1

Course Notes – Chemistry A & B Chemistry is called the “central science.” • It connects physical sciences with life and applied

sciences. • Without understanding of chemistry, there is no

physics, no astronomy, no biochemistry and no material science.

• Chemistry is about the food that you eat, the shampoo you wash your hair with, the clothes you wear, how you cook your food and how your house was built.

History Chemistry is as old as the human race. Even so, new findings keep popping up every day. We’re going to start at the beginning. • The word “atom” comes from the Greek term

“atomos” which means “undivided.” • Democritus, an ancient Greek philosopher, lived

between about 460 to 370 BC. He is considered the first atomist – a practitioner of the atomos school. He is credited with the atomos theory, that all matter could be divided into tiny particles.

• Atomos theory taught that the particles had

2

different characteristics and combined in different combinations.

• Only four elements were proposed by the atomists: earth, wind, fire, water. They believed the combinations gave all the substances in the world.

• The idea of four elements is found in most eastern cultures including Chinese, Hindu and Greek.

• To these four, a mystical fifth element eventually was added -- the “quintessence.” “Quint-“ means “fifth,” and “-essence” referred to an intangible, “vital” something that “animated” substances. To the ancients, the quintessence was why heat was released when wood burned, and why life existed. It became mixed with mysticism including witchcraft, religion and secret rites, including the search for the “philosopher’s stone,” and the “Holy Grail.” The word remains with us today.

• The belief in “combinations” of the elements eventually led to the rise of “alchemy,” which was classified as a ”black art” – if the practitioner was unable to use his science for the benefit of his patron, usually a king.

• The main goal of alchemy was turning “base” substances like lead into gold – transmuting. This seemed plausible -- if there were four

3

“elements” and the difference between lead and gold, for example, was just the quantity of each element mixed. One of the oldest transmutations was soap-

making. It is a process believed to be several thousand years old. Lets see what happens when we try it today. We will take a mixture of oil or fat, mix it with an alkaline material (either a pure chemical like sodium hydroxide -- also called “lye” – or an impure source like wood ashes) and turn it into soap.

If we start with impure fat, we boil it with water to take out water-soluble substances such as proteins. Two layers will form, with the fat on top of the water layer (Why is that?). When cool, the fat solidifies. If we are using oil, it will be pretty pure to begin with and probably don’t need the washing step.

We separate off the top layer, either as an oil or solid. This is mixed with the alkali. If we get it from wood ash, we have to purify it. We do that by soaking the ash in water, then filtering the liquid to take out particles. The liquid is now pretty alkaline and might actually have a slippery feel to it if you touch

4

it. It feels slippery because the alkali reacts with some of the fat in your skin making it soap. If we use ash as a source of alkali, it is called “potash,” and the main alkaline metal that it contains is potassium. (Is there a connection between “potash” and the name of the element “potassium?”)

We will boil the fat/alkali mixture to react fully. In times past, people would taste the mixture to know when all the alkali was used up. When the mixture no longer felt like it was cutting the tongue, all the alkali had reacted. We can also use pH indicator in the class. (What is “pH?”) If lye was used, a hard soap results. If potash, a liquid, brown soap was produced. This “soft soap,” as it was called, could be hardened by mixing with table salt, sodium chloride (NaCl). Solid soap would form a layer on top.

• The word “Alchemy,” is derived from the Arabic word “al-kimia.” The word origin extends further to the ancient Greek word “chemia,” and may even have been the root of the Egyptian name for Egypt, “keme.”

• While the alchemists may have been chasing

5

rainbows, the work they did to investigate matter and chemical reactions formed the basis for chemistry as we know it today.

Why didn’t people understand that there were more than four elements? They did – about 2,000 years after Aristotle said that matter was not discrete. • Aristotle lived from 384 to 322 BC. He studied,

wrote and taught logic, physics, biology, mathematics, ethics and medicine, among other fields. He developed huge prestige, especially as he was the tutor of Alexander the Great.

• The life-times of Democritus and Aristotle barely intersected, but Aristotle was aware of the former’s atomos beliefs. He rejected them. Because Aristotle was a very learned man, he was considered a “greater” philosopher than Democritus, so his ideas prevailed. Interestingly, it was because of a mathematician that Aristotle developed his own ideas about matter.

• What is a philosopher? Look at the roots of the word: “philos-“ is from the Greek word for “love” and “-sophic” means “learning” or knowledge. Hence a philosopher was a learned person and is why today we grant the university degree “Doctor of

6

Philosophy,” or Ph.D. • The mathematician was Xeno (About 490 to 430

BC). Xeno set forth paradoxes. (What is a “paradox?”) One was that Hercules could cover half of the distance to his destination in a single bound. That means his first step takes him half-way. The second step takes him another half, or one fourth the distance, then one eighth and so on. If you add up these steps, you see that Hercules gets closer and closer to the destination, but can never reach it. Democritus thought this argument was silly; any fool could see that the end point could be reached. Therefore, he believed, matter was composed of discrete, tiny particles. Nevertheless, Aristotle pointed to the paradox as evidence that matter would go on forever.

• By the mid 1600’s, experimenters – the alchemists, metallurgists, dyers and healers, among others – realized that there were some substances that could not be broken down further; that the combinations postulated by Aristotle and others failed to explain their observations of the real world.

• Some historians suggest that Aristotle (and his followers) did not actually conduct experiments to validate and explain their beliefs. Instead, they

7

conclude that Aristotle arrived at his conclusions by “logic.” This is why for 2000 years people believed that heavier objects fell faster than less heavy ones. Other historians try to make the case that Aristotle did practice “a form” of what we call the “scientific method.” Let’s try something: You know what a leaf looks like, don’t you?

THINK for a minute about what the leaf looks like. Now, DRAW what you think the leaf looks like. Be very careful and as detailed as you can be. Is your drawing full-size or not? If not, what is the scale of your drawing. After you have made your drawing, list what you think are the “attributes” of the leaf. (What is an “attribute?”)

Now, examine an actual leaf. Study it carefully from all angles, close-up and at arm’s length. Draw it. Then, smell it -- perhaps even taste it. Feel it. Crush and tear it. What did you discover? List your findings based on actual observations. What attributes of the leaf do you see that are or are not on your original list?

Next, compare your first drawing to the second. What are the differences? What are

8

the similarities? Write a paragraph or two about your comparison. Include how actually using your senses other than sight influenced your comparison. Would you say one or both of your drawings and lists of attributes were arrived at “scientifically?” If one, which one? If both, please explain your logic.

What are some tests that you could perform on the leaf, if you had other substances or materials to work with? Why would you perform the tests? Could you conceive of the tests if you had not handled and examined the leaf? Explain.

From what you have already done, and what you may have learned in other classes, list what you think are the “steps” in the scientific method. (Keep this list for later, when we discuss the scientific method in writing lab reports.)

The Birth of Modern Chemistry The modern era of chemistry started around 1650. From that date to today, chemistry has dramatically altered and improved life for people around the world. Other species, and our biosphere itself, have also been

9

impacted by chemistry, often in ways that have not been beneficial. Recognition of how important chemistry is, as well as the role humans play in the world, are finally being applied to solving the problems that we have created. Here is a “birth-timeline” of chemists and their most significant discoveries, important to our study of chemistry:

Chemist Born-Died

Known For Note

Robert Boyle 1627 - 1691

Boyle’s law- Volume of gas decreases with increasing pressure.

First “modern” chemist; pioneer of scientific method; roots in alchemy.

Daniel Fahrenheit

1686 - 1736

Thermometer Reliable temperature measurement, defined scale.

Joseph Black 1728 - 1799

Discovered carbon dioxide Identified “latent heat” and foundation of calorimetry

Joseph Priestly

1733 - 1804

Discovered oxygen Championed “phlogistron” theory; Unitarianism founder

Charles-Augustin de Coulomb

1736 - 1806

Developed Coulomb’s law; explains how particles attract or repell each other electrostatically.

Successful physicist; contributions to science including geotechnical engineering.

Antoine Lavoisier

1743 - 1794

Discovered nitrogen; named oxygen

Published first modern chemistry textbook, Traité Élémentaire de Chimie

Jacques Charles

1746- 1823

Charles’ law – Volume of gas increases with increasing temperature at constant pressure

Consulted to Montgolfiere Brothers in France who developed first hot-air balloon

John Dalton 1766 - 1844

Dalton’s 5 postulates First modern scientific description of atomic theory.

Amedeo Avogadro

1776 - 1856

Equal volumes of gases (at same pressure and temperature) contain equal number of particles.

Avogadro’s number named in his honor; 6.02 x 1023 particles per mole.

Joseph Louis Gay-Lussac

1778 - 1850

Water is two atoms of hydrogen to one of oxygen.

Also credited with gas law relating pressure and temperature.

Sir Humphry Davy

1778 - 1829

Isolated elements by electrolysis

Emile Clapeyron

1799 - 1864

Developed the “Ideal Gas Law”

French Engineeer and Physicist

10

Friedrich Wohler

1800 - 1882

Synthesizes urea from inorganic sources.

Disproves “vital life source” to make organic compounds.

Lord Kelvin 1824 - 1907

Establishes that absolute zero exists

Temperature at which all molecular motion stops.

Dimitri Mendeleev

1834 - 1907

First “modern” periodic table

Henry Louis Le Chatlier

1850 - 1936

Le Chatelier’s principle of chemical equilibrium

William Ramsay

1852 - 1916

Discovered noble gases Nobel prize in chemistry

John Joseph Thompson

1856 - 1940

Discovered the electron Discovered isotopes; invented mass spec; Nobel prize; taught Ernest Rutherford.

Robert Millikan

1868 - 1953

Measured the charge on the electron

Oil-drop experiment is one of the most important in science.

Antonius van den Broek

1870 - 1926

Said that atomic number was the number of protons in the atomic nucleus.

Practiced law; Amateur physicist - realized atomic number was related to tangible particles in the atomic nucleus.

Ernest Rutherford

1871 - 1937

The atomic nucleus is a tiny volume of the atom

Discovered radioactive transmutation; Nobel prize.

Gilbert N. Lewis

1875 - 1946

Discovered covalent bond; invented “dot diagrams” in 1916.

Lewis acids and bases; may have been blocked as recipient of Nobel Prize.

Albert Einstein

1879 - 1955

Photoelectric Effect Energy = mass times speed of light squared.

Erwin Madelung

1881 - 1972

Madelung Rule for filling atomic orbitals

Professor at University of Göttingen

Niels Bohr 1885 - 1962

Discovered atomic structure

Awarded Nobel prize in 1922. Was dyslexic.

Henry Moseley

1887 - 1915

Organized periodic table by atomic number

Left graduate studies to fight in British Army in WWI; killed in action in Gallipoli, Turkey

James Chadwick

1891 - 1974

Discovered the neutron in 1932

Crucial for the fission of uranium 235.

Linus Pauling 1901 - 1994

Invented electronegativity scale for atomic bonding

Questionable political activities

• In 1789, Antoine Lavosier published what is

considered the first “modern” chemistry textbook. He listed 34 elemental substances that could not be broken down further by chemical means. Hence,

11

the name “element.” Included were substances he had isolated himself, as well as those discovered by others. His list included oxygen, nitrogen, hydrogen, phosphorus, mercury, zinc, and sulfur.

• John Dalton finally ended the Aristotelian era and the nonsense of transmutation by setting up the postulates on which modern chemistry is founded. In 1803 he published his “five postulates.” They held that (1) All matter is composed of tiny particles that he termed “atoms” after the concepts of Democritus, (2) All atoms of a given element are identical, (3) The atoms of any given element are different in relative weight from those of any other element, (4) Atoms of one element can combine with atoms of another element to form chemical compounds, (5) Atoms cannot be created, divided into smaller particles or destroyed by chemical reactions.

• Dalton was not completely correct, but the proof of that did not come during his lifetime. He did not know that atoms have various isotopes – same number of protons, but different number of neutrons. (In fact, he did not know about the sub-atomic particles that we now know make up atoms. The basic particles are protons, neutrons and

12

electrons, and that is all we will study, but there seem to be others, as well.) The sum of the protons and the neutrons is what we call atomic weight of the atom. The average atomic weight is the relative mass of the element. It is based on the sum of the mass fraction of each isotope. We’ll discuss this in more detail when we talk about isotopes.

• Dalton likewise did not know about nuclear energy and particle accelerators. Today we know we can tear atoms apart, fuse them together and take out parts. But it is true that this does not happen in conventional chemical reactions.

• Dalton DID know that elements could combine with each other in ratios of small, whole numbers. This clearly showed that atoms were discrete, not continuous and is now known as “Dalton’s First Law.” Dalton also knew that the same elements could

sometimes form other compounds. For example, copper and chlorine can form compounds called “cuprous chloride,” with the formula CuCl, as well as “cupric chloride” having the formula CuCl2. This is known as “Dalton’s Second Law.”

13

Periodic Table • Demitri Mendeleev is known as the father of the

periodic table. A Russian chemistry professor, he was writing a new textbook and determined to see if he could organize the elements better. The story is that he wrote information about all, known elements on cards – there were 65 in 1868 – and saw a pattern if the cards were placed in order of increasing relative atomic mass: elements with similar properties, could be arranged in columns. Here is what Mendeleev’s table looked like in the 1898 version of his text book:

There were many omissions because Mendeleev expected new elements to be discovered. In fact, he predicted the mass of some of the unknown elements.

14

• Notice that Mendeleev did not list what we today call the “noble gases.” This is because they had not yet been discovered. Since the noble gases do not normally react with other elements, chemists were unaware that they existed.

• There were still problems with the periodic table. Some of the masses of different elements were so close that chemists were not quite sure if the order was correct. They sensed something was wrong because experimental data suggested it. However, basing the table on mass was all chemists had at this time. Henry Moseley solved the problem. In 1911 Rutherford noted the presence of a

small, charged nucleus within atoms (Standard 1e). He used a radioactive source to shoot alpha particles – today known as the nuclei of helium atoms – at a gold foil target. Most of the particles passed right through the foil and impacted a photographic film behind it. There they released energy causing the plate to develop spots. The intensity of the spots corresponded to the number of alpha particles hitting the plate at a given spot.

Rutherford’s graduate assistants, the story goes, misunderstood what he wanted them to

15

do, arranging the film in a circle. When the film was developed, to everyone’s surprise, some spots appeared near the point from where the alpha particles were fired.

The only way these spots could be explained was if something very, very tiny and very, very dense occupied the center of the gold atom. If it were large, more of the spots would have been reflected 180 degrees; if it was not dense, the energetic alpha particles would have pushed it out of the way.

Eventually, the size of the nucleus was calculated as about one trillionth (1/1,000,000,000,000 or 1.0 x 10-12) the volume of the atom!

• Van den Broek, an amateur Dutch physicist, knew of the Rutherford’s work. In 1911, one month after Rutherford published his findings, he proposed that the “atomic number” was not just a convenient way of ordering the elements, but was exactly the number of positive charges in the atomic nucleus. By 1920 Rutherford began using the term “proton” to describe the entity within the atomic nucleus with a charge equal to and opposite that of the electron discovered by Thompson and

16

identical to the atomic number. Moseley was intrigued by van den Broeks

idea. He used (the newly developed field of x-ray diffraction) to investigate. This means he looked at an image of dots on a photographic plate, and then figured out what sort of structure would scatter x-rays to make the observed pattern. In 1913 Moseley found that the wavelength of the x-rays diffracted by elements was related to the atomic number of that element, not its atomic mass. When elements were arranged by atomic number the confusion disappeared. Thus, although Demitri Mendeleev may have been the “father” of the periodic table, Henry Moseley certainly has to be considered its “favorite son.” Let’s try something to better understand what Moseley did:

Spectroscopy

An ordinary compact disk can be used as a diffraction grating. If a laser is shined at it, the light will bounce off the mirrored surface and can be viewed on a suitable screen. Dots will be seen corresponding to the spacing of the

17

lines on the CD1. Beside the CD and laser, you will need a ruler,

a binder clip, a pen/pencil and graph paper. Two people need to work together. One person is

the “operator.” They mount and move the CD and shine the laser. The other person is the “recorder.” They measure and record distance between the dots on the screen and between the CD and the screen.

1. Place the CD in the binder clip to make a stand.

2. Tape the paper to a suitable backing to form a screen.

3. Stand the CD about 12 inches from the screen, reflective side facing the screen.

4. Shine the laser on the shiny side of the CD and watch where the reflected light hits the screen.

5. How many dots do you see? If only one, adjust the angle of the CD to see more. (You might see two or three.)

6. Mark the location on the screen between two

1 For an “advanced” experiment, use artificial “jewels.” Place them in the laser beam and examine the diffraction pattern. Follow the same instructions for examining diffraction from the CD. In fact, you are being provided with a prepared CD and jewels. Half of the CD (and the jewels) have had the aluminized (reflective) backing removed (How was this done?), and the jewels glued to the clear part of the CD.

18

of the dots. Then measure the distance between them. Label this d1.

7. Measure the distance between screen and CD. This is L1

8. Move the CD forward or back from the screen, or change it’s angle. Shine the laser at it and again mark where dots hit the screen, the distance between the dots (d2) and the distance from the laser to the screen (L2).

9. The equation that relates the distance between the dots (d), the wavelength of the laser light (λ), the length (L) and the grid spacing of the lines on the CD (w) is: w=Lλ/d

For a red laser the wavelength (λ ) is approximately 6.5 x 10-7 m.

10. What is the grid spacing of the CD lines? Does this spacing appear to change when the distance from the screen to the CD changes? What about if the angle of the CD is changed?

11. What is the ratio L1:L2? What about the ratio d1:d2? Is there a relationship?

12. This method of analysis is called “scattering.” How do you think Moseley used scattering of x-rays to conclude that different

19

elements scattered x-rays based on their atomic number?

So far you have seen how the periodic table was developed. This is all part of Chemistry Educational Standard 1, “The periodic table displays the elements in increasing atomic number and shows how periodicity of the physical and chemical properties of the elements relates to atomic structure. “ • You know that the periodic table displays the

elements in increasing atomic number (thanks to the work of Henry Moseley) and why they are not listed by atomic mass. You understand that elements show “periodicity.” That is, the physical and chemical properties of the elements appear to form repeating series on the periodic table. This is the point, after all, that Demitri Mendeleev found. (Standard 1a) But we haven’t talked at all about how the periodic table relates to atomic structure! Before we do, there are a few other things to discuss:

• How are types of elements grouped on the periodic table? For that matter, what groups of elements are there? (Standard 1b)

• Does the arrangement of elements on the periodic table tell us anything about their energy, or

20

stability, or how they can bond to other elements? (Standard 1c, 1d)

• Let’s look at how the elements are grouped. Look at a copy of the periodic table. How many (vertical) columns are there? Compare the number of columns to the number of columns in Mendeleev’s original table. What do you notice? Don’t just scan the tables; actually study them and write down what you see in the modern periodic table compared to Mendeleev’s.

• Most of the vertical columns constitute “families” of elements. This designation is based on similarities of the elements in a given column to other elements in the same column. For example, lithium, sodium and potassium are all very reactive with water. They form solutions that are alkaline, or basic, turning litmus paper blue. They all meet the definition of a metal – that is, they conduct heat and electricity although they may not be especially malleable or ductile. (What do these two terms mean?) For this reason, the metals in that vertical group or “family” are called the “alkali metals.” Later on, we’ll see how their reactivity is due to their atomic structure. Hydrogen is at the top of the column, but not

21

because it is a metal; it is there because the element has a reactivity that is similar to the metals in the group and a structure that causes the reactivity.

• Next to the alkali metals, the second column contains a family that was named by the ancient alchemists. They called these substances “earths” because when they burned, their oxides (element plus oxygen) resembled dirt. They also produced water solutions that were alkaline, although much less reactive than the alkali metals. And they also conduct electricity or heat, so they are metals. The choice was clear: these elements were termed “alkaline earth metals.”

Let’s demonstrate. We will take a strip of magnesium, an element in the second column of the periodic table. 1. Check to see if it conducts electricity using a Volts-Ohms-Milliamp Meter (VOM). Set the VOM to ohms. If the strip conducts electricity, the VOM will register zero resistance. 2. Will it conduct heat? We can gently bend the strip into a “U” shape and hang it into a beaker filled with hot water. Does the end not in the water become warm?

22

3. Now, take the strip out of the beaker. Carefully straighten it – it might break regardless of your care, and if this happens, just use one of the pieces. Put on safety goggles. Then, hold the strip at one end with tongs, and put the other end in a flame. When it starts to burn, hold it over the beaker of water to collect the residue that falls off. 4. What does the residue look like? 5. Try stirring the residue in the water. Does

it dissolve very well? 6. Check the pH of the water. Use a pH meter,

litmus paper, phenolphthalein, or other indicator. Is it acid or basic?

7. What do you think has happened, starting with the magnesium strip?

• Columns 3 through 11, called “transition metals,” do not usually have names given to each column. Column 12 and parts of columns 13 through 16 are termed “poor metals,” or “post-transition metals.” They are called this because the transition metals have electrons that go into what are called “d” orbitals. How many electrons can all these “d” orbitals

23

hold? Each orbital can hold 2 electrons. Look at periods 4 through 7. How many columns are there between 2 and 12? (Don’t get confused by the numbering of “periods” and “columns!”) So, how many individual “d” orbitals are there in each period?

Column 12 elements (Zinc, Cadmium, Mercury and element 112, Copernicium, get the last electron in the “d” orbitals. Because the orbitals are now considered “full,” the elements are sometimes included with elements in columns 13 through 16 as “post-transition” metals. Or not!

Columns 12 through 16 really are confusing. They have elements classified as metals, but also “metalloids” and “non-metals.” Look at this table:

13 14 15 16 2 B - boron C - carbon N - nitrogen O- oxygen 3 Al -aluminum Si – silicon P - phosphorus S - sulfur 4 Ga - gallium Ge – germanium As – arsenic Se - selenium 5 In - indium Sn – tin Sb – antimony Te - tellurium 6 Tl – thallium Pb – lead Bi – bismuth Po – polonium 7 Uut – ununtrium Fl – flerovium Uup - ununpentium Lv - livermorium

Boron, silicon, germanium, arsenic, antimony, tellurium and polonium (light shading) are semimetals. Under the right circumstances, they can conduct an electric

24

current. They are used in electronic circuits. But carbon, a non-metal (dark shading) can also conduct an electric current when it is graphite!

To see how carbon conducts an electric current, draw a line on paper with a soft lead pencil (i.e. – no. 2). Set the VOM to “ohms.” Put one electrode on one end of the pencil line. Then, move the other electrode at the other end, and slowly move it toward the first. What do you read on the VOM scale?

The other elements, aluminum, gallium, indium, thallium, tin, lead and bismuth are metals.

The elements at the bottom of the table are laboratory creations, not fully characterized yet. They PROBABLY have characteristics typical of the rest of the table.

• Column 17 does have a name. It is called “halogen.” The word comes from the Greek words for “salt” and “generate” or “to make.” In other words, a halogen is a salt-former. None of the halogens are metals. Would you expect them to be?

• Column 18 also has a name. It is called “noble gases.” Before you begin thinking that these gases

25

are royalty, let me just remind you that they have a full outer layer – orbital – of electrons. Because of this, the noble gases are stable. They have no reason to enter into a chemical bond to become more stable ---- because they can’t!

• Now look at those two rows of elements at the bottom of the periodic table. These are called the Lanthanide and Actinide series. They follow and take their name from elements 57, Lanthanum, and 89, Actinium, respectively.

I call these elements “Cinderella’s step-sisters.” Do you remember how the step-sisters tried to fit their big feet into the little, dainty glass slipper? They bent their foot so the heel and toe fit, but the rest of the foot bowed up. The Lanthanides and Actinides sort of bow up, too. They have electrons in what we call the “f” orbitals. If the periodic table were stretched out to make room for these two series, it would be 14 columns wider and there would be a gap between elements 21, Scandium, and 22, Titanium, and between elements 39, Yttrium, and 40, Zirconium.

Most often, however, the Lanthanides, along with Lanthanum itself, as well as Yttrium

26

and Scandium are called “rare earths” because mining them is so difficult.

Subatomic particles So far, the periodic table tells us all about itself by what is written on it. We know the names of the elements and the number of electrons each has. Since the electrons have negative charge, and the atom overall is neutral, we know that there are protons to balance the charge. For each electron there is a proton. But atoms other than hydrogen have more mass than can be accounted for by just the sum of the weights of the protons and electrons. In fact, the mass of the electron is just 1/1836th the mass of a proton. • In 1932, James Chadwick, already a well-known

physicist, discovered an uncharged particle within the atomic nucleus. It became known as a “neutron” because it was “neutral” – had no charge. The neutron lacked charge, but it had mass

almost exactly the same as the proton. • Now the picture of the atom was much clearer! The

atomic mass was known from the work of Mendeleev and others; the number of protons was learned from Henry Moseley’s studies, and the electron had been investigated by Sir J.J.

27

Thompson. Thompson’s work was fascinating, and a good

description of it is in your textbook. Here, we will just note that he used a Crooke’s or “cathode-ray” tube in his study of electricity. He found that “particles” were released by the cathodic end of the tube (hence, “cathode ray”) and flowed to the anodic end. These particles apparently carried a negative electrical charge and interacted with matter. They also had the ability to cause phosphors on a screen to glow and could be deflected by both magnetic and electric fields. Thompson discovered that regardless of what the anode was made of, the particles were the same mass. It was others who named the particle “electron.”

• To determine the number of neutrons in an atom, all one needs to do is subtract the atomic number – the number of protons – from the overall atomic mass, the sum of the neutrons and the protons. The mass of the electrons can be ignored.

28

The atomic mass listed on the periodic table is an average. It is usually not an integer because some atoms of an element have more or fewer neutrons than other atoms of the same element. These atoms are termed “isotopes” and were discovered by Thompson.

When a student is asked to find the number of neutrons, given the atomic mass, they must first round off to the nearest whole number. Then, they subtract the atomic number from the rounded mass number. For example, how many neutrons in a gold atom, atomic mass 196.97, atomic number 79? First, round off the atomic mass to 197, then subtract the number of protons (atomic number 79) for the answer, 118.

Since, in a neutral atom, the number of

electrons is equal to the number of protons, students can find the number of neutrons by subtracting the number of electrons from the rounded mass number.

• The periodic table also tells us where the electrons are

29

located in the atom. By that, I mean the orbital or – which is just about the same thing – the energy level of the electron.

• There are four types of orbitals currently known. In a few tens of billions of years, there could be more. (I’ll let you know!) These are called “s,” “p,” “d,” and “f.” The letters are abbreviations for “sharp,” “principal,” “diffuse,” and “fine,” respectively. The four words describe the lines visible in light emission spectrums made by elements. Emission spectrums of elements are made by

exciting samples of the elements either by electricity or by heating them. The energy input to them causes them to emit photons of light. If the light is looked at using a spectroscope, it appears as lines of various colors on a dark background. The location and intensity of the lines is characteristic for a given element and relates to the type of orbital occupied by the electrons of the atom. Information on electrons and their orbitals was first learned by use of spectroscopes.

We will make and use a student spectroscope. The body of the ‘scope is cardboard. You have a template in these notes, which can be

30

photocopied onto light cardboard, folded and used. You will also need 1/8 of a CD. It is cut as a wedge-shape and goes inside the ‘scope to diffract incoming light. Instructions are included with the template.

Bohr Model

Please read in your text-book about Neils Bohr. His research determined that the electrons occupy orbitals around the nucleus of the atom. Further, he found that the orbitals are of discrete energy levels. Study, especially, the section describing how electrons are boosted to a higher energy level and how they emit light as they return to what is termed the “ground state.”

Bohr’s description of the atom showed that two electrons could be placed in the orbitals which produced “sharp” spectroscopic lines. These “s” orbitals were eventually mathematically determined to be spherical in shape. The other orbitals were also found to have specific, mathematically discoverable shapes. The “p”

31

orbitals were found to be at right angles to each other. In 3-dimension, then, this is three orbitals, each of which can hold two electrons for a total of six. There are five “d” orbitals, which cumulatively can hold 10 electrons, and the “f” orbitals – populated by the Lanthanide and Actinide elements – can hold 14 electrons.

Clearly, the shapes of the “d” and “f” orbitals are pretty complex! Bohr himself could not determine the mathematics for any atom other than hydrogen. Still, when we draw representations of these other elements, with electrons around a central nucleus, we still call them “Bohr models.” This is distinguished from Ernest Rutherford’s “plum pudding” model in which the electrons were embedded in a nuclear matrix with positively charged protons.

• A fundamental question remained: What was the relationship between electrons and light? While other investigators were instrumental in researching this question, it was Albert Einstein who described the “photoelectric effect.” This gave rise to the wave/particle duality nature of electrons.

32

It is a disconcerting idea, but one that fits all the data. Specifically, photons of light exhibit both wave (energy) as well as particle (matter) characteristics. Photons hitting matter can liberate electrons

from their surface. If you don’t think this is important, recognize that this is why photosynthesis, upon which life on earth is based, works.

Further, the intensity of the light is not what liberates the electrons; the intensity is related to the amount of electrons that flow, but the frequency of light is critical to electron release. If the wavelength is too long, the photons are not energetic enough to release electrons and nothing happens.

Energy in Light • We will conduct an experiment to investigate this

effect. You are provided with a cardboard tube to block

out light. The tube has a cap at one end on the inner surface of which is glued a luminous star. This star will absorb light energy and glow for a while after the light source is

33

removed. You are also provided with a set of light emitting diodes (LED), in red, yellow, green, blue and white (which is supposed to provide ultraviolet light). 1. Look in the open end of the tube. What do you see? If the tube has been left on a table, open end up, you may see the star glowing. Turn it over so no light reaches the star. When it fades, you are ready to start. 2. Turn on the red LED. Place the tube over the LED and let the star be exposed to the light for 15 seconds. Then look through the tube. What do you see? If the star is glowing, record how long it continues to glow. Repeat the activity for 30 seconds, then for 1 minute. If the star glows, let it fade before the next step. 3. Turn on the yellow light in the same manner as you did for the red, for the same time intervals. Record your observations. 4. Repeat for the other color LEDs. 5. Answer these questions: a. What is the wavelength ( ) of each LED? b. What is the frequency of the light from each LED? ( f = C/ ), C = 299,792 meters/second. c. Is there a relation between how long the star

34

was exposed to the LEDs and how long it glowed afterward? d. What do you conclude from this experiment regarding the photoelectric effect?

• The type and characteristics of the orbitals teach us why and how atoms combine with each other (Standard 1d).

It’s all about energy. Basically, atoms are lazy. The only time electrons populate higher energy levels is when the lower ones (that is, the orbitals closer to the nucleus) are full. If these, outer electrons are then energized, they bound to higher energy levels yet, but then fall back down as far as they can go, emitting a photon of light in the process.

The characteristic of “lazy” also relates to HOW atoms combine with each other. First, it is only the outermost or “valence” electrons that enter into reactions. Second, if there are only one or two electrons in the outer orbital, then the atom is more likely to give them up to form a bond, instead of trying to accumulate six or seven to complete that

35

orbital. Classic examples of atoms that give up

electrons include the alkali metals, such as lithium and sodium. They each have one electron in their outer orbital. By losing the electron, they take on the configuration of the nearest noble gas; helium for lithium and neon for sodium.

Notice that the “magic number” seems to be eight. In most cases, stability of atoms occurs when the “p” orbitals contain eight electrons and take on the configuration of noble gases.

For atoms whose electrons occupy the “d” and “f” orbitals, they are at lower energies than the “p” orbital, so they must fill first.

Boron is one of the exceptions to the eight-electron rule. This is because boron, atomic number 5, has two electrons in its first “s” orbital, two in its second “s” orbital, and one in its “p” orbital. The electron in one of the “p” orbitals is easiest to lose because six can be held. With only one electron, elemental boron is unstable. It takes less energy for boron to kick out an electron than it is for

36

beryllium, which has two electrons in its outer orbital to begin with. (When we study ionization energy, you will see this.) But that is not all! For greatest stability, the two more electrons in the second “s” orbital have to be donated to some other atom, as well. With those electrons gone, the “outer shell” now becomes the first “s” orbital and it is filled by just two electrons.

At the other extreme, we have the halogens. They all have seven electrons in their outer orbital and can become stable if they acquire another electron. It is easier (less energy) to gain an electron than to strip off all seven already in the outer orbital. “Mr. Sodium, meet Ms. Chlorine. You give an electron; she accepts it and both of you now appear “noble.” So long as you stay together, you have a 50:50 relationship. You remain neutral in the world of ionic chemistry. But if your relationship gets all wet, then you may part; she keeping your electron, and being quite negative about it! You on the other hand, get a charge as well,

37

but remain positive until you two come together again.”

Carbon is right in the middle, with four electrons in its outer orbital. It can go either way. Caught this way, carbon shares its electrons by forming covalent bonds, not the donating and accepting method of ionic compounds. “Covalent” literally means, “sharing outer valence electrons.”

Periodic Table – Other Uses There are at least three other features of the periodic table that are not necessarily apparent from its outward form. These are electronegativity, ionization energy and atomic size (Standard 1c). • Ionization energy is, literally, the energy needed to

remove an electron from the parent atom. This always, always, always results in more protons in the nucleus than electrons remaining around it so the resulting ion is positive. Ionization energy has been determined for

almost all of the elements. There is an equivalent energy, called “electron affinity,” and it is the energy released when atoms gain

38

electrons. From what you have already learned about “lazy” atoms, you probably realize that only those atoms which will tend to gain electrons to begin with have an appreciable eletron affinity. Most often, only halogens and their neighbors starting with oxygen are really able to demonstrate electron affinity. For us, the most important thing to note about electron affinity is that the values are “negative” to indicate the release of energy, instead of positive, which shows that energy is needed to ionize the atom.

In many cases, the energy to remove a second, third or more electrons has also been determined. It takes much more energy to remove the second electron than the first, and so forth. In the case of electron affinity, only elements in column 16, starting with oxygen, even aspire to collect a second electron.

Ionization Energy, Electronegativity • Ionization energy is very useful in determining

how readily elements will combine with each other and, remembering John Dalton’s law of Multiple Proportions, teaches us why.

39

Please make a chart of the ionization energy of the elements using the ionization table included in these notes: 1. Use graph paper. With a ruler, draw

vertical and horizontal axes. Label the vertical axis “Ionization Energy, kJ/mol. ” Label the horizontal axis “Atomic Number.”

2. Mark off the vertical axis in equal units from 300 to 2,400. Mark the horizontal axis from 1 to 104: You may need to tape additional sheets of graph paper together to be able to list all of the atomic numbers.

3. Plot the 1st ionization energy for each element on the chart. You probably will not have room to list the symbol for each element, but attempt to list at least every 4th one.

4. When you have plotted the energies, draw a straight line from the first to the second point plotted, from the second to the third and so on.

5. After you have completed this chart, study it and answer the following questions: a. Which element has the highest ionization energy? What is its family?

40

Why do you think its ionization energy is so high? b. What other elements also have high ionization energies? Are they in the same family as the one you identified in ”a”? Can you connect these elements with another line? What is the trend of the line? What does this suggest to you? c. What elements have the lowest ionization energies? In what family, or families, are they? What is the characteristic of that family or families? d. Do elements between the highs and lows rise or fall smoothly? How can you account for the shape of the peaks on the chart? e. Why does the ionization energy decrease as the atomic number increases?

• At this point you might be wondering if there is a relationship between what is called “electronegativity” and ionization energy. Actually, there is more of a relationship between electronegativity and electron affinity.

Electron affinity is the energy to attract an electron to a neutral atom.

41

Electronegativity measures the tendency of an atom in a molecule to pull electrons from another atom in that molecule to itself.

Electronegativity is a dimensionless number ranging from 0.7 to 4.0. The concept was introduced by Linus Pauling in 1932 who observed that covalent bonds between various atoms was stronger than was otherwise predicted by valence bond theory.

• The more electronegative an atom, the more it pulls an electron from a partner atom. This causes the bond between the two atoms to be “polar.” A polar bond is one in which the electrons

shared by two atoms – otherwise covalent – are shared unequally. This is “polar covalent” and causes one “end” of the molecule to be negative and the other end to be positive. Water is a great example of a polar covalent bond.

• The greater the difference in electronegativity, the more ionic is the bond between the atoms; the less the difference, the more covalent is the bond. What kind of bond is between two oxygen

atoms? Between two fluorine atoms? Both are reactive, but they are the same type of atoms so the bond is totally covalent!

42

What kind of bond is between sodium (0.9) and chlorine (3.0)? The difference is 2.1. This is clearly a polar bond, but is it completely ionic? We do call salt an “ionic solid,” but in fact there remains some covalent character to the bond. I haven’t found any source that WANTS to set a definitive point at which a bond stops being covalent and IS ionic!

• You can see that electronegativity is informative, although perhaps not as much as ionization energy is. We study electronegativity mainly as a periodic table “trend.” That is, it increases left to right and bottom to top of the table. This makes sense when you compare the most electronegative element, fluorine (4) in the top, right corner, with the least electronegative francium (0.7), in the bottom left. The trend has everything to do with the way the

electrons are arrayed in atomic orbitals. Each successive layer of electrons shields the outer most electrons – the ones in the valence layer – from the positive charges in the nucleus. Francium, being an alkali metal – albeit rare and radioactive -- has only one electron in its outer orbital. It is in period 7 so there are six

43

layers of electrons between it and the 87 protons in its nucleus. This means the outer electron is a substantial distance from the nucleus. Regardless of the number of protons, this causes the electron to be only weakly held.

Contrast francium to lithium, which only has three protons, and two orbitals of electrons. It’s electronegativity is 0.98 indicating a stronger grip on its valence electron. Clearly, its size will be less than francium.

Finally, contrast lithium to fluorine. Is fluorine larger or smaller than lithium? From the electronegativity of 4, we know that the atom holds its electrons tightly. It is in the same period as lithium, so there are no additional orbitals to shield electrons, but there are additional protons. These protons exert attraction for the valence electrons. Consequently, we would predict from the electronegativity values that fluorine is smaller than lithium. Look at this table: Element Electronegativity Atomic radius fluorine 4.0 50 – 60 pm lithium 0.98 145 pm francium 0.7 270 - 280 pm

44

Our expectations are realized.

• So the concept of electronegativity teaches us about periodic table trends. You can understand WHY as well as HOW atomic size changes across, as well as up and down the table. We’ve also discussed ions. A good question is

what can we learn about ionic size? Does electronegativity help us there, too?

Think about it: When an atom gains an electron, that electron goes into the outer orbital. Remember that the number of protons has not changed, but now the outer orbital will feel less attraction because it has an additional electron. This means the outer orbital has to be greater in size. Let’s compare the ions of lithium, fluorine and francium: Ion Ionic diameter fluoride 133 pm lithium 76 pm francium 194 pm

Compared to the size of the neutral atom,

lithium and francium ions are smaller. Why? (Hint: remember what happens when atoms

45

ionize.) For the same reason, the fluoride ion is larger.

Clearly then, trends are established, and electronegativity played a major role in the understanding.

• Please use your textbooks to study the figures and diagrams showing the size of atoms and ions, and how they trend in one direction or another on the periodic table. Also, review the periodic table showing electronegativity, appended, and understand, in context of these notes, what the table shows. Spend time in analyzing the ionization energy chart that you made to understand how it relates ionization energy to stability of the various orbitals and sub orbitals.

• The periodic table is “LOGICAL.” Studying it teaches scientists and students how well it makes chemistry understandable. One feature is how each element has a specific number of electrons (based on atomic number of the neutral atom), and – more importantly – how they are arrayed in the various orbitals and suborbitals of the atom.

46

The ionization energy chart helps to visualize how the atom stabilizes itself by the location of electrons.

Another diagram, based on the Aufbau principle (“the electrons in an atom are so arranged that they occupy orbitals in the order of

their increasing energy,” shows how electrons actually fill orbitals by energy levels. The diagram was developed by Edwin Madelung and is the Madelung Diagram. Follow the arrows. It is based on the seven periods of the periodic table. Each “s” orbital can hold two electrons, each “p” orbital can hold eight, the “d” orbitals hold 10 and the “f” orbitals, occupied only by the Lanthanide and Actinide series of elements, hold 14 electrons.

Example: What is the electron configuration for selenium, atomic number 34?

47

1. Since the atomic number is the same as the number of protons, and the number of protons in a neutral atom is the same as the number of electrons, there are 34 electrons. 2. Following the arrows in the Madelung diagram, the configuration for selenium is: [1s2][2s2][2p6][3s2][3p6][4s2][3d10][4p4]

Using your knowledge of how many electrons fit in the “s,” “p,” “d” and “f” orbitals, and the Madelung Diagram, please write the full electron configuraton for the following elements: 1. Carbon 2. Helium 3. Manganese 4. Iodine

5. Sulfur 6. Thallium 7. Radon 8. Yttrium

Chemical Bonds At the start of this course I noted that EVERYTHING is based on chemistry. Now, we are going to see exactly what is meant by that statement. Standard 2 is “Chemical Bonds.” Quoting from the standard, “Biological, chemical, and physical properties

48

of matter result from the ability of atoms to form bonds from electrostatic forces between electrons and protons and between atoms and molecules.”

• Lets analyze that statement: Do you understand the words? What is a “biological property?” What is a (strictly) “chemical property?” And, how does chemical bonding affect physical properties of matter? Please write your analysis. Give it some

thought. Use your textbook for support if you need to.

I think you know by now that atoms can form bonds with each other. By extension, they can also form bonds with molecules. If you do not understand that, please go back and review notes for Standard 1d.

And everything, then, seems to depend on what is titled “electrostatic forces.” We have a clue to understanding this term in that some force interacts between protons, with positive charge, and electrons, with negative charge.

Just what IS an electrostatic force? The first part of the word, “electro-“ clearly relates to