methane reacts with two parts oxygen to produce one part carbon dioxide and two parts water." This ratio always remains the same. For example, if we start with two parts methane, then we will consume four parts O2 and generate two parts CO2 and four parts H2O. If we start with excess of any of the reactants (e.g., five parts oxygen when only one part methane is available), the excess reactant will not be consumed:
CH4 + 5O2 arrow CO2 + 2H2O + 3O2

Excess reactants will not be consumed.

In the example seen above, 3O2 had to be added to the right side of the equation to balance it and show that the excess oxygen is not consumed during the reaction. In this example, methane is called the limiting reactant.

Although we have discussed balancing equations in terms of numbers of atoms and molecules, keep in mind that we never talk about a single atom (or molecule) when we use chemical equations. This is because single atoms (and molecules) are so tiny that they are difficult to isolate. Chemical equations are discussed in relation to the number of moles of reactants and products used or produced (see our The Mole module). Because the mole refers to a standard number of atoms (or molecules), the term can simply be substituted into chemical equations. Thus, the balanced methane equation above can also be interpreted as reading, "one mole of methane reacts with two moles of oxygen to produce one mole of carbon dioxide and two moles of water."

Conservation of Matter
The law of conservation of matter states that matter is neither lost nor gained in traditional chemical reactions; it simply changes form. Thus, if we have a certain number of atoms of an element on the left side of an equation, we have to have the same number on the right side. This implies that mass is also conserved during a chemical reaction.

The total mass of the reactants, 36.04g, is exactly equal to the total mass of the products, 36.04g (if you are confused about these molecular weights, you should review the The Mole lesson). This holds true for all balanced chemical equations.

External Resources

• Periodic Table Playing Cards

• Chemical Demonstrations : A Handbook for Teachers of Chemistry Vol 4

• Other Recommended Products

Carbohydrates

2009-02-02T19:01:00.000-08:00

Carbohydrates

by Anthony Carpi, Ph.D.

In many ways, our bodies can be thought of as chemical processing plants. Chemicals are taken in, processed through various types of reactions, and then distributed throughout the body to be used immediately or stored for later use. The chemicals used by the body can be divided into two broad categories: macronutrients, those substances that we need to eat regularly in fairly large quantities, and micronutrients, those substances that we need only in small amounts. Three major classes of macronutrients are essential to living organisms: carbohydrates, fats, and proteins. In this lesson, we will discuss the carbohydrates; fats and proteins are discussed in another lesson (see our Fats and Protetin module).

Carbohydrates
Carbohydrates are the main energy source for the human body. Chemically, carbohydrates are organic molecules in which carbon, hydrogen, and oxygen bond together in the ratio: Cx(H2O)y, where x and y are whole numbers that differ depending on the specific carbohydrate to which we are referring. Animals (including humans) break down carbohydrates during the process of metabolism to release energy. For example, the chemical metabolism of the sugar glucose is shown below:
C6H12O6 + 6 O2 arrow 6 CO2 + 6 H2O + energy

Animals obtain carbohydrates by eating foods that contain them, for example potatoes, rice, breads, and so on. These carbohydrates are manufactured by plants during the process of photosynthesis. Plants harvest energy from sunlight to run the reaction just described in reverse:
6 CO2 + 6 H2O + energy (from sunlight) arrow C6H12O6 + 6 O2

A potato, for example, is primarily a chemical storage system containing glucose molecules manufactured during photosynthesis. In a potato, however, those glucose molecules are bound together in a long chain. As it turns out, there are two types of carbohydrates, the simple sugars and those carbohydrates that are made of long chains of sugars - the complex carbohydrates.

Simple Sugars

All carbohydrates are made up of units of sugar (also called saccharide units). Carbohydrates that contain only one sugar unit (monosaccharides) or two sugar units (disaccharides) are referred to as simple sugars. Simple sugars are sweet in taste and are broken down quickly in the body to release energy. Two of the most common monosaccharides are glucose and fructose. Glucose is the primary form of sugar stored in the human body for energy. Fructose is the main sugar found in most fruits. Both glucose and fructose have the same chemical formula (C6H12O6); however, they have different structures, as shown (note: the carbon atoms that sit in the "corners" of the rings are not labeled):
glucose fructose
Glucose Fructose

Disaccharides have two sugar units bonded together. For example, common table sugar is sucrose, a disaccharide that consists of a glucose unit bonded to a fructose unit:
sucrose molecule - Sucrose

Sucrose

Complex Carbohydrates
Complex carbohydrates are polymers of the simple sugars. In other words, the complex carbohydrates are long chains of simple sugar units bonded together (for this reason the complex carbohydrates are often referred to as polysaccharides). The potato we discussed earlier actually contains the complex carbohydrate starch. Starch is a polymer of the monosaccharide glucose:
starch molecule1 - Starch (n is the number of repeating glucose units and ranges in the 1,000\'s)

Starch
(n is the number of repeating glucose units and ranges in the 1,000's)

Starch is the principal polysaccharide used by plants to store glucose for later use as energy. Plants often store starch in seeds or other specialized organs; for example, common sources of starch include rice, beans, wheat, corn, potatoes, and so on. When humans eat starch, an enzyme that occurs in saliva and in the intestines called amylase breaks the bonds between the repeating glucose units, thus allowing the sugar to be absorbed into the bloodstream. Once absorbed into the bloodstream, the human body distributes glucose to the areas where it is needed for energy or stores it as its own special polymer - glycogen. Glycogen, another polymer of glucose, is the polysaccharide used by animals to store energy. Excess glucose is bonded together to form glycogen molecules, which the animal stores in the liver and muscle tissue as an "instant" source of energy. Both starch and glycogen are polymers of glucose; however, starch is a long, straight chain of glucose units, whereas glycogen is a branched chain of glucose units, as seen in the illustrations linked below:

The Starch Molecule • The Glycogen Molecule

Another important polysaccharide is cellulose. Cellulose is yet a third polymer of the monosaccharide glucose. Cellulose differs from starch and glycogen because the glucose units form a two-dimensional structure, with hydrogen bonds holding together nearby polymers, thus giving the molecule added stability. Cellulose, also known as plant fiber, cannot be digested by human beings, therefore cellulose passes through the digestive tract without being absorbed into the body. Some animals, such as cows and termites, contain bacteria in their digestive tract that help them to digest cellulose. Cellulose is a relatively stiff material, and in plants cellulose is used as a structural molecule to add support to the leaves, stem, and other plant parts. Despite the fact that it cannot be used as an energy source in most animals, cellulose fiber is essential in the diet because it helps exercise the digestive track and keep it clean and healthy.

The Cellulose Molecule

Related Modules

Organic Chemistry

Fats and Proteins

External Resources

• Carbohydrate Chemistry (Oxford Chemistry Primers, 99)

• The sweet science of glycobiology: complex carbohydrates, molecules that are particularly important for communication among cells, are coming under systematic ... An article from: American Scientist

The Mole

2009-02-02T18:58:00.000-08:00

The Mole
Its History and Use

by Anthony Carpi, Ph.D.

Simply put, the mole represents a number. Just as the term dozen refers to the number twelve, the mole represents the number 6.02 x 1023. (If you're confused by the form of this number refer to our The Metric System module).

Now that's a big number! While a dozen eggs will make a nice omelet, a mole of eggs will fill all of the oceans on earth more than 30 million times over. Think about it: It would take 10 billion chickens laying 10 eggs per day more than 10 billion years to lay a mole of eggs. So why would we ever use such a big number? Certainly the local donut store is not going to "supersize" your dozen by giving you a mole of jelly-filled treats.

The mole is used when we're talking about numbers of atoms and molecules. Atoms and molecules are very tiny things. A drop of water the size of the period at the end of this sentence would contain 10 trillion water molecules. Instead of talking about trillions and quadrillions of molecules (and more), it's much simpler to use the mole.

History of the Mole
The number of objects in one mole, that is, 6.02 x 1023, is commonly referred to as Avogadro's number. Amadeo Avogadro was an Italian physics professor who proposed in 1811 that equal volumes of different gases at the same temperature contain equal numbers of molecules. About fifty years later, an Italian scientist named Stanislao Cannizzaro used Avogadro's hypothesis to develop a set of atomic weights for the known elements by comparing the masses of equal volumes of gas. Building on this work, an Austrian high school teacher named Josef Loschmidt calculated the size of a molecule of air in 1865, and thus developed an estimate for the number of molecules in a given volume of air. While these early estimates have since been refined, they led to the concept of the mole - that is, the theory that in a defined mass of an element (its atomic weight) there is a precise number of atoms: Avogadro's number.

Molar Mass
A sample of any element with a mass equal to that element's atomic weight (in grams) will contain precisely one mole of atoms (6.02 x 1023 atoms). For example, helium has an atomic weight of 4.00. Therefore, 4.00 grams of helium will contain one mole of helium atoms. You can also work with fractions (or multiples) of moles:

Other atomic weights are listed on the periodic table (see our Periodic Table page). For each element listed, measuring out a quantity of the element equal to its atomic weight in grams will yield 6.02 x 1023 atoms of that element.

The atomic weight of an element identifies both the mass of one mole of that element and the total number of protons and neutrons in an atom of that element. How can that be? Let's look at hydrogen. One mole of hydrogen atoms will weigh 1.01 grams.
Hydrogen atom - A Hydrogen Atom

A Hydrogen Atom

Each hydrogen atom consists of one proton surrounded by one electron. But remember, the electron weighs so little that it does not contribute much to an atom's weight. Ignoring the weight of hydrogen's electrons, we can say that one mole of protons (H nuclei) weighs approximately one gram. Since protons and neutrons have about the same mass, a mole of either of these particles will weigh about one gram. For example, in one mole of helium, there are two moles of protons and two moles of neutrons - four grams of particles.

Molecular Weight
If you stand on a scale with a friend, the scale will register the combined weight of both you and your friend. When atoms form molecules, the atoms bond together, and the molecule's weight is the combined weight of all of its parts.

For example, every water molecule (H2O) has two atoms of hydrogen and one atom of oxygen. One mole of water molecules will contain two moles of hydrogen and one mole of oxygen.

Mole and Weight Relationships of Water and Its Parts. bottle filled with exactly 18.02 g water will contain 6.02 x 1023 water molecules. The concept of fractions and multiples described above also applies to molecules: 9.01 g of water would contain 1/2 mole, or 3.01 x 1023 molecules. You can calculate the molecular weight of any compound simply by summing the weights of atoms that make up that compound.

Molecular Weight Calculator

External Resources

• Periodic Table of the Elements Poster, 34" x 24"

• The mole concept in chemistry (Selected topics in modern chemistry)

• Other Recommended Products

Organic Chemistry

2009-02-02T18:57:00.000-08:00

Organic Chemistry
An Introduction

by Anthony Carpi, Ph.D.

To understand life as we know it, we must first understand a little bit of organic chemistry. Organic molecules contain both carbon and hydrogen. Though many organic chemicals also contain other elements, it is the carbon-hydrogen bond that defines them as organic. Organic chemistry defines life. Just as there are millions of different types of living organisms on this planet, there are millions of different organic molecules, each with different chemical and physical properties. There are organic chemicals that make up your hair, your skin, your fingernails, and so on. The diversity of organic chemicals is due to the versatility of the carbon atom. Why is carbon such a special element? Let's look at its chemistry in a little more detail.

Carbon (C) appears in the second row of the periodic table and has four bonding electrons in its valence shell (see our Periodic Table module for more information). Similar to other non-metals, carbon needs eight electrons to satisfy its valence shell. Carbon therefore forms four bonds with other atoms (each bond consisting of one of carbon's electrons and one of the bonding atom's electrons). Every valence electron participates in bonding, thus a carbon atom's bonds will be distributed evenly over the atom's surface. These bonds form a tetrahedron (a pyramid with a spike at the top), as illustrated below:
carbon bonds - Carbon forms 4 bonds

Carbon forms 4 bonds

Organic chemicals get their diversity from the many different ways carbon can bond to other atoms. The simplest organic chemicals, called hydrocarbons, contain only carbon and hydrogen atoms; the simplest hydrocarbon (called methane) contains a single carbon atom bonded to four hydrogen atoms:
carbon-methane - Methane - a carbon atom bonded to 4 hydrogen atoms

Methane - a carbon atom bonded to 4 hydrogen atoms

But carbon can bond to other carbon atoms in addition to hydrogen, as illustrated in the molecule ethane below:
carbon-ethane - Ethane - a carbon-carbon bond

Ethane - a carbon-carbon bond

In fact, the uniqueness of carbon comes from the fact that it can bond to itself in many different ways. Carbon atoms can form long chains:
carbon-hexane - Hexane - a 6-carbon chain

There appears to be almost no limit to the number of different structures that carbon can form. To add to the complexity of organic chemistry, neighboring carbon atoms can form double and triple bonds in addition to single carbon-carbon bonds:
c-ethane c-ethene c-ethyne

Keep in mind that each carbon atom forms four bonds. As the number of bonds between any two carbon atoms increases, the number of hydrogen atoms in the molecule decreases (as can be seen in the figures above).

Simple Hydrocarbons
The simplest hydrocarbons are those that contain only carbon and hydrogen. These simple hydrocarbons come in three varieties depending on the type of carbon-carbon bonds that occur in the molecule. Alkanes are the first class of simple hydrocarbons and contain only carbon-carbon single bonds. The alkanes are named by combining a prefix that describes the number of carbon atoms in the molecule with the root ending "ane". The names and prefixes for the first ten alkanes are given in the

The chemical formula for any alkane is given by the expression CnH2n+2. The structural formula, shown for the first five alkanes in the table, shows each carbon atom and the elements that are attached to it. This structural formula is important when we begin to discuss more complex hydrocarbons. The simple alkanes share many properties in common. All enter into combustion reactions with oxygen to produce carbon dioxide and water vapor. In other words, many alkanes are flammable. This makes them good fuels. For example, methane is the principle component of natural gas, and butane is common lighter fluid.

CH4 + 2O2 arrow CO2 + 2H2O

The combustion of methane

The second class of simple hydrocarbons, the alkenes, consists of molecules that contain at least one double-bonded carbon pair. Alkenes follow the same naming convention used for alkanes. A prefix (to describe the number of carbon atoms) is combined with the ending "ene" to denote an alkene. Ethene, for example is the two- carbon molecule that contains one double bond. The chemical formula for the simple alkenes follows the expression CnH2n. Because one of the carbon pairs is double bonded, simple alkenes have two fewer hydrogen atoms than alkanes.
carbon-ethene - Ethene

Ethene

Alkynes are the third class of simple hydrocarbons and are molecules that contain at least one triple-bonded carbon pair. Like the alkanes and alkenes, alkynes are named by combining a prefix with the ending "yne" to denote the triple bond. The chemical formula for the simple alkynes follows the expression CnH2n-2.

Isomers
Because carbon can bond in so many different ways, a single molecule can have different bonding configurations. Consider the two molecules illustrated here:

Both molecules have identical chemical formulas (shown in the left column); however, their structural formulas (and thus some chemical properties) are different. These two molecules are called isomers. Isomers are molecules that have the same chemical formula but different structural formulas.

Functional Groups
In addition to carbon and hydrogen, hydrocarbons can also contain other elements. In fact, many common groups of atoms can occur within organic molecules, these groups of atoms are called functional groups. One good example is the hydroxyl functional group. The hydroxyl group consists of a single oxygen atom bound to a single hydrogen atom (-OH). The group of hydrocarbons that contain a hydroxyl functional group is called alcohols. The alcohols are named in a similar fashion to the simple hydrocarbons, a prefix is attached to a root ending (in this case "anol") that designates the alcohol. The existence of the functional group completely changes the chemical properties of the molecule. Ethane, the two-carbon alkane, is a gas at room temperature; ethanol, the two-carbon alcohol, is a liquid.
carbon-ethanol - Ethanol

Ethanol

Ethanol, common drinking alcohol, is the active ingredient in "alcoholic" beverages such as beer and wine.

External Resources

• Organic Molecular Model Kit

• The Most Beautiful Molecule: The Discovery of the Buckyball

• Other Recommended Products

Electron Configuration Shorthand:

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Electron Configuration Shorthand:
For elements in groups labeled A in the periodic table (IA, IIA, etc.), the number of valence electrons corresponds to the group number. Thus Li, Na, and other elements in group IA have one valence electron. Be, Mg, and other group-IIA elements have two valence electrons. B, Al and other group-IIIA elements have three valence electrons, and so on. The row, or period, number that an element resides in on the table is equal to the number of total shells that contain electrons in the atom. H and He in the first period normally have electrons in only the first shell; Li, Be, B, and other period-two elements have two shells occupied, and so on. To write the electron configuration of elements, scientists often use a shorthand in which the element's symbol is followed by the element's electron shells, written as a right-hand parentheses symbol ")". The number of electrons in each shell is then written after the ) symbol. A few examples are shown below.

Electron Configuration and the Table

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Electron Configuration and the Table

The "periodic" nature of chemical properties that Mendeleev had discovered is related to the electron configuration of the atoms of the elements. In other words, the way in which an atom's electrons are arranged around its nucleus affects the properties of the atom.

Bohr's theory of the atom tells us that electrons are not located randomly around an atom's nucleus, but they occur in specific electron shells (see our Atomic Theory II module for more information). Each shell has a limited capacity for electrons. As lower shells are filled, additional electrons reside in more-distant shells.

The capacity of the first electron shell is two electrons and for the second shell the capacity is eight. Thus, in our example discussed above, oxygen, with eight protons and eight electrons, carries two electrons in its first shell and six in its second shell. Fluorine, with nine electrons, carries two in its first shell and seven in the second. Neon, with ten electrons, carries two in the first and eight in the second. Because the number of electrons in the second shell increases, we can begin to imagine why the chemical properties gradually change as we move from oxygen to fluorine to neon.

Sodium has eleven electrons. Two fit in its first shell, but remember that the second shell can only carry eight electrons. Sodium's eleventh electron cannot fit into either its first or its second shell. This electron takes up residence in yet another orbit, a third electron shell in sodium. The reason that there is a dramatic shift in chemical properties when moving from neon to sodium is because there is a dramatic shift in electron configuration between the two elements. But why is sodium similar to lithium? Let's look at the electron configurations of these elements.

As you can see in the illustration, while sodium has three electron shells and lithium two, the characteristic they share in common is that they both have only one electron in their outermost electron shell. These outer-shell electrons (called valence electrons) are important in determining the chemical properties of the elements.

An element's chemical properties are determined by the way in which its atoms interact with other atoms. If we picture the outer (valence) electron shell of an atom as a sphere encompassing everything inside, then it is only the valence shell that can interact with other atoms - much the same way as it is only the paint on the exterior of your house that "interacts" with, and gets wet by, rain water.

The valence shell electrons in an atom determine the way it will interact with neighboring atoms, and therefore determine its chemical properties. Since both sodium and lithium have one valence electron, they share similar chemical

Water as a Solvent

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Water as a Solvent

The partial charge that develops across the water molecule helps make it an excellent solvent. Water dissolves many substances by surrounding charged particles and "pulling" them into solution. For example, common table salt, sodium chloride, is an ionic substance that contains alternating sodium and chlorine ions.
NaCl-crystal - Sodium chloride contains Na+ and Cl- ions.

Sodium chloride contains Na+ and Cl- ions.

When table salt is added to water, the partial charges on the water molecule are attracted to the Na+ and Cl- ions. The water molecules work their way into the crystal structure and between the individual ions, surrounding them and slowly dissolving the salt. The water molecules will actually line up differently depending on which ions are being pulled into solution. The negative oxygen ends of water molecules will surround the positive sodium ions; the positive hydrogen ends will surround the negative chlorine ions.
NaCl-dissolve - Table Salt Dissolving in Water

Table Salt Dissolving in Water

In a similar fashion, any substance that carries a net electrical charge, including both ionic compounds and polar covalent molecules (those that have a dipole), can dissolve in water. This idea also explains why some substances do not dissolve in water. Oil, for example, is a nonpolar molecule. Because there is no net electrical charge across an oil molecule, it is not attracted to water molecules and therefore does not dissolve in water.

From: http://www.visionlearning.com/library/module_viewer.php?c3=&mid=57&l=

The Periodic Table of Elements

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The Periodic Table of Elements

by Anthony Carpi, Ph.D.

In 1869, the Russian chemist Dmitri Mendeleev first proposed that the chemical elements exhibited a "periodicity of properties." Mendeleev had tried to organize the chemical elements according to their atomic weights, assuming that the properties of the elements would gradually change as atomic weight increased. What he found, however, was that the chemical and physical properties of the elements increased gradually and then suddenly changed at distinct steps, or periods. To account for these repeating trends, Mendeleev grouped the elements in a table that had both rows and columns.

The modern periodic table of elements is based on Mendeleev's observations; however, instead of being organized by atomic weight, the modern table is arranged by atomic number (z). As one moves from left to right in a row of the periodic table, the properties of the elements gradually change. At the end of each row, a drastic shift occurs in chemical properties. The next element in order of atomic number is more similar (chemically speaking) to the first element in the row above it; thus a new row begins on the table.

For example, oxygen (O), fluorine (F), and neon (Ne) (z = 8, 9 and 10, respectively) all are stable nonmetals that are gases at room temperature. Sodium (Na, z = 11), however, is a silver metal that is solid at room temperature, much like the element lithium (z = 3). Thus sodium begins a new row in the periodic table and is placed directly beneath lithium, highlighting their chemical similarities.

Rows in the periodic table are called periods. As one moves from left to right in a given period, the chemical properties of the elements slowly change. Columns in the periodic table are called groups. Elements in a given group in the periodic table share many similar chemical and physical properties. The link below will open a copy of the periodic table of elements in a new window.

From: http://www.visionlearning.com/library/module_viewer.php?c3=&mid=57&l=

Water Properties and Behavior

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Water
Properties and Behavior

by Anthony Carpi, Ph.D.

In many ways, water is a miracle liquid. It is essential for all living things (on this planet at least), and it is often referred to as a universal solvent because many substances dissolve in it. These unique properties of water result from the ways in which individual H2O molecules interact with each other.
watermolecule2
Electronic Distribution in H2O

In the Chemical Bonding lesson we discussed the dipole that forms across the water molecule as a result of the polar covalent bonding between hydrogen and oxygen (see our Chemical Bonding module for more information). Because the bonding electrons are shared unequally by the hydrogen and oxygen atoms, a partial negative charge (ð-) forms at the oxygen end of the water molecule, and a partial positive charge (ð+) forms at the hydrogen ends. Since the hydrogen and oxygen atoms in the molecule carry opposite (though partial) charges, nearby water molecules are attracted to each other like tiny little magnets. The electrostatic attraction between the ð+ hydrogen and the ð- oxygen in adjacent molecules is called hydrogen bonding.
hydrogen - bond - Hydrogen Bonding between Water Molecules

Hydrogen Bonding between Water Molecules

Hydrogen bonding makes water molecules "stick" together. While hydrogen bonds are relatively weak compared to other types of bonds, they are strong enough to give water many unique properties. For example, hydrogen bonds sank the Titanic, and hydrogen bonds allow the Basilisk lizard to walk on water (as a result, the Basilisk has earned the nickname "Jesus" lizard).

Just how does hydrogen bonding do this? Well, let's start with the Titanic. The Titanic sank because it hit an iceberg - a chunk of ice floating on the surface of the ocean. The reason ice floats is because of hydrogen bonding. In water's liquid form, hydrogen bonding pulls water molecules together. As a result, liquid water has a relatively compact, dense structure. The animation below illustrates this idea.

Liquid Water and Hydrogen Bonding

Concept simulation - Reenacts hydrogen bonding between molecules of liquid water.

As water freezes into ice, the molecules become frozen in place and begin to arrange themselves in a rigid lattice structure, as shown in the animation linked below.

Ice and Hydrogen Bonding

Concept simulation - Reenacts hydrogen bonding between molecules of solid water.

The structure that forms in the solid ice crystal actually has large holes in it. Therefore, in a given volume of ice, there are fewer water molecules than in the same volume of liquid water. In other words, ice is less dense than liquid water and will float on the surface of the liquid. Throw in one really big chunk of ice and a cruise ship, and you begin to see the problems that can arise.

Surface Tension: As we just discussed, neighboring water molecules are attracted to one another. Molecules at the surface of liquid water have fewer neighbors and, as a result, have a greater attraction to the few water molecules that are nearby. This enhanced attraction is called surface tension. It makes the surface of the liquid slightly more difficult to break through than the interior.

When a small object that would normally sink in water is placed carefully on the surface, it can remain suspended on the surface due to surface tension. The Basilisk lizard makes use of the high surface tension of water to accomplish the incredible feat of walking on water's surface. The Basilisk can't actually walk on water; rather, it runs on water, moving its feet before they break through the surface.

from: http://www.visionlearning.com/library/module_viewer.php?c3=&mid=57&l=

Covalent Bonding

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Covalent Bonding

The second major type of atomic bonding occurs when atoms share electrons. As opposed to ionic bonding in which a complete transfer of electrons occurs, covalent bonding occurs when two (or more) elements share electrons. Covalent bonding occurs because the atoms in the compound have a similar tendency for electrons (generally to gain electrons). This most commonly occurs when two nonmetals bond together. Because both of the nonmetals will want to gain electrons, the elements involved will share electrons in an effort to fill their valence shells. A good example of a covalent bond is that which occurs between two hydrogen atoms. Atoms of hydrogen (H) have one valence electron in their first electron shell. Since the capacity of this shell is two electrons, each hydrogen atom will "want" to pick up a second electron. In an effort to pick up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H2. Because the hydrogen compound is a combination of equally matched atoms, the atoms will share each other's single electron, forming one covalent bond. In this way, both atoms share the stability of a full valence shell.

Unlike ionic compounds, covalent molecules exist as true molecules. Because electrons are shared in covalent molecules, no full ionic charges are formed. Thus covalent molecules are not strongly attracted to one another. As a result, covalent molecules move about freely and tend to exist as liquids or gases at room temperature.

Multiple Bonds: For every pair of electrons shared between two atoms, a single covalent bond is formed. Some atoms can share multiple pairs of electrons, forming multiple covalent bonds. For example, oxygen (which has six valence electrons) needs two electrons to complete its valence shell. When two oxygen atoms form the compound O2, they share two pairs of electrons, forming two covalent bonds.

Lewis Dot Structures: Lewis dot structures are a shorthand to represent the valence electrons of an atom. The structures are written as the element symbol surrounded by dots that represent the valence electrons. The Lewis structures for the elements in the first two periods of the periodic table are shown below.

Lewis structures can also be used to show bonding between atoms. The bonding electrons are placed between the atoms and can be represented by a pair of dots or a dash (each dash represents one pair of electrons, or one bond). Lewis structures for H2 and O2 are shown below.

Polar and Nonpolar Covalent Bonding
There are, in fact, two subtypes of covalent bonds. The H2 molecule is a good example of the first type of covalent bond, the nonpolar bond. Because both atoms in the H2 molecule have an equal attraction (or affinity) for electrons, the bonding electrons are equally shared by the two atoms, and a nonpolar covalent bond is formed. Whenever two atoms of the same element bond together, a nonpolar bond is formed.

A polar bond is formed when electrons are unequally shared between two atoms. Polar covalent bonding occurs because one atom has a stronger affinity for electrons than the other (yet not enough to pull the electrons away completely and form an ion). In a polar covalent bond, the bonding electrons will spend a greater amount of time around the atom that has the stronger affinity for electrons. A good example of a polar covalent bond is the hydrogen-oxygen bond in the water molecule.
water molecule-with caption Water molecules contain two hydrogen atoms (pictured in red) bonded to one oxygen atom (blue). Oxygen, with six valence electrons, needs two additional electrons to complete its valence shell. Each hydrogen contains one electron. Thus oxygen shares the electrons from two hydrogen atoms to complete its own valence shell, and in return shares two of its own electrons with each hydrogen, completing the H valence shells.

Polar covalent bonding simulated in water

The primary difference between the H-O bond in water and the H-H bond is the degree of electron sharing. The large oxygen atom has a stronger affinity for electrons than the small hydrogen atoms. Because oxygen has a stronger pull on the bonding electrons, it preoccupies their time, and this leads to unequal sharing and the formation of a polar covalent bond.

The Dipole
Because the valence electrons in the water molecule spend more time around the oxygen atom than the hydrogen atoms, the oxygen end of the molecule develops a partial negative charge (because of the negative charge on the electrons). For the same reason, the hydrogen end of the molecule develops a partial positive charge. Ions are not formed; however, the molecule develops a partial electrical charge across it called a dipole. The water dipole is represented by the arrow in the pop-up animation (above) in which the head of the arrow points toward the electron dense (negative) end of the dipole and the cross resides near the electron poor (positive) end of the molecule.

Ionic Bonding

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Ionic Bonding
In ionic bonding, electrons are completely transferred from one atom to another. In the process of either losing or gaining negatively charged electrons, the reacting atoms form ions. The oppositely charged ions are attracted to each other by electrostatic forces, which are the basis of the ionic bond.

For example, during the reaction of sodium with chlorine:
Sodium&Chlorine-transfer sodium (on the left) loses its one valence electron to chlorine (on the right),
arrow-down resulting in
SodiumChlorineIons a positively charged sodium ion (left) and a negatively charged chlorine ion (right).

The reaction of sodium with chlorine

Concept simulation - Reenacts the reaction of sodium with chlorine.

(Flash required)

Notice that when sodium loses its one valence electron it gets smaller in size, while chlorine grows larger when it gains an additional valence electron. This is typical of the relative sizes of ions to atoms. Positive ions tend to be smaller than their parent atoms while negative ions tend to be larger than their parent. After the reaction takes place, the charged Na+ and Cl- ions are held together by electrostatic forces, thus forming an ionic bond. Ionic compounds share many features in common:

* Ionic bonds form between metals and nonmetals.
* In naming simple ionic compounds, the metal is always first, the nonmetal second (e.g., sodium chloride).
* Ionic compounds dissolve easily in water and other polar solvents.
* In solution, ionic compounds easily conduct electricity.
* Ionic compounds tend to form crystalline solids with high melting temperatures.

This last feature, the fact that ionic compounds are solids, results from the intermolecular forces (forces between molecules) in ionic solids. If we consider a solid crystal of sodium chloride, the solid is made up of many positively charged sodium ions (pictured below as small gray spheres) and an equal number of negatively charged chlorine ions (green spheres). Due to the interaction of the charged ions, the sodium and chlorine ions are arranged in an alternating fashion as demonstrated in the schematic. Each sodium ion is attracted equally to all of its neighboring chlorine ions, and likewise for the chlorine to sodium attraction. The concept of a single molecule becomes blurred in ionic crystals because the solid exists as one continuous system. Forces between molecules are comparable to the forces within the molecule, and ionic compounds tend to form crystal solids with high melting points as a result.From: http://www.visionlearning.com/library/module_viewer.php?c3=&mid=55&l=

Chemical Bonding

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Chemical Bonding

by Anthony Carpi, Ph.D.

Though the periodic table has only 118 or so elements, there are obviously more substances in nature than 118 pure elements. This is because atoms can react with one another to form new substances called compounds (see our Chemical Reactions module). Formed when two or more atoms chemically bond together, the resulting compound is unique both chemically and physically from its parent atoms.

Let's look at an example. The element sodium is a silver-colored metal that reacts so violently with water that flames are produced when sodium gets wet. The element chlorine is a greenish-colored gas that is so poisonous that it was used as a weapon in World War I. When chemically bonded together, these two dangerous substances form the compound sodium chloride, a compound so safe that we eat it every day - common table salt!
sodium metal + chlorine gas arrow salt&shaker

In 1916, the American chemist Gilbert Newton Lewis proposed that chemical bonds are formed between atoms because electrons from the atoms interact with each other. Lewis had observed that many elements are most stable when they contain eight electrons in their valence shell. He suggested that atoms with fewer than eight valence electrons bond together to share electrons and complete their valence shells.

While some of Lewis' predictions have since been proven incorrect (he suggested that electrons occupy cube-shaped orbitals), his work established the basis of what is known today about chemical bonding. We now know that there are two main types of chemical bonding; ionic bonding and covalent bonding.

from: http://www.visionlearning.com/library/module_viewer.php?c3=&mid=55&l=

Ionic Molecules

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Ionic Molecules

Ionic and covalent bonds constitute the ends of a spectrum of bonding possibilities. Ionic molecules are formed by the transfer of one or more electrons from one atom to another. This transfer results in the formation of cations and anions. These oppositely charged particles are then held tightly together by very strong Coulombic attractions. Equation 2 illustrates the formation of sodium chloride from the reaction of a sodium atom with a chlorine atom.

Exercise 1 What is the valence shell electron configuration of a sodium ion? What is the valence shell electron configuration of a chloride ion?

Exercise 2 If the bond in a sodium chloride molecule were covalent, i.e. if the sodium and chlorine atoms shared a pair of electrons, what would the valence shell electron configuration of the sodium be? What would the valence shell electron configuration of the chlorine be?

There are two properties that are characteristic of ionic molecules: 1. They are generally high melting solids 2. When dissolved in water, they produce solutions which conduct an electrical current.

From: http://usm.maine.edu/~newton/Chy251_253/Lectures/AcidBase/DynamicsFS.html

What are Acids and Bases - on a molecular level?

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What are Acids and Bases - on a molecular level?

Acids:	Bases:
Arrhenius Definition (1887) - covers the dissociation of acids and bases in water.
Produce hydrogen ions (H⁺)	Produce hydroxide ions (OH^-)
Example: HCl HCl ---> H⁺) + Cl^-	NaOH NaOH ---> Na⁺ + OH^-
In the graphic, the first reaction shows citric acid producing hydrogen ions, therefore it is an acid.

Bronsted-Lowery Definition (1923) - The definition of acids and bases involving hydrogen and hydroxide ions, respectively is much too limiting. A broader definition was proposed by Bronsted and Lowry in 1923. The main effect of the definition is to increase the number of substances that act as bases.

Acid	Base
Donates hydrogen ions	Accepts hydrogen ions.
HCl +	HOH --->	H₃O⁺ + Cl^-
HOH +	NH₃--->	NH₄⁺ + OH^-

The determination of a substance as a Bronsted-Lowery acid or base can only be done by observing the reaction. In the case of the HOH it is a base in the first case and an acid in the second case.

Link to Chime animation of ammonium ion to water transfer - Jeremy Harvey, University of Bristol, England

See the graphic on the left for an example:

To determine whether a substance is an acid or a base, count the hydrogens on each substance before and after the reaction. If the number of hydrogens has decreased that substance is the acid (donates hydrogen ions). If the number of hydrogens has increased that substance is the base (accepts hydrogen ions). These definitions are normally applied to the reactants on the left.

If the reaction is viewed in reverse a new acid and base can be identified. The substances on the right side of the equation are called conjugate acid and conjugate base compared to those on the left.

Also note that the original acid turns in the conjugate base after the reaction is over.

From: http://www.elmhurst.edu/~chm/vchembook/180acidsbases.html

Acids and Bases

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Acids and Bases

Introduction and Definitions:

Acids and bases are encountered frequently both in chemistry and in everyday living. They have opposite properties and have the ability to cancel or neutralize each other. Acids and bases are carefully regulated in the body by the lungs, blood, and kidneys through equilibrium processes.

What are acids and bases?

Observational definitions:

Acids:	Bases:
Taste sour.	Taste bitter.
Give sharp stinging pain in a cut or wound.	Feels slippery
Turn blue litmus paper red.	Turn red litmus paper blue.
Turn phenolphthalein colorless.	Turn phenolphthalein pink.
React with metals to produce hydrogen gas.
React with carbonates or bicarbonates to produce carbon dioxide gas.

See the graphic on the left as a base is wiped over the paper which already has the indicator phenolphthalein on it. The drawing changes to pink.

From: http://www.elmhurst.edu/~chm/vchembook/180acidsbases.html

Covalent Molecules

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Covalent Molecules

As we saw in our discussion of valence bond theory, covalent bonds involve the sharing of electrons between atoms rather than the transfer of electrons from one atom to another. Equation 3 depicts the formation of a covalent bond between a hydrogen atom and a chlorine atom.

Exercise 3 What is the valence shell electron configuration of the hydrogen in HCl? What is the valence shell electron configuration of the chlorine?

Exercise 4 If the bond in a hydrogen chloride molecule were ionic, i.e. if the hydrogen transferred its valence electron to the chlorine, what would the valence shell electron configuration of the hydrogen be? What would the valence shell electron configuration of the chlorine be?

In contrast to ionic compounds, covalent molecules may be gases, liquids, or solids. Those that are solids generally have melting points below 300^oC. Furthermore, aqueous solutions of most covalent molecules do not conduct a current. There are some notable exceptions however. HCl is one.

Molecular Basis for the Indicator Color Change

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Molecular Basis for the Indicator Color Change:

Color changes in molecules can be caused by changes in electron confinement. More confinement makes the light absorbed more blue, and less makes it more red.

How are electrons confined in phenolphthalein? There are three benzene rings in the molecule. Every atom involved in a double bond has a p orbital which can overlap side-to-side with similar atoms next to it. The overlap creates a 'pi bond' which allows the electrons in the p orbital to be found on either bonded atom. These electrons can spread like a cloud over any region of the molecule that is flat and has alternating double and single bonds. Each of the benzene rings is such a system.

See the far left graphic - The carbon atom at the center (adjacent to the yellow circled red oxygen atom) doesn't have a p-orbital available for pi-bonding, and it confines the pi electrons to the rings. The molecule absorbs in the ultraviolet, and this form of phenolphthalein is colorless.

In basic solution, the molecule loses one hydrogen ion. Almost instantly, the five-sided ring in the center opens and the electronic structure around the center carbon changes (yellow circled atoms) to a double bond which now does contain pi electrons. The pi electrons are no longer confined separately to the three benzene rings, but because of the change in geometry around the yellow circled atoms, the whole molecule is now flat and electrons are free to move within the entire molecule. The result of all of these changes is the change in color to pink.

Chime: Phenolphthalein

Link to further explanation - General Chemistry Online! Water to Wine by Fred Senese

Many other indicators behave on the molecular level in a similar fashion (the details may be different) but the result is a change in electronic structure along with the removal of a hydrogen ion from the molecule. Plant pigments in flowers and leaves also behave in this fashion.

from: http://www.elmhurst.edu/~chm/vchembook/186indicator.html

Magic Pitcher Demonstration:

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Magic Pitcher Demonstration:

Phenolphthalein is an indicator of acids (colorless) and bases (pink). Sodium hydroxide is a base, and it was in the pitcher at the beginning, so when added to the phenolphthalein in beakers 2 and 4, it turned pink (top half of the graphic).

Explanation:

Equilibrium: HIn --> H⁺ + In^-
colorless red

The equilibrium shifts right, HIn decreases, and In^- increases. As the pH increase between 8.2 to 10.0 the color becomes red because of the equilibrium shifts to form mostly In^- ions.

The third beaker has only the NaOH but no phenolphthalein, so it remained colorless. The first beaker contain acetic acid and is skipped over at first.

After pouring beakers 2, 3, 4 back into the pitcher it give a pink solution.

Bottom half of the graphic: When the pitcher is then poured back into beakers 2, 3, 4 it is a pink solution.

In the first beaker, a strange thing happens in that the pink solution coming out of the pitcher now changes to colorless. This happens because the first beaker contains some vinegar or acetic acid which neutralizes the NaOH, and changes the solution from basic to acidic. Under acidic conditions, the phenolphthalein indicator is colorless.

Neutralization: HC₂H₃O₂ + NaOH --> Na(C₂H₃O₂) + HOH

Explain the color indicator change:

Use equilibrium principles to explain the color change for phenolphthalein at the end of the demonstration.

Solution:

The simplified reaction is: H⁺ + OH^- --> HOH

As OH^- ions are added, they are consumed by the excess of acid already in the beaker as expressed in the above equation. The hydroxide ions keep decreasing and the hydrogen ions increase, pH decreases.

See lower equation: The indicator equilibrium shifts left, In^- ions decrease. Below pH 8.2 the indicator is colorless. As H⁺ ions are further increased and pH decreases to pH 4-5, the indicator equilibrium is effected and changes to the colorless HIn form.

Equilibrium: HIn --> H⁺ + In^-
colorless red

Acid and Base Indicators

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Acid and Base Indicators

Acid - Base Indicators:

The most common method to get an idea about the pH of solution is to use an acid base indicator. An indicator is a large organic molecule that works somewhat like a " color dye". Whereas most dyes do not change color with the amount of acid or base present, there are many molecules, known as acid - base indicators , which do respond to a change in the hydrogen ion concentration. Most of the indicators are themselves weak acids.

The most common indicator is found on "litmus" paper. It is red below pH 4.5 and blue above pH 8.2.

Color	Blue Litmus	Red Litmus
Acid	turns red	stays same
Base	stays same	turns blue

Other commercial pH papers are able to give colors for every main pH unit. Universal Indicator, which is a solution of a mixture of indicators is able to also provide a full range of colors for the pH scale.

A variety of indicators change color at various pH levels. A properly selected acid-base indicator can be used to visually "indicate" the approximate pH of a sample. An indicator is usually some weak organic acid or base dye that changes colors at definite pH values. The weak acid form (HIn) will have one color and the weak acid negative ion (In^-) will have a different color. The weak acid equilibrium is:

HIn --> H+ + In^-
For phenolphthalein: pH 8.2 = colorless; pH 10 = red

For bromophenol blue: pH 3 = yellow; pH 4.6 = blue

See the graphic for more indicators, colors, and pH ranges.

from: http://www.elmhurst.edu/~chm/vchembook/186indicator.html

Lewis Acid-Base Experiment

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Introduction

Lewis acids, which are electron pair acceptors, and Lewis bases, which are electron pair donors, react to form adducts in which a coordinate covalent bond is formed. This type of bond is usually represented by an arrow. The strength of the interaction between a Lewis acid and a Lewis base is controlled by at least two factors, electronic and steric (1). Electron donating groups on an atom can increase the Lewis basicity of that atom, while electron-withdrawing groups can increase the Lewis acidity. Trivalent nitrogen and boron compounds are common examples of Lewis bases and acids, respectively. Trimethylamine, N(CH₃)₃, is a Lewis base and because of the electron donating methyl groups, it should be a stronger base than NH₃ when both are compared with a common Lewis acid. Similarly, one would expect BF₃ to be a stronger Lewis acid than B(CH₃)₃.

Steric factors affect both Lewis acids and bases in similar ways. A trivalent boron compound is triangular planar with an sp²hybridized boron atom. When a Lewis acid-base adduct is formed, the boron atom is rehybridized to sp³ and the three groups on boron are forced closer together. Large, bulky groups on boron will not favor being forced closer together and this will affect the strength of the interaction. As well, a trivalent nitrogen compound is pyramidal and the nitrogen is sp³. No rehybridization of the nitrogen occurs upon adduct formation, but large, bulky groups on nitrogen will prefer a more planar geometry to minimize non-bonded repulsions. When the Lewis acid-base adduct is formed, the groups on nitrogen are also forced closer together which again, affects the strength of the interaction. Finally, large, bulky groups on both boron and nitrogen can prevent the close approach of the bonding atoms. This is sometimes referred to as F-strain.

One measure of the strength of a Lewis acid-base interaction is the heat of reaction, DH_rxn. Heats of reaction can be estimated from the heats of formation, DH_f, of the reactants and product. Recall that, DH_rxn = S[DH_f(products)] - S[DH_f (reactants)]. The more negative the DH_rxn, the stronger the Lewis acid-base interaction. The heats of formation are calculated from a semi-empirical calculation. In this experiment, the modeling program CAChe was used and the MOPAC application was used to obtain thermodynamic values. The resulting heats of reaction may not agree well with experimental results in all cases, but the trends in DH_rxn observed for a single series of reactions are consistent with expected results of acid-base strength.

Modeling and Calculation Exercises

In each of the studies (1-5) below it is necessary to build the reactants and products in the CAChe Editor, mimimize the geometry of each with Molecular Mechanics and then minimze geometry again with MOPAC ( AM1). The MOPAC-results will produce the heat of formation for each molecule. Also, from the MOPAC-minimized-geometry atom distances and bond angles can be obtained using the Adjust menu in Editor. The acid-base bond must be a Coordinate bond from base to acid.

Boron Trihalides + Ammonia
Consider the reactions of BX₃ (X = F, Cl, Br) with NH₃. Calculate the heats of reaction and rank the boron halides in acid strength. Also, it may be informative to obtain the B-N bond distance for each adduct as well as the values of some bond angles which may undergo a change upon reaction. Discuss your results in terms of the electronic properties of the boron halides (2,3). It will be necessary to consider p-bonding effects in the boron halides. The results obtained for this series of reactions are shown in Table I below.

(In Table I below and all others that follow, LA = Lewis acid, LB = Lewis base and all values are in kcal/mol):

Table I

LA DH_f(LA) LB DH_f(LB) DH_f(adduct) DH_rxn

BF₃ -272.15 NH₃ -7.28 -291.81 -12.38
BCl₃ -97.01 NH₃ -7.28 -140.14 -35.85
BBr₃ -50.05 NH₃ -7.28 -91.39 -34.06

Trialkylboron Compounds + Ammonia

Consider the reactions of BR₃ [R = CH₃ (Me), C₂H₅ (Et), C₃H₇ (n-Pr and i-Pr)] with ammonia. Calculate the heats of reactions as before from the heats of formation. Discuss your results in terms of both the changing electronic and steric effects on the boron atom. The four trialkylboron Lewis acids are shown below and the results obtained for this series of reactions are contained in Table II.

Table II

LA	DH_f(LA)	LB	DH_f(LB)	DH_f(adduct)	DH_rxn

BMe₃	-24.60	NH₃	-7.28	-54.48	-22.60
BEt₃	-38.13	NH₃	-7.28	-70.74	-25.33
B(n-Pr)₃	-58.27	NH₃	-7.28	-90.05	-24.50
B(i-Pr)₃	-45.07	NH₃	-7.28	-70.60	-18.25

Trialkylboron Compounds + Trialkylamines

In this series of reactions the Lewis base remains the same at trimethylboron while the alkyl groups on nitrogen increase in size. Again, obtain the heats of formation and calculate the heats of reaction. Interpret the trend in heats of reaction in terms of the changing electronic and steric effects on the nitrogen atom. The four trialkyamines are shown below and the results obtained for this series of reactions are contained in Table III.

Table III

LA	DH_f(LA)	LB	DH_f(LB)	DH_f(adduct)	DH_rxn

BMe₃	-24.60	NMe₃	-1.71	-26.01	0.30
BMe₃	-24.60	NEt₃	-14.93	-38.09	1.44
BMe₃	-24.60	N(n-Pr)₃	-35.89	-58.04	2.45
BMe₃	-24.60	N(i-Pr)₃	-19.40	-14.96	29.40

Pyridine and Substituted Pyridines + Trimethylboron

In this study pyridine (Py), 4-methylpyridine (4-MePy), 2-methylpyridine (2-MePy) and 2,6-dimethylpyridine (2,6-diMePy) will be allowed to react with trimethylboron. This series of pyridines represents an interesting "control" of steric and electronic factors. Interpret the heats of reaction results in terms of both the changing electronic and steric factors of the four pyridines which are shown below (4). The thermodynamic results are contained in Table IV.

Table IV

LA	DH_f(LA)	LB	DH_f(LB)	DH_f(adduct)	DH_rxn

BMe₃	-24.60	Py	32.05	3.43	-4.02
BMe₃	-24.60	4-MePy	24.17	-4.73	-4.30
BMe₃	-24.60	2-MePy	25.72	2.33	1.21
BMe₃	-24.60	2,6-DiMePy	19.51	2.32	7.41

Other Interesting Studies

Basicity of Triethylamine vs. Quinuclidine.- Consider reactions of both triethylamine [NEt₃] and quinuclidine [N(CH₂CH₂)₃CH] with trimethylboron (4). The two bases, which are shown below, should have similar electronic effects, but they should have different steric constraints. The nitrogen lone pair on quinuclidine should be more accessible to a Lewis acid than that of triethylamine. The thermodynamic results are shown in Table V.

Basicity of Trimethylamine vs. Trisilylamine.- Compare the reactions of trimethylboron (BMe₃) with both trimethylamine (NMe₃) and trisilylamine (N(SiH₃)₃). The MOPAC-minimized-geometry of N(SiH₃)₃ shows the compound to be planar with an sp² nitrogen, while the geometry of N(CH₃)₃ is pyramidal with an sp³ nitrogen. The planar geometry in N(SiH ₃)₃ is presumably due to pp- dp bonding between a nitrogen 2p orbital and empty 3d orbitals on Si (5). This p-bonding decreases the Lewis basicity of N(SiH₃)₃ by delocalizing the nitrogen lone pair. The steric effects between N(SiH₃)₃ and NMe₃ should be very similar, and any differences in the heats of reacton can be attributed to electronic effects. The thermodynamic results are collected in Table V.

Table V

LA	DH_f(LA)	LB	DH_f(LB)	DH_f(adduct)	DH_rxn

BMe₃	-24.60	NEt₃	-14.93	-38.09	1.44
BMe₃	-24.60	N(CH₂CH ₂)₃CH	-8.18	-34.49	-1.71
BMe₃	-24.60	NMe₃	-1.71	-25.98	0.32
BMe₃	-24.60	N(SiH₃)₃	-30.57	-32.58	22.58

References

Muetterties, E. L. The Chemistry of Boron and Its Compounds; John Wiley & Sons, Inc.: New York, 1967; pp. 9, 232.
Greenwood, N. N.; Earnshaw, A. Chemistry of the Elements; Pergamon Press: New York, 1984; p. 220.
Brown, H. C.; Holmes, R. R. J. Am. Chem. Soc. 1956, 78, 2173.
Cotton, F. A.; Wilkinson, G.; Gaus, P. L. Basic Inorganic Chemistry; 3rd ed.; John Wiley & Sons, Inc.: New York, 1995; p 228.
Livant, P.; McKee, M. L.; Worley, S. D. Inorg. Chem. 1983, 22, 895-901.

The Greek Concept of Atomos: The Indivisible Atom

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The Greek Concept of Atomos: The Indivisible Atom

Page references are to S. Sambursky (1956) "The Physical World of the Greeks" Princeton University Press. ISBN 0-691-02411-1.

I continue to grow in my knowledge. Atomistic theory is prominent in some of the Hindu teachings in India.

Around 440 BC, Leucippus of Miletus originated the atom concept. He and his pupil, Democritus (c460-371 BC) of Abdera, refined and extended it in future years. There are five major points to their atomic idea. Almost all of the original writings of Leucippus and Democritus are lost. About the only sources we have for their atomistic ideas are found in quotations of other writers.

Democritus [16K GIF] is known as the "Laughing Philosopher" because of his joyous spirit. He was a big man (relatively speaking) and enjoyed life tremendously. He also was very widely traveled.

This map [16K GIF] shows the important towns of Greece around the time the atom concept was developed. It is about 250 miles as the crow flies between the Abdera and Miletus.

At this time Greek philosophy was about 150 years old, having emerged early in the sixth century BC, centered in the city of Miletus on the Ionian coast in Asia Minor (now Turkey). The earliest known Greek philosopher was Thales of Miletus.

The work of Leucippus and Democritus was further developed by Epicurus (341-270 BC) of Samos, who made the ideas more generally known. Aristotle (384-322 BC) quotes both of them extensively in arguing against their ideas. Much of what we know about their ideas comes to us in a poem titled "De Rerum Natura" (On the Nature of Things) written by Lucretius (c95-55 BC). This poem, lost for over 1000 years, was rediscovered in 1417.

You may find the Leonard translation of "De Rerum Natura" on-line, but it is a big file.

On the left is a 18K GIF of Aristotle and to the right is a 19K GIF of Epicurus.

Point #1 - All matter is composed of atoms, which are bits of matter too small to be seen. These atoms CANNOT be further split into smaller portions.

Democritus quotes Leucippus: "The atomists hold that splitting stops when it reaches indivisible particles and does not go on infinitely." (p. 107)

In other words, there is a lower limit to the division of matter beyond which we cannot go. Atoms were impenetrably hard, meaning they could not be divided. In Greek, the prefix "a" means "not" and the word "tomos" means cut. Our word atom therefore comes from atomos, a Greek word meaning uncuttable.

Democritus reasoned that if matter could be infinitely divided, it was also subject to complete disintegration from which it can never be put back together. However, matter can be reintegrated.

Even though matter can be destroyed by repeated splitting, new things can be made by joining simpler pieces of matter together. The process of disintegration & reintegration is reversible.

The idea of reversibility means that there must be a lower limit to the splitting of matter. If matter can be split infinitely, there is nothing to stop it from going on forever and destroying all matter.

Only with a definite and finite lower limit to splitting do we keep a permament foundation of ultimate particles with which to build up everything we see. As Epicurus says:

"Therefore, we must not only do away with division into smaller and smaller parts to infinity, in order that we may not make all things weak, and so in the composition of aggregate bodies be compelled to crush and squander the things that exist into the non-existent...." (p. 108)

Epicurus also insisted on an upper limit for atoms - they are always invisible. Although no reason is given, it seems obvious enough: all matter that can be seen by humans is still divisible, therefore cannot be atoms.

Point #2 - There is a void, which is empty space between atoms.

Aristotle quotes Leucippus: "Unless there is a void with a separate being of its own, 'what is' cannot be moved-nor again can it be 'many', since there is nothing to keep things apart." (p. 108)

In other words, there is empty space between atoms. In modern times, we would use the word vacuum, although the Greeks did not.

Given that all matter is composed of atoms (the ultimate and unchanging particles), then all changes must be as a result of the movement of atoms. However, in order to move there must be a void-a space entirely empty of matter-through which atoms can move from place to place.

Point #3 - Atoms are completely solid.

It then follows that there can be no void inside an atom itself. Otherwise an atom would be subject to changes from outside and could disintegrate. Then, it would not be an atom.

We know this is incorrect. In 1911, Ernest Rutherford discovered the nucleus, demonstrating in the process that a single atom is mostly empty space.

Point #4 - Atoms are homogeneous, with no internal structure.

The absolute solidity of the atoms also leads to the notion that atoms are homogeneous, or the same all the way through. Another way to express this is that an atom would have no internal structure.

Although there was speculation about sub-atomic structure in the 1800's after John Dalton introduced the atom idea on a solid scientific basis, it was not until 1897 and J.J. Thomson's discovery of the electron that the atom was shown to have an internal structure.

Point #5 - Atoms are different in ...

1) ...their sizes. See the Democritus quote just below.

2) ...their shapes. According to Aristotle: "Democritus and Leucippus say that there are indivisible bodies, infinite both in number and in the varieties of their shapes...." (p. 110)

Democritus says of atoms: "They have all sorts of shapes and appearences and different sizes.... Some are rough, some hook-shaped, some concave, some convex and some have other innumerable variations." (p. 110-111)

3) ...their weight. Again from Aristotle: "Democritus recognized only two basic properties of the atom: size and shape. But Epicurus added weight as a third. For, according to him, the bodies move by necessity through the force of weight." (p. 111)

Concluding Remarks

The idea of the atom was strongly opposed by Aristotle and others. Because of this, the atom receeded into the background. Although there is a fairly continuous pattern of atomistic thought through the ages, only a relative few scholars gave it much thought.

Due to complex circumstances beyond the scope of this lesson, the Catholic Church accepted Aristotle's position and came to equate atomistic ideas with Godlessness. For example, "Democritus of Abdera said that there is no end to the universe, since it was not created by any outside power."

It was not until 1660 that Pierre Gassendi succeeded in separating the two and not until 1803 that John Dalton put the atom on a solid scientific basis. The atom concept is often presented as laying fallow between Democritus and Dalton.

Atomic Structure from Democritus to Dalton

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Atomic Structure from Democritus to Dalton

In 1803, John Dalton of England introduced the atomic idea to chemistry (and is called the Father of Modern Atomic Theory for his efforts). However, it would be false to assume that atomic ideas disappeared completely from the intellectual map for over 2000 years. For, although atomic thinkers between the Greeks and Dalton were few, there is a fairly continuous line from the Greeks to John Dalton.

Much of the following is based on these articles:

1) "The Origins of the Atomic Theory" by J.R. Partington. Annals of Science, vol 4, no. 3 (July, 1939)
2) "The Atomic View of Matter in the XVth, XVIth, and XVIIth Centuries" by G.B. Stones. Isis, vol. 10, part 2, No. 34 (January 1928).

I. Atomism in Antiquity

The atomic ideas of Leucippus and Democritus (from about 440 BC) were opposed by Aristotle about 100 years or so later. Those who acknowledged Aristotle as their master opposed atoms. Since Epicurus was an atomist, he was opposed by his rivals, the Stoics. Cicero, Seneca and Galen all spoke against atoms.

Hero of Alexandria (150 A.D.?) makes use of atoms to explain compression and rarefaction (to thin something out; become less dense). Hero denied the existence of an extended vacuum, but allowed for a vacuum between atoms. One proof he cited was that fire could enter into a material, showing that it had openings, i.e., a vacuum. He pointed out that the pores of a diamond were too small to let in fire and so the stone was incombustible. (In the 1700's, both Lavosier and Priestly were able to burn diamonds with large lenses that concentrated the sunlight.)

Important figures within the Church spoke against atoms. Dionysios (Bishop of Alexandria 200 A.D.), Lactantius (died 324 A.D.) and Augustine (354-430 A.D.) are names cited by Partington.

II. Atomism in the Middle Ages

Isidore, Bishop of Seville (560-636), the Venerable Bede (672-735), and Hrabanus Maurus (776-856) all used the word "atom" to refer to discontunities in bodies. William of Conches (1080-1154) and Vincent of Beauvais (died ca. 1264-8) both showed knowledge of atomic thinking in their writing. William openly taught about the ideas of Democritus. Vincent wrote a great encyclopedia, but only gave short quotations about atoms.

The works of Aristotle were rediscovered by Western Europe about 1200, in Latin translations of Arabic translations from the Greek. Much scholastic discussion followed among such people as St. Thomas Aquinas (1225-74) and Roger Bacon (1214-92). Over time, the Catholic Church began to elevate Aristotle's writings to the same level as Scripture and had associated atomic thinking with Godlessness. (Quite frankly, the ChemTeam does not know how the process took place, but it did. On a side issue, the Church also did the same thing with Ptolemy in astronomy. When Galileo opposed the Church (in the 1630's?), he was found guilty of heresy. Only recently (around the late 1980's-early 1990's) has the Church formally admitted its error.)

De Rerum Natura was rediscovered in 1417 (and printed in 1473, reprinted in 1486) and became the prime source (still true today) for the ideas of Leucippus and Democritus. You may ask how William of Conches knew of Democritus. Scattered about the libraries of churches in Europe were a few copies of De Rerum Natura. Stones, in his article, cites three known to have existed in William's lifetime. Other copies certainly also existed at that time.

III. Atomism in the Renaissance

A) Nicholas of Cusa (1401-1464) wrote:

"What dost thou understand by an atom?
"Under mental consideration that which is continuous becomes divided into the ever divisible, and the multitude of parts progresses to infinity. But by actual division we arrive at an actually indivisible part which I call an atom. For an atom is a quantity, which on account of its smallness is actually indivisible."

B) Girolamo Fracastoro (1478-1553) was a physician who wrote about atomism. In fact, the phrase "seeds of disease" is asociated with his name. In discussion the mechanism of infection, he supposed the existence of minute indivisible substances which convey the disease. he called these semina. Interestingly, Lucretius (in Book VI) refers to seeds helpful to life and seeds which cause disease and death. In a different book, Fracastoro indicates his agreement with Democritus and puts forward an atomistic point of view concerning chemical reactions.

C) Peter Ramus (1515-1572) broke with Aristotle early in his life. (Remember, the Catholic Church had long ago elevated Aristotle's works to Scripture. In essence, both were considered to be infallible.) At age 21, he presented a thesis based on this idea: "all that Aristotle has said is false." His opponents could not just appeal to the authority of Aristotle to refute him, since that would be begging the question. After attacking his ideas for a whole day and being refuted, Ramus was finally awarded his degree with honors.

In 1543, he wrote two books (aganist Aristotle) that provoked violent reaction. Their publication was banned, the books were burned, and Ramus was silenced by order of the Pope, Francis I. After the Pope died a year later, Ramus resumed teaching and was appointed professor in 1551.

However, he embraced the Reformed faith (Martin Luther had nailed his "95 Theses" to church door at the University of Wittenberg on October 31, 1517.) and was forced to flee from Paris. His home was pillaged and his library burned. He returned eventually, but ultimately died in the massacre of St. Bartholomew in Paris in 1572.

Although it appears that Ramus did not write about atomism as such, he was in the forefront of the attack on the authority of Aristotle.

D) In 1588, Giordano Bruno wrote:

"The division of natural things has a limit; an indivisible something exists. The division of natural things attains the smallest and last parts which are not perceptible by the aid of human instruments."

E) Partington lists five other names of people alive through in the 1500's and 1600's who wrote about atoms. Of interest is Sebastin Basso, who wrote of particles of the first, second, and third order, that is to say, structures BUILT UP by bringing atoms together. What we might call a molecule today. J.C. Magnenus attempted to calculate the size of an atom.

F) Daniel Sennert (1572-1637) was an atomist during the time Rene Descartes (1596-1650) and Francis Bacon (1561-1626) were alive. Both Bacon and Descartes, although intellectual giants of that era, were not too convinced about atomism.

Sennert taught that there must be atoms of more than one type and that atoms joined together to form composite bodies (I think he called these secondary atoms, but I am not sure). He used the fact that vapor from wine penetrated 4 layers of paper to show the smallness of atoms. Another example was that a large volume of vapor yielded a small drop of liquid.

He also taught that atoms retain their essential form. For example, melt some pure gold and pure silver together until completely mixed. On treating the mixture with nitric acid, the silver is dissolved and the gold remains.

G) Partington dates the real beginning of the revival of atomic thinking to the invention of the barometer in 1634 by Evangalestia Torricelli. Above the mercury of the barometer was a vacuum. An important position of Aristotle (and the Church) was that the vacuum did not exist. This invention (and the air pump by Otto von Guericke in 1654) dealt a severe, if not crippling, blow to the non-existence of the vacuum.

IV. Pierre Gassendi (1592-1655)

Gassendi is considered by many to be the reviver of atomism, but as you have seen, atomism never really went away, it was just on the fringes. However, Gassendi was successful in making atomism more widely known and acceptable, especially by separating a belief in atomism from athesism.

Before going into his teachings, it is interesting to note that in 1624, the Parliment of Paris had issued a decree that anyone holding or teaching a position opposed to Aristotle (including atomism) was liable to be put to death. Gassendi has influential friends, so he was left alone.

In 1649 he published his major work on atomism: Syntagma philosophiae Epicuri. It is divided into three sections: Logic, Physics, and Ethics.

Before even discussing atoms, Gassendi devotes three chapters to discussing the void and its necessity. He dwells on Torricelli and his experiments at length.

He describes the Greek position: atoms cannot be created nor destroyed, they are solid, they have weight, and cannot be subdivided. Gassendi taught that atoms are not just geometric points, but that they have a definite size, although it is very small.

However, he differs from the Greeks in that atoms have not been in existence forever, but were made by God. The atoms move not a se ipsis (of themselves), but Dei gratia (as a gift of God). This is the idea which freed atomism from athesism.

Gassendi allows for the union of atoms to form groups, which he calls moleculae or corpuscula. However, these groups are not held together by attractive forces, but by mechanical forces such as hooks-and-eyes or antlers.

V. From Gassendi to Dalton: Just Under 150 years

Robert Boyle (1627-91) was an atomist, although he liked the word "corpuscle." In 1661, published the "Sceptical Chymist." In it, he insists that the chemical elements must be actual, physical substances rather than the "principles" the alchemists thought of (the "principle of salt", the "principle of gold" and so on). Boyle says:

"I can easily enough sublime gold into the form of red Chrystalls of a considerable length; and many other ways may Gold be disguis'd, and help to constitute Bodies of very different Natures both from It and from one another, and nevertheless be afterwards reduc'd to the self-same Numerical, Yellow, Fixt, Ponderous, and Malleable Gold as it was before its commixture."

Later on in the book, he says of atoms (oops, sorry Bob, corpuscles) of gold:

"though they may not be primary Concretions of the most minute Particles of matter, but confessedly mixt Bodies, are able to concurre plentifully in the composition of several very differing bodies without losing their own Nature or Texture, or having their cohesion violated by the divorce of their associated parts or ingredients.

Again, he says:

"the difference of Bodies may depend meerly upon that of the schemes whereinto their Common matter is put . . . so that according as the small parts of matter reccede from each other, or work upon each other . . . a Body of this or that denomination is producd."

Incidently, two of the last non-believers in the reality of atoms were Wilhelm Ostwald and Ernst Mach. (I am not including those who are not in the mainstream of science, Ostwald and Mach were both respected scientists.) In 1908, Ostwald explicitly stated his belief in the reality of atoms in the introduction to his textbook Outline of General Chemistry. In 1915, Mach was still writing in an anti-atomistic way. The following year, Mach died, aged 78.

Since then, no one of any scientific substance has questioned the reality of atoms.

John Dalton (1766-1844): The Father of the Chemical Atomic Theory

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John Dalton (1766-1844): The Father of the Chemical Atomic Theory

Before delving into Dalton, I want to draw a difference between physical and chemical atomism.

The path that Dalton took to the chemical atomic theory is complex. As Leonard K. Nash points out, there are three more-or-less contemporary descriptions of how Dalton developed his ideas. This includes Dalton's own words on the subject. Add to this the several theories advanced over the years by historians and there is a lot to read, study and ponder about. And then, it turns out that the three contemporary descriptions are mutually contradictory and none are consistent with information available elsewhere.

The ChemTeam believes that the most satisfactory answer to Dalton's path to the chemical atomic theory was by way of his studies on vapor pressure, gas solubility and gas mixtures. I intend to develop this topic in the future.

John Dalton wrote his first table of atomic weights in his notebook dated September 1803. In 1830, in a paper read to the Manchester Literary and Philosophical Society, he said:

"A series of Essays read before this society and afterwards published in the 5th Vol. of their Memoirs gradually led me to the consideration of ultimate particles or atoms & their combinations. Under the date of Sept 3d, 1803, I find in my notebook 'Observations on the ultimate particles of bodies and their combinations,' in which the atomic symbols I still use [were] introduced. On the 23rd of October the same year[I r]ead my Essay on the absorption of Gases [by Water] at the conclusion of which a series of atomic [weights] was given for 21 simple and compound elements .. . . "

t was not until 1805 that the above essay was published and it was not until 1808 that Dalton himself discussed his methods for atomic weight determination. He published his theories on the atmosphere and gas behavior in a book titled A New System of Chemical Philosophy. Only in the last few pages (Chapter III) did he discuss his atomic theory.

Modern scholarship has identified four basic ideas in Dalton's chemical atomic theory.

1) chemical elements are made of atoms
2) the atoms of an element are identical in their masses
3) atoms of different elements have different masses
4) atoms only combine in small, whole number ratios such as 1:1, 1:2, 2:3 and so on.

1) elements are made of atoms
Elements are made up of minute, discrete, indivisible, and indestructible particles called atoms. These atoms maintain their identity through all physical and chemical changes. This, of course, is not a new idea to Dalton. This basic idea goes back to the Greeks. However, please keep in mind that atoms, as such, were not part of the chemical mainstream in the early 1800's.

Dalton's idea of an element is what we believe today - an element is a chemical substance that cannot be decomposed further by chemical means (i.e. heat, electricity, reacting with another chemical). This definition traces to Lavoisier.

Daltonian atoms are usually taught as being similar to featureless billiard balls. In truth, Dalton never ruled out the possibility of subatomic structure. He just knew that the state of the art in the early 1800's did not allow the physical structure of an atom to be probed.

2) the atoms of an element are identical in their masses
Atoms of the same element have the same properties, such as weight. Atoms of different elements have different properties, including a different weight. The idea that all atoms of a given chemical element weigh the same is known today to be incorrect, but in 1803 the concept of isotopes was just over 100 years in the future.

Also, the concept of chemical combination in 1803 was much, much different than what Dalton was proposing. Although Dalton was well-known at the time, the most authoritative chemist of the period was Claude Louis Berthollet and his ideas were phrased thus:

"Berthollet has shown also, that every body, how weak soever its affinity for another maybe, is capable of abstracting part of that other from a third, how strong soever the affinity of that third is, provded it be applied in sufficient quantity."

Even many years later, Berthollet resisted the idea of the atom: that elements combine in small, whole number ratios that are fixed. Even into the late 1800's, there were French chemists who used their authority to punish lesser colleagues and students who publically supported the chemical atomic theory.

Incidently, please do not think that Berthollet was a total loser. On the contrary. He was the first, in 1798, to observe a reversible reaction and ideas like the ones expressed in the above quote worked perfectly well for chemical substances reacting, just not for determining atomic weights. Some of the negative reaction to his anti-atom stance seems to have spilled over into unjusly ignoring his other work. Consequently, it was not until the mid-1860's that chemical equilibrium began to be explored in depth.

3) atoms of different elements have different masses
Although this idea is implicit in Dalton's theory, it is not original with him. Once again, the Greeks. developed this general idea. This idea was even discussed in chemical textbooks of Dalton's time, so the idea was "in the air," so to speak.

However, here is where we meet the original contribution of Dalton. Arnold W. Thackray says (his original is in italics):

"The particular development of Dalton, which distinguishes his chemical atomic theory from earlier work, was his devising of an effective system to obtain these relative particle weights from currently available chemical data. . . ."

In other words, while it was claimed atoms of different elements had different weights, no one could figure out what the different weight values were. Dalton was the first to do so.

So, exactly how did he determine atomic weights? I intend to develop this topic in the future.

4) atoms combine in small, whole number ratios
Chemical combination between two or more atoms occur in simple, numerical ratios (i.e., 1 to 1, 1 to 2; 2 to 3; etc.).

This point gives immediate explanation to the Law of Definite Proportions, announced by Joseph Louis Proust in 1797.

During his research, Dalton discovered the Law of Multiple Proportions, another law which is easily explained by his atomic theory.

Dalton discovered this law while studying some of the oxides of nitrogen.I will defer discussion of this work to a future time. The law, in modern terminology, is:

Atoms of the same element can unite in more than one ratio with another element to form more than one compound.

A fifth idea implicit In Dalton's theory, but usually not discussed is this: atoms can be neither created nor destroyed. An element's atoms do not change into other element's atoms by chemical reactions. For example, nitrogen and oxygen atoms stay as themselves even when combined. They can be recovered by decomposing the substance. As Dalton says:

"We might as well attempt to introduce a new planet into the solar system, or to annihilate one already in existence, as to create or destroy a particle of hydrogen. All the changes we can produce, consist in separating particles that are in a state of cohesion or combination, and joining those that were previously at a distance."

As with most of Dalton's theory, this idea is not original to Dalton. It is Lavoisier who is responsible for the Law of Conservation of Mass in chemical reactions.

Finally, a few Dalton factoids to close:

1) In "A New System of Chemical Philosophy," he proposed standard symbols for the elements. He was the first to do so.
2) He was the first to identify color-blindness. He, himself, suffered from red-green color-blindness. Still today, "daltonism" is often used to name this problem.
3) The unit for atomic weight was called a "dalton" for many years. In modern times, you most often hear it used in biochemical circles, as in "The atomic weight of that protein is 35,000 daltons."

J.J. Thomson and the Discovery of the Electron

2009-01-28T19:22:00.001-08:00

J.J. Thomson and the Discovery of the Electron

On April 30, 1897, Joseph John (J.J.) Thomson (1856-1940) [13K GIF] announced that cathode rays were negatively charged particles which he called 'corpuscles.' He also announced that they had a mass about 1000 times smaller than a hydrogen atom, and he claimed that these corpuscles were the things from which atoms were built up. Later in 1897, he wrote:

"...we have in the cathode rays matter in a new state, a state in which the subdivision of matter is carried very much farther than in the ordinary gaseous state: a state in which all matter-that is, matter derived from different sources such as hydrogen, oxygen, etc.-is of one and the same kind; this matter being the substance from which the chemical elements are built up." (J.J. Thomson (1897). "Cathode Rays," Philosophical Magazine 44, 295.)

He had leaped to the conclusion that the particles in the cathode ray (which we now call electrons) were a fundamental part of all matter. This was reaching quite far beyond what he had actually discovered. As he was to recall much later:

"At first there were very few who believed in the existence of these bodies smaller than atoms. I was even told long afterwards by a distinguished physicist who had been present at my [1897] lecture at the Royal Institution that he thought I had been `pulling their legs.' " (J.J. Thomson (1936). Recollections and Reflections. G. Bell and Sons: London. p. 341.)

Thomson's corpuscle hypothesis was not generally accepted, even by British scientists, until he spoke of it again in 1899. By this time, George Francis FitzGerald (1851-1901), [14K GIF] an Irish physicist, had suggested that Thomson's 'corpuscles' making up the cathode ray were actually free electrons. In fact, this suggestion was published as a commentary to the publication of Thomson's April 30, 1897 lecture in which he first announced his results. Thomson himself continued to use the term corpuscle until 1913.

Other people had measured the e/m ratio or suggested that the cathode rays were composed of particles, but Thomson was the first to say that the cathode ray was a building block of the atom. It was a risky thing, but he was proved right and for his courage he is remembered as the discoverer of the electron.

Walter Kaufmann deserves special mention before leaving this subject. In 1897, he had better data than Thomson and had it months before him. However, Kaufmann was a follower of a philosophy called "positivism," championed at that time by Ernst Mach (whose name is used today to signify the speed of sound - Mach one; twice the speed of sound - Mach two, and so on). Positivism allowed explanations of events which were based on sensory experience only. Since submicroscopic particles were not seen by the human senses, but rather inferred from the data, Kaufmann could not bring himself to the "corpuscle hypothesis" that Thomson announced. So it was that Kaufmann missed out on a great discovery and become a footnote to history. By the way, Kaufmann was convinced by 1901 of the electron's existence and became a leading experimenter working to determine more about it.

The Thomson Model of the Atom

2009-01-28T19:22:00.000-08:00

The Thomson Model of the Atom

In 1897, J.J. Thomson discovered the electron, the first subatomic particle. He also was the first to attempt to incorporate the electron into a structure for the atom. The internal structure of the atom had been a source of speculation for thousands of years. The Greeks taught that the atom was solid, as did Dalton. Although Dalton did allow for the fact that there might be a sub-atomic structure of which he was unaware.

Thomson faced two major problems: (1) how to account for the mass of the atom when the electron was only about 1/1000 the mass of the hydrogen atom (the more modern figure is 1/1836) and (2) how to create a neutral atom when the only particle available was negatively charged.

His solution was to rule the scientific world for about a decade and Thomson himself would make a major contribution to undermining his own model.

I. Leadup to Thomson's 1904 Model of the Atom

Thomson had been in the business of proposing atomic models since at least 1881, which is when he proposed his "vortex" model of the atom. We will not go into details about it.

The first seed of the model we are discussing appear in his famous 1897 announcement of the discovery of the electron. He wrote:

"The explanation which seems to me to account in the most simple and straightforward manner for the facts is founded on a view of the constitution of the chemical elements which has been favourably entertained by many chemists: this view is that the atoms of the different chemical elements are different aggregations of atoms of the same kind. In the form in which this hypothesis was enunciated by Prout, the atoms of the different elements were hydrogen atoms; in this precise form the hypothesis is not tenable, but if we substitute for hydrogen some unknown primordial substance X, there is nothing known which is inconsistent with this hypothesis, which is one that has been recently supported by Sir Norman Lockyer for reasons derived from the study of the stellar spectra.
If, in the very intense electric field in the neighbourhood of the cathode, the molecules of the gas are dissociated and are split up, not into the ordinary chemical atoms, but into these primordial atoms, which we shall for brevity call corpuscles; and if these corpuscles are charged with electricity and projected from the cathode by the electric field, they would behave exactly like the cathode rays. "

And a few paragraphs later:

"If we regard the chemical atom as an aggregation of a number of primordial atoms, . . . ."

However, he does not go into the presence of a positive force, although he must have been aware of its necessisity.

Here is what he then said in 1899:

"I regard the atom as containing a large number of smaller bodies which I will call corpuscles, these corpuscles are equal to each other.... In the normal atom, this assemblage of corpuscles forms a system which is electrically neutral. Though the individual corpuscles behave like negative ions, yet when they are assembled in a neutral atom the negative effect is balanced by something which causes the space through which the corpuscles are spread to act as if it had a charge of positive electricity equal in amount to the sum of the negative charges of the corpuscles.... The detached corpuscles behave like negative ions, each carrying a constant negative charge which we shall call for brevity the unit charge; while the part of the atom left behind behaves like a positive ion with the unit positive charge and a mass large compared with that of the negative ion."

This last portion is interesting in that it proposes the correct mechanism for ionization; a negative electron is removed leaving behind a positive atom.

II. Thomson's Mature Model

His next statement on the structure of the atom comes in a 1904 article. The first half of the article is filled with detailed calculations about the stability of corpuscles moving about in a positive environment. In fact, Thomson is only able to make calculations where all the corpuscles are limited to roatating in a ring. Moving from ring to sphere proves too difficult a challenge.

Here is a quote from the 1904 article:

We suppose that the atom consists of a number of corpuscles moving about in a sphere of uniform positive electrification . . . .

That seems pretty straighforward, but the problem will soon become the electrons and their mass.

By the way, this is often referred to as Thomson's "plum pudding model," where the pudding represents the sphere of positive electricity and the bits of plum scattered in the pudding are the electrons. The ChemTeam likes to call it the "blueberry muffin" model. All those round little blueberries surrounded by the bread of the muffin. Ummmm, good. Some butter on top of a muffin hot from the oven and some nice, COLD milk. Oh my.

You can read more of Thomson's 1904 article in the classic papers section.

However, not everyone is convinced this is the right answer. Savante Arrhenius (the 1903 Nobel Prize winner in Chemistry) had this to say about Thomson's model in 1907:

"This conception has hitherto remained only a formal one, and has led to no new results."

Arrhenius goes on to several criticisms of the Thomson Model.

Before leaving this topic, I want to make a point about how the Thomson Model is presented today. Sometimes teachers, and even textbooks, will represent the Thomson Model as a mixture of protons and electrons, like on the right-hand side of this image:

Make sure you have the correct idea firmly in mind. The Thomson Model has negative partices (electrons) and a sphere of positive charge. There are NO protons in the Thomson Model of the atoms. Be careful, a teacher might try to trip you up on a test question. (Those teachers sure are evil, aren't they??)

The Thomson Model will hold sway for a few years, until Ernest Rutherford announces the nuclear model of the atom in 1911. This tutorial: A Brief History of Rutherford's Experiment starts the story.

Interest in the Thomson Model fell off rapidly after 1911, although in 1914 and 1915 attempts were made to resurrect it. These efforts came to nothing and the Thomson Model assumed its place in history as the first modern attempt to construct a theory of atomic structure.

Blueprints on Fabric

2009-01-27T19:52:00.000-08:00

	1. Layout the design so you know exactly what to expect.
	2. Mix the solution for making the blueprint (for 6 yards of fabric, use 1/2 gallon of water, 4 oz. of potassium ferricyanide, 8 oz. of ferric ammonium citrate. Use glass or plastic tools for mixing.
	3. Apply the solution to the fabric on the desired area. Apply thoroughly and evenly. Do not worry if you apply some to areas outside of the desired area, just cover it before exposure to the sun.
	4. Lay out your design and pin the design.
	5. Take your work out in the Sun. The ideal time is from 11:00 am to 2:00 pm on a sunny day.
	6. When your blue print turns a little darker than the shade you want your blueprint to be, take it out of the Sun and remove all the designs. Rinse the fabric several times until the water runs clear. 7. Let your blueprint dry and enjoy your creation.
	Safety--Wear your goggles, rubber gloves, and particle dust mask. When rinsing, wear rubber gloves.
Kathy Kitzmann (Mercy High School, Farmington Hills, MI) has used this activity with her students. She has written a different Recipe Page [Word \| Acrobat] and a Handout [Word \| Acrobat].
	References: Blueprints on Fabric -- Innovative Uses for Cyanotype by Barbara Hewitt, Interweave Press, Inc., 201 East Fourth Street, Loveland, CO 80537 ISBN 0-934026-91-2

2009-01-27T19:48:00.000-08:00

1). Atom Radius
 It is the distance from the atomic nucleus to the electrons out in the skin.
 The fingers atom is influenced by the size of the atomic number element.
 The number of atomic elements multitude, the more the number of skin electron, so that the larger the fingers atom.
So: in one group (from top to bottom), fingers atom bigger.
 In one period (from left to right), increasing the number atom means increasing cargo core, while the number of skin electron remain. As a result of pulling the core electrons the big, causing the small fingers atom.
So: in one period (from left to right), spoke atom the less.

2). Ion radius
 Ion fingers have a significantly different (significant) compared with the fingers neutral atom.
 Ion be positive (cation) have fingers of smaller, whereas had a negative ion (anion) have fingers larger if compared with the fingers netralnya atom.

3). Ionization energy (kJ.mol-1)
 It is the minimum required energy neutral atoms in gas form to release one electron so that the ion had formed +1 (cation).
 If the atom is to deliver the second electron will be required a greater energy (called the second ionization energy), dst.

Group

2009-01-27T19:44:00.001-08:00

Group
 periodic system consisting of 18 vertical columns, called the
 There are 2 ways for naming groups:
a) the System 8
According to this way, periodic system is divided into 8 groups, namely the main (class A) 8 and the transition (group B).
b) the System 18
According to this way, periodic system is divided into 18 groups, namely the 1 to 18, starting from the far left column.
 elements that have the same valence electrons are placed on the same.
 For the elements of a location in accordance with the periodic system:
number of Group = number of valence electrons

 elements A has the name that is:
a. Group IA = Alkali
b. The group IIA = Alkali Land
c. The group IIIA = Boron
d. The IVA = Carbon group
e. The VA = Nitrogen
f. The group VIA = Oxygen
g. Group VIIA = Halida/Halogen
h. = Group VIIIA = Inert Gas

Period

2009-01-27T19:37:00.000-08:00

Period
o It is a horizontal row-row on periodic table.
o Modern SPU consists of period 7. Each period the number of states / atom skin many elements that occupy the period-period.
So:

Period number = Number of Skin Atom

o The number of elements in each period:

Period

Total Element

Atom ( Z )

1

2

1 – 2

2

8

3 – 10

3

8

11 – 18

4

18

19 – 36

5

18

37 – 54

6

32

55 – 86

7

32

87 – 118
Note:
a) Period 1, 2 and 3 called for a short period of relatively few elements
b) Period 4 and so-called long period
c) The period of 7 is called the period not yet complete because it has not come to class VIII A.
d) For a number of elements based on a number atomnya, you only need to know the atomic number elements that start each period

o elements that have 1 skin (skin K only) is located in the period 1 (line 1), the elements that have 2 skin (skin K and L) located in the period 2 .

Example :

9F = 2 , 7 period 2

12Mg = 2 , 8 , 2 period 3

31Ga = 2 , 8 , 18 , 3 period 4

Periodic System Elements

2009-01-27T19:29:00.000-08:00

1). The basis of the Non-Metal and Metal
1. raised by Lavoisier
2. The still is very simple, because the elements of the metal itself there are still many differences.

2). Law Triade Dobereiner
 raised by Johan Wolfgang Dobereiner (Germany).
 elements are grouped into three groups of elements called Triade.
 Dasarnya: similarity physics and chemical nature of these elements.

Kins of Triade :
a.Triade Litium (Li), Natrium (Na) dan Kalium (K)

Elelment Atom Mass Wujud
Li 6,94 Padat
Na 22,99 Padat
K 39,10 Padat

b. Triade Kalsium ( Ca ), Stronsium ( Sr ) dan Barium ( Ba )
c. Triade Klor ( Cl ), Brom ( Br ) dan Iod ( I )

3). Law octave Newlands
 raised by John Newlands (UK).
 elements are grouped based on the relative increase in atomic mass (Ar).
 Element-8 to have the nature of chemicals similar to the first element, the element to-9 has a similar nature with the elements to-2 ff.
 nature-nature element that is found after a regular or periodic elements called Law 8 octave.

H Li Be B C N O
F Na Mg Al Si P S
Cl K Ca Cr Ti Mn Fe

4). Mendeleev Periodic System (Short Periodic System)
 Two chemist, Lothar Meyer (Germany) and Dmitri Ivanovich Mendeleev (Russia) based on the principles of Newlands, perform classification elements.
 Lothar Meyer prefer the nature chemical elements of Mendeleev more while the increase in atomic mass.
 According to Mendeleev: attributes element is a function of periodic relative atomic mass. This means: if the elements are developed according to the relative increase in atomic mass, the specific nature will be repeated periodically.
 elements that have similar attributes are placed in a vertical column, called Group.
 The horizontal row, for the elements based on the relative increase in atomic mass and is called Period.

5). Periodic System Modern System (Periodic length)
• presented by Henry G Moseley, who argued that the nature of a periodic function of the number atomnya.
• This means: a basic element of nature is determined by the number atomnya not by its relative atomic mass (Ar).

Link Learning Materials Chemistry with Animation

2009-01-27T19:27:00.000-08:00

Link Learning Materials Chemistry with Animation

In addition to some that I've met and I've put in the column next to this blog (Link Chemical) and also on the previous posting, here are some links that I find new subjects related to chemical materials, be very suitable material for study or for high school students students at the beginning of the semester chemistry department or the other.

1. http://www.mpcfaculty.net/mark_bishop/default.htm
Contain chemicals that lessons are presented in the form of presentation files, animation, pdf files.
2. http://lc.brooklyn.cuny.edu/smarttutor/
Not only the chemical materials in this site are provided, but also other available materials Sain and also mathematics and computer programming and other materials. Not too deep, but quite the interesting things, including interactive learning model.
3. www.chemmybear.com
At this link can be found some chemical materials from the general chemical study, which tersajikan in the form of animation, guide the experiment, study card in pdf format and some other links to study chemistry.
4. www.chem.ox.ac.uk / vrchemistry
Virtual Chemistry, this is a virtual learning chemistry, complete with a tutorial in the form of a movie that can be interactive with the view that is quite interesting and not boring.
5. http://ull.chemistry.uakron.edu/genchem/animations/index.html
Not too full, but can be a complement:). Served in the mov file. Some link it can not work.
6. http://chemmovies.unl.edu/ChemAnime/index.htm
Simple and interactive, contains a discussion of acid base, attractive style, orbital atom, atomic structure, electron chemistry, molecular structure, solvent, solid.

From which I write all over my coffee / copy it to disk so that it can be used offline (no need to connect to the Internet). To be able to open the perfect sometimes necessary software shockwave, and / or Quicktime, and / or java, atausoftware another type according to the files on each site. If interested to copy the contents of the website link above I recommend using HTTrack website copier, good for under windows or linux available.

Indeed, for some time that the website is needed for a long time to download. But it's not a problem if we have a fairly wide bandwidth. For those who do not have enough time or access the Internet that I can kopikan sufficient to CD or DVD (just replace the cost and tell pengkopian chip to CD / DVD). May be used for additional or alternative materials in the classroom lessons. Who are interested (seriously) can leave a message through the field posting a comment below or send to my email address: mr dot Urip at gmail dot com. Or can also through a short message to 081349 224050.

Sorry edges so well: D, but simply to help fellow chemistry teacher or student-student may not have that and are very need.

Mass relative molecular (Mr)

2009-01-26T09:03:00.000-08:00

Mass relative molecular (Mr)
• It is the comparison between the mass of a molecule with a standard.
• The relative molecular mass (Mr) = amount of a substance relative atomic mass (Ar) of the atom-atom molecular oxygen in it.
• Special compound to be used ion term Relative Formula Mass (Mr) because the compound does not consist of ion molecule.
• Mr = Sum of Ar
Example:
Note: relative atomic mass (Ar) H = 1; C = 12, N and O = 14 = 16.
What is the relative molecular mass (Mr) of CO (NH2) 2
Answer:
Mr CO (NH2) 2 = (1 x Ar C) + (1 x Ar O) + (2 x Ar N) + (4 x Ar H)
= (1 x 12) + (1 x 16) + (2 x 14) + (4 x 1)
= 60

MASS RELATIVELY ATOM (Ar)

2009-01-26T09:00:00.000-08:00

Ø It is the comparison between the atomic mass of 1 atom of the other.

Ø In general, some elements of the isotope in the determination of relative atomic mass (Ar) is used the average mass of the isotopes

Ø According to IUPAC, used as a benchmark atom C-12 is 1 / 12 the mass of 1 atom C-12;

ELECTRON CONFIGURATION

2009-01-26T08:55:00.000-08:00

ELECTRON CONFIGURATION

 electron distribution in the skins atom called configuration.
 Skin's first atom (the most close to the core) is given the symbol K, skin-2 to be the symbol of L
 The maximum number of electrons on the skin of each meet 2n2 formula (n = number skin).
Example:
Skin K (n = 1) a maximum of 2 x 12 = 2 electrons
Skin L (n = 2) a maximum of 2 x 22 = 8 electrons
Skin M (n = 3) maximum of 2 x 32 = 18 electrons
Skin N (n = 4) a maximum of 2 x 42 = 32 electrons
Skin O (n = 5) a maximum of 2 x 52 = 50 electrons
Note:
Although the skin O, P and Q can accommodate more than 32 electrons, but the fact is skins are not fully used.

Steps Writing electron configuration:
1. Skin skin-filled skin from K and L
2. Especially for the primary (class A):
The number of skin = number of
The number of valence electrons = group number
3. The maximum number of electrons on the outside skin (valence electrons) is 8.
o valence electrons have a role in the formation of bonds between atoms in a compound.
o The nature of a chemical element is also determined by electron valence. Therefore, the elements that have the same valence electrons, will have a similar chemical nature.
4. For the main elements (group A), electron configuration can be determined as follows:
a) Some may be filled with full leather electrons.
b) Determine the number of electrons remaining.
 If the number of electrons remaining> 32, filled with leather next 32 electrons.
 If the number of electrons remaining <32, the next skin filled with 18 electrons.
 If the number of electrons remaining <18, the next skin filled with 8 electrons.
 If the number of electrons remaining <8, all the electrons on the skin next.

ISOTOPE, ISOBAR AND ISOTONIC

2009-01-26T08:53:00.000-08:00

1). Isotope
Is the atom-atom from the same elements that (have the same atomic number) but different numbers mass.
Example:

2). ISOBAR
Is the atom-atom from a different element (atom has a different number) but have the same mass number.
Example:

3). ISOTONIC
Is the atom-atom from a different element (atom has a different number) but have the same number of neutrons.
Example:

ION STRUCTURE

2009-01-26T08:49:00.000-08:00

ION STRUCTURE
 An atom can lose / release electrons or get / receive additional electrons.
 Atom a loss / release electrons, will be a positive ion (cation).
 get the Atom / receive electrons, will be a negative ion (anion).
 In the Ion, which changed only the number of electrons only, while the number of proton and neutron remain.

Spesi Proton Electron Neutron
Atom Na 11 11 12
Ion 11 10 12
Ion 11 12 12

General formula for calculating the number of proton, neutrons and electrons:
1). To nuklida neutral atom:
: P = Z
e = Z
n = (A-Z)

2). To nuklida cation:
: P = Z
e = Z - (+ y)
n = (A-Z)

3). To nuklida anion:
: P = Z
e = Z - (-y)
n = (A-Z)

MASS NUMBER

2009-01-26T08:47:00.000-08:00

NUMBER MASS
 Indicates the number of proton and neutrons in the atomic nucleus.
 Proton particles and neutrons as the composer called the atomic nucleus Nucleon.
 nucleon in the number of atomic elements is expressed as a number Mass (given the symbol of the letter A), so that:
A = mass number
= Number of proton (p) + number of neutrons (n)
A = p + n + n = Z
 The single atom with atomic number on the left of the mass and number in the top left of the symbol of the atom. This kind of notation called Nuklida.

Description:
X = = A symbol of atomic mass number
Z = atomic number Example:

ATOM NUMBER

2009-01-26T08:46:00.001-08:00

ATOM NUMBER
 says the number of proton in the atom.
 For neutral atoms, the number of proton = number of electrons (atomic number also says the number of electrons).
 be symbols letter Z
 Atomic release of electrons becomes a positive ion, which receives the electron becomes a negative ion.
Example: 19K

5). Modern Atomic Model

2009-01-26T08:45:00.001-08:00

5). Modern Atomic Model
Developed based on the theory of quantum mechanics is called wave mechanics; initiated by 3 experts:
a) Louis Victor de Broglie
States that the material has a duality that is the nature of the material and as a wave.
b) Werner Heisenberg
Propose the principle of uncertainty for the material that is as particles and waves. Distance or location of the electron-electron encircle the core can only be determined with the possibility - the possibility only.
c) Erwin Schrödinger (perfect model Bohr Atom)
Successfully preparing for the electron wave equation by using the principle of wave mechanics. Electron-electron that encircle the core there is an orbital in the 3-dimensional area around the nucleus where the electrons with a certain energy can be found with the largest possible.

Modern atomic model:
a) Atom atom consists of a core containing a proton and neutrons, while electron-electron atoms move around the core and are on the orbital-orbital particular form of skin atom.
b) Orbital 3 dimensions, namely in the area around where the core electrons with a certain energy can be found with the largest possible.
c) The position on the orbital electrons orbital-expressed with the quantum.

Orbital

 Orbital described as the electron cloud: forms a space in which electrons likely found.
 The meeting electron cloud then the more likely found electrons and vice versa.

4). Niels Bohr Atom Model

2009-01-26T08:42:00.000-08:00

4). Niels Bohr Atom Model
• Model atomic based on quantum theory to explain the spectrum hydrogen gas.
• According to Bohr, the spectrum shows that the electrons only occupy the high-level atomic energy in particular.
According to him:
a) consists of a core Atom had a positive and circulating around the electron-electron be negative.
b) the electrons circulating in the atomic nucleus beset particular orbit, known as the movement of stationer (still) hereinafter referred to as the main energy level (skin electron) stated that the main quantum number (n).
c) During the electron is in orbit stationer, energy will remain, so there is no light emanated.
d) electrons can only move from the track stationer to a lower trajectory stationer higher if absorb energy. Conversely, if the electrons move from the track stationer higher to lower energy release occurred.
e) In normal circumstances (without outside influence), electrons occupy the lowest energy level (called a basic level = ground state).

The weakness of the Niels Bohr Atom Model:
1. Can only explain the spectrum of the atom or ion containing one electron and not in accordance with the atom or ion spectrum of manelectrony.
2. Not able to explain that the atom can form molecules through chemical bonds.

Rutherford Atom Model

2009-01-26T08:41:00.000-08:00

Rutherford Atom Model
a) Rutherford found evidence that there are atoms in the core of the atom be positive, the size smaller than the size of atoms and atomic mass is almost entirely derived from its core mass.
b) Atom atom consists of a core that had positive and are on the center atom and the electrons move across the core (such as a planet in the solar system).
c) Atom are neutral.
d)'s finger-core atom and fingers are atom can be determined.

Weakness Model Rutherford Atom:
 The failure to explain why electrons do not fall due to the core of atom style electrostatics pull the core of the electron.
 According to the Maxwell theory, if the particles as electrons orbit the nucleus had a cargo that has the opposite orbite will involute and will lose the power / energy in the form of radiation so that it eventually fell to the core.

2). Thomson Atom Model

2009-01-26T08:39:00.001-08:00

2). Thomson Atom Model
After finding electron by JJ Thomson, Thomson disusunlah atom model that was a refinement of the model atom Dalton. According to Thomson:
a) Atom consists of materials and be positive in scattered electrons (like raisins in raisin bread)
b) Atom are neutral, that is positive and the cargo load negative same amount

ATOM DEVELOPMENT THEORY

2009-01-26T08:37:00.000-08:00

ATOM development THEORY

1). Dalton Atom Model
a) Atom is described as pejal the ball is very small.
b) Atom is the smallest particles that can not be parsed again.
c) Atom has the same elements of a similar nature, while the atom different elements, different in mass and nature.
d) if the compound form atom join with each other.
e) Reaction chemistry is just reorganization of atom-atom, so there is no atom has changed due to chemical reactions.

Figure Model Atom Dalton

Dalton atomic theory is supported by 2 natural law, namely:
1. Legal Immortality Massa (Lavoisier law): mass substances before and after the reaction is the same.
2. Law Comparison Stay (legal Proust): comparison of mass elements composing a substance that is permanent.

The weakness of the Atomic Model Dalton:
1) Can not explain the difference between the elements of a single atom with the other elements
2) Can not explain the nature of the power of the materials
3) can not describe how each atom-atom berikatan
4) According to the theory of atomic Dalton number 5, have not changed due to the heavy chemical reaction. Now that chemical reaction with nuclear, an atom can be changed to another atom.

Dalton's Atomic Theory

2009-01-26T08:36:00.001-08:00

Basic Postulat Dalton's Atomic Theory:
1) Any material consisting of particles called atoms
2) element is the material that consists of a kind of atom
3) Atom is an element identical but with different atomic elements other (have different masses)
4) compound is a material that consists of 2 or more types of atoms with a specific comparison
5) Atom can not be created or destroyed and can not be changed to another atom through chemical reactions normal. Chemical reactions are just re-regulation (reorganization) atom-atom involved in the reaction is

The weaknesses of the theory Atom postulat Dalton:
1) Atom is not something that is not divided, but consists of particles subatom
2) Atom-atom from the same element, can have a different mass (called isotope)
3) Atom of an element can be changed into other elements through nuclear Nuclear Reaction
4) Some of the elements does not consist of atom-atom molecule-molecule but

ATOM STRUCTURE

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Particles MATERIALS
The smallest particles of the materials mentioned.
Some of the particles in the material:
1. According to Democritus, the materials are discontinuous (if the material is divided and continues to be divided and eventually have the smallest particles can not be divided again = called Atom)
2. According to Plato and Aristoteles, a continuous distribution of materials (the division can continue without limit)

Chemical formula

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Declares the type of atom and the relative amount of a substance that sort.
Be divided into 3:
a. Molecular formula
Declares the type and number of atoms of an oxygen molecule prepare
Example: molecular formula of water (H2O)

b. Chemical formula compound Ion
Declares the type and number of atoms that arrange a compound ion
Characteristic of compound ion is one of the composer atom compound are metal (located in front)
Example: Mg (NO3) 2; BaCl2; CuSO4; NaCl

c. Empirical formula
The formula is also called, and the type of claim is a simple comparison of the atom-atom in a compound
Example: Etuna with molecular formula C2H2 and have empirical formula CH
The formula is a chemical compound ion empirical formula
Example: salt (NaCl)

Basic Materials particles

2009-01-26T08:30:00.001-08:00

Can be:
1) Atom
 Atom is the smallest particle of an element that still has elements of nature that
 Atom elements are given a badge with the symbol of the element
 Example: Na, Mg, Ba, Ca, Fe

2) molecule
 molecule is neutral particles that consist of 2 or more atoms, the atom and the atom is a kind of different.
 molecule that consists of a kind of atom molecule called Element
 molecule that consists of atom-atom molecule called a different compound
 Example: H2O, CO2; H2SO4

3) Ion
 Ion is the atom or group of atoms that had electricity
 Ion who had called Kation positive, while the negative ion which had called anion
 Ion consists of 1 atom called Ion Single (monoatom), while the ion consists of 2 or more atoms called Ion Poliatom
 Examples:
Single Kation: Na +, K +
Kation Poliatom: NH4 +, H3O +
Single anion: Cl-, S2 -
Anion Poliatom: NO3-,-OH

Element particles (atom can form, can be a molecule)
a. In general, every element, including metal particles have a form of Atom
b. Only a few non-metallic elements that form partikelnya molecules (eg hydrogen H2; fosforus P4; sulfur S8)
c. Molecules that consist of 2 atom molecule called Diatomik (eg, molecular hydrogen, nitrogen)
d. Molecule that consists of more than 2 atom molecule called Poliatomik (eg molecular fosforus, sulfur)

Compound particles (can be a molecule, can form ion)
o Can a molecule (called a molecular compound) or Ion (compound called Ion)
o compound of metal ion including the compound, while the compound of the non-metallic elements including molecular compound.
Example molecular compound: water (H2O); compound ion: Calcium carbonic (CaCO3)

Compound

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 compound chemical form by commitments from two or more types of elements.
 The nature of a compound with different elements of nature redactor.
For example: compound H2O (l) and NaCl (s)

Element

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A. ELEMENT
• Element is a single substance that can not be explained again in a chemical oxygen-oxygen other, more simple.
• Element is a single substance from the most simple materials.
For example: H, C, N, P, Fe, Au, Mg

o Symbol Element (Symbol Atom)
According Jöns Jakob Berzelius (Sweden):
 Each element is represented with a letter from the initial name of the Latin elements and is written with capital letters / capital.
 elements that have the same initial letter, symbol be distinguished by adding one letter of the name of Latin element, which was written with lowercase letters.

Working in the laboratory techniques

2009-01-26T08:22:00.000-08:00

o The handling of the chemicals:
a) Avoiding direct contact with chemicals
b) Avoid direct steam to smell chemicals
c) Use of gloves

o If you want to move the chemicals:
a) Read the label chemicals (at least 2 times)
b) Move to the amount required
c) Do not use excessive
d) If any of the rest, do not return to the chemicals in the bottle to prevent contamination
e) Using the tools that are not corrosive chemicals to move the compact
f) For liquid chemicals, move carefully in order not to spill

o If the material is exposed to chemicals:
a) Be quiet and do not panic
b) Request the help of friends near you
c) Clearing of the direct contact (washed with clean water)
d) Do not scratch the skin is exposed to chemicals
e) Go to a place that is oxygen
f) Contact the paramedics as soon as possible

o The problem of handling chemical waste:
a) the form of waste organic matter should be removed in order to separate recyclable
b) The liquid waste is not dangerous but can be removed should be diluted with water prior secukupnya
c) Waste liquid that does not dissolve in water and toxic waste must be collected in a bottle and container labeled
d) solid waste must be separate because it can clog the water channel
e) Soap, detergents and liquids are not dangerous in the water can be removed through the channel of water and rinsed with dirty water secukupnya
f) Use substances / chemicals secukupnya

Preparation work in the laboratory

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Preparation work in the laboratory:
1. Plan your experiment to be conducted before the start practicum
2. Using work equipment (glasses, practicum jacket, gloves and closed shoes)
3. For women who long, mandatory binding hair
4. Prohibited from eating, drinking and smoking
5. Keep the table practicum and environmental laboratory
6. Sink to wash hands with soap and water, especially after practicum
7. When skin is exposed to chemicals, not touch not to spread
8. Ensure that tap gas would not leak when using bunsen
9. Make sure that the tap water is always in a closed before and after doing practicum

Materials Chemistry

2009-01-26T08:19:00.001-08:00

Based on the type of chemicals are:
a) unstable (explosive)
b) pengoksidasi (oxidizing)
c) karsinogenik (carcinogenic: trigger the occurrence of cancer cells)
d) dangerous for the environment (dangerous to the environment)
e) up easily (flammable)
f) poisonous (toxic)
g) corrosive (corrosive)
h) causing irritation (irritant)

Development of Chemistry

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1) Around 3500 years BC, in Ancient Egypt already practicing chemical reactions (eg how to make wine, the corpse preservation). 2) In the 4th century BC, the Greek philosophical and Democritus the Aristoteles try to understand the substance material. According to Democritus o = every material consists of small particles called atoms. According to Aristoteles o = material form of 4 types of elements, namely: land, water, air and fire. 3) mid-Century (years 500-1600), which dipelopori by chemical experts Arabic and Persian.  Chemical lead to more practical terms. Produced various types of substances such as alcohol, arsen, zink acid iodida, acid sulphate and nitrate acid.  Name chemistry of birth, of the words in Arabic (al-kimiya = material changes) by the Arab scientist Jabir ibn Hayyan (700-778 years). 4) to the 18 Century, the term appears Modern Chemistry. Dipelopori by French chemist Antoine Laurent Lavoisier (years 1743-1794) revealed a successful legal immortality mass. 5) In 1803, a British chemist named John Dalton (year 1766-1844) put the theory of atoms for the first time. Since then, the chemistry continues to grow rapidly at this time.
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Branchs Chemistry

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Branchs Chemistry
Include:
1) Chemical Analysis
= Learn about the analysis of chemicals found in a product.
2) Physical Chemistry
= Studies focus on determining the form attached to the energy of a chemical reaction, the nature and substance fisis changes chemical compound.
3) Organic Chemistry
= Learn the chemicals found in living things.
4) inorganic Chemistry
= Opposite of organic chemistry; learn things off.
5) Green Chemistry
= Learn about everything that happens in the environment, particularly related to environmental pollution and how penanggulangannya.
6) Chemistry Core (Radiokimia)
= Learn substance-radioactive substances.
7) Biokimia
= Branch of chemistry that is closely related with the biological sciences.
8) Food Chemistry
= Learn how to improve the quality of food.
9) Pharmaceutical Chemistry
= Studies form the focus of research and development of materials that contain drugs.

Benefits of Learning Chemistry

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Benefits of Learning Chemistry
Include:
a. We become more understanding of the nature of good and different processes that took place in it.
b. Have the ability to process natural materials into a useful product for people.
c. Help us in the framework of the formation of attitudes.

• In particular, the chemistry has very important role in the areas of: health, agriculture, livestock, law, biology, geology and architecture. (A role in the field of chemistry-field!)

• Behind the great contributions to our lives, must be honestly acknowledged that the development of chemistry also provides for the negative impact of human life. (Say for example!)

Metallic Copper + Concentrated Nitric Acid

2009-01-26T07:54:00.000-08:00



I adapted this lecture experiment from the "Coin-Operated Reaction" developed by Ron Perkins. I have made video records of a few lecture demonstrations, not to replace doing them in class, but rather so students can go back and review the experiments later in the year. I wanted an interesting system with which to make observations on the first day of class. I also wanted to give students an opportunity to explain what they saw in a logical manner. This system is rich with ideas such as gas laws, air pressure, complex ions, oxidizing acids, oxidation reduction, and gas solubility in water. I have the gas bubble into an aquarium since I have no hood in my classroom. The first time I tried this set-up, I was surprised when the gas cooled and the water was pushed back into the reaction flask. Here is a three-minute streaming video version of the lecture experiment (Real Audio plug-in required). Note: there is sound, but we are not saying anything. We took this movie with a digital camera and then our Stage Crew chief (James Jontz) made the 700 mb video into a 7 mb streaming video using his Macintosh (technically unexplainable by me).
The gas is bubbled into the aquarium so it won't get into the air. I also place a wet cloth over the opening of the aquarium.	What You Do... To a 500 mL florence flask, add about 50 mL of concentrated nitric acid. Place a coil of copper wire into the acid and stopper with a one-holed rubber stopper fitted with a long tube. The end of the tube is placed into a large container filled with water. (I use a plastic aquarium.)[1]

As soon as the Cu^o contacts the HNO₃(conc) the red-brown NO₂(g) forms.	What You See... Many changes occur during this demonstration. When the copper wire (or use two pre-1982 pennies) is added to the colorless nitric acid, the solution turns green and a large amount of red-brown gas is formed. The air being displaced by the gas formation can be seen bubbling through the water. The flask gets VERY warm. When enough gas is formed, it bubbles through the water (keep the liquid stirred so most of it will dissolve. The gas that makes it to the top is noxious.
When the water siphons back into the flask, the blue Cu(H₂O)₄²⁺ forms.	Later... The gas in the flask begins to cool and therefore contracts. (I am reading my Ira Remsen story and allow students to notice the change.) As the pressure inside the flask decreases, the outside air pressure begins to push the water back toward the original flask. In addition, the red-brown gas dissolves in the water. Eventually, the water rushes into the flask, the solution turns characteristic blue, and the red-brown gas disappears as it is dissolved.

The Set Up...

Equipment:

500 mL florence flask
ring stand
large ring placed below the flask
small ring that fits over the neck of the flask
one-hole rubber stopper
60 cm glass tubing
large container of water

The glass tube is bent in such a way as to connect the top of the flask with the bottom of the water container. The water in the container can be stirred by hand or with a magnetic stirrer. It needs to be stirred, however, or else the NO2 gas collects above the liquid (as it did when I took this picture...whew!).

The Chemistry...

Oxidation of copper metal with a strong oxidizing agent, conc. nitric acid.: In a classic experiment, copper metal is turned into copper(II) ion while the nitrogen(V) in the nitrate ion becomes nitrogen(IV) in the nitrogen dioxide gas.
Charles' Law: As the temperature from the reaction warms the gas, it expands. Later, as it cools, the gas contracts.
Nonmetal oxides are acid anhydrides (also link to acid rain): Although the nitrogen dioxide gas is noxious and toxic, it dissolves readily in water and make the solution acidic. This can be shown by adding a little indicator to the water and making the water slightly basic before the copper is added to the acid.
Air pressure: As the pressure in the flask is decreased as it cools, the outside pressure pushes the water up the tubing toward the flask. The nitrogen dioxide gas is not pulling the water in.
Descriptive chemistry--copper solutions are green and blue: The colored solutions come from complexes of copper(II) ion in solution. Aqueous copper ion is blue, Cu(H₂O)₄²⁺ The green must be copper surrounded by nitrates(?)

Discussion...

During this demonstration I read the reminiscence by Ira Remson quoted in Bassam Shakhashiri's demonstration book.[2]

While reading a textbook of chemistry I came upon the statement, "nitric acid acts upon copper." I was getting tired of reading such absurd stuff and I was determined to see what this meant. Copper was more or less familiar to me, for copper cents were then in use. I had seen a bottle marked nitric acid on a table in the doctor's office where I was then "doing time." I did not know its pecularities, but the spirit of adventure was upon me. Having nitric acid and copper, I had only to learn what the words "act upon" meant. The statement "nitric acid acts upon copper" would be more than mere words.

All was still. In the interest of knowledge I was even willing to sacrifice one of the few copper cents then in my possission. I put one of them on the table, opened the bottle marked nitric acid, poured some of the liquid on the copper and prepared to make an observation. But what was this wonderful thing which I beheld? The cent was already changed and it was no small change either. A green-blue liquid foamed and fumed over the cent and over the table. The air in the neighborhood of the performance became colored dark red. A great colored cloud arose. This was disagreeable and suffocating. How should I stop this?

I tried to get rid of the objectionable mess by picking it up and throwing it out of the window. I learned another fact. Nitric acid not only acts upon copper, but it acts upon fingers. The pain led to another unpremeditated experiment. I drew my fingers across my trousers and another fact was discovered. Nitric acid acts upon trousers. Taking everything into consideration, that was the most impressive experiment and relatively probably the most costly experiment I have ever performed... It was a revelation to me. It resulted in a desire on my part to learn more about that remarkable kind of action. Plainly, the only way to learn about it was to see its results, to experiment, to work in a laboratory.[3]

Safety and Disposal...

The solution is highly acidic. I pour it out into a large beaker or battery jar and add excess sodium carbonate. The carbon dioxide bubbles indicate neutralization and the resulting copper carbonate precipitate is filtered, placed in a baggie and thrown away. The neutralized filtrate can be disposed of as you would any simple salt solution. Procedures may vary from location to location.

References:: [1] This demonstration is based on one shown by Ron Perkins called the "Coin Operated Demonstration"; [2] Shakhishiri, B.Z. "Chemical Demonstrations Volume 1--A Handbook for Teachers of Chemistry"; The University of Wisconsin Press: Madison, Wisconsin, 1983; [3] Getman, F.H. "The Life of Ira Remsem"; Journal of Chemical Education: Easton, Pennsylvania, 1940; pp9-10.

CHEMICAL SCIENCE

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1.1 Scope Chemistry
Definition:
- In short, Chemistry is the science of engineering materials that is a material change into another material.
- The full, Chemistry is a science to learn about:
a. The order covers materials = components forming material, and a comparison of each
component.
b. Structure includes the structure of the material particles of a material or it illustrates how
atom-atom redactor material mutual berikatan.
c. The nature of the material covering the nature fisis = (shape and appearance) and the
nature of chemicals. The nature of the material is influenced by: the order and structure of
materials.
d. The changes include changes in the material = fisis / physics (form) and chemical changes
(produce new substances).
e. Energy changes that accompany the material = the number of energy that accompanies a
number of materials and the origin of that energy.

 Chemistry developed by chemical experts to answer the question "what" and "why" about the nature of the material that is in the nature.
 Knowledge born of the effort to answer the question "what" is a fact that is: nature of the material observed by the same person will result Descriptive Knowledge.
 Knowledge born of the effort to answer the question "why" of a material nature have a certain theoretical knowledge will result.

LA	DH_f(LA)	LB	DH_f(LB)	DH_f(adduct)	DH_rxn

BF₃	-272.15	NH₃	-7.28	-291.81	-12.38
BCl₃	-97.01	NH₃	-7.28	-140.14	-35.85
BBr₃	-50.05	NH₃	-7.28	-91.39	-34.06