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PRINCIPLES OF ALCHEMY
EARTH

Introduction to the Periodic Table

I promised you a way to keep track of all these elements and their properties. Most of this lesson will explain how we Alchemists do just that.
You may have noticed some common properties for some elements by now.

Yeah. Some elements have the same kind of outer shell so they make bonds in similar ways. Right?

Aye, you're right. That's a very important observation.
The number of protons (atomic number) defines the element and give it its properties. But the electrons are the real "workers" in Alchemy because the electrons are responsible for all the chemical bonding. We Alchemists have learned to keep track of the electrons, especially the outer shell electrons, because they are the most important for making molecules.

It's kind of an "inside out" way of looking at atoms, isn't it?

How do you mean?

Well, we think first of the atom's nucleus. First we count the protons to determine the atomic number and that defines the atom.

Yes, from the atomic number we figure out the kind of atom (actually element) we are dealing with. What do you mean by an "inside out" look?

Once I know the proton count I figure out how many electrons it should have. Then I work from the inside outwards filling shells until all the electrons are in shells. The outer shell is where all the action is. Right?

Aye.

When I imagine the atom, I start with the nucleus but end up spending most of my time thinking about the outer shell. Inside out!

Oh, I see what you mean! Yes, that is true. Very good. Inside out. You always start with the nucleus, the protons, in order to get to the outer shell.

So, when you we say that the protons define the element we mean it determines what kind of outer shell it has. Right?

Right!
It wasn't until the 20th century that Alchemists knew about shells and other important features of atoms. Before then, all that Alchemists had to work with was a long list of different chemical properties and the ratios of elements in molecules.

What do you mean by "ratios of elements in molecules"? Like what?

Like the fact that two chlorines can bind with magnesium (to make MgCl2), but only one chlorine binds with sodium (to make NaCl). Stuff like that use to drive Alchemists crazy! Why do you get twice as much chlorine from magnesium chloride as you do from sodium chloride?

Because magnesium has a +2 charge so it can hold two chlorines. Sodium has a +1 charge so it can hold only one. It's just the atoms trying to arrange their shells to look noble. In the case of these two elements, they do it by loosing one or two electrons and forming ionic bonds.

Yes. Exactly. During the 1800's many Alchemists started to look for patterns in these ratios. It was not easy. They did not have a complete understanding of atoms and they thought that atomic MASS was the key, not atomic NUMBER! However, even that mistaken system got them somewhere.
In 1866 a fellow named Newlands noticed that when (some) of the elements are listed in ever increasing mass, they tended to repeat their properties every eighth time. Newlands called this his "Law of Octaves" because it reminded him of the octaves used in music. He said "the eight element, starting from a given one, is a repetition of the first."

Newlands' Law of Octaves is just what I was saying before about elements with the same kind of outer shell making the same kind of bonds! Because the L, M and N-shells can all hold eight electrons they will have similar bonding depending on how full they are. An element with one electron in its L-shell will be like an element with one electron in its M or N-shell.

Yes, that's absolutely right. Newlands' Law of Octaves is based upon that idea, although Newlands didn't know it. That's also the basis for Lewis structures.

So this Law of Octaves is just another way of looking at Lewis structures and his "octet rule"?

Yes, it could be used that way. Unfortunately as you go to other, larger, elements the Law of Octaves (and Lewis structures) starts to come apart. It just doesn't work for gold or iron and many other metals. Three years after Newlands' Law of Octaves a Russian Alchemist named Dimitri Mendeleev developed a slightly better theory about the chemical periodicity. He drew a table of elements. We call it (Mendeleev's) Periodic Table of the Elements.

So, Mendeleev fixed Newlands' Law of Octaves, and then all of Alchemy suddenly made sense?

Well, sort of.
Mendeleev set up the foundation of the Periodic Table, but he still had some errors because he was still arranging the elements according to their atomic mass, instead of their atomic number.

That's so stupid! The behavior of the elements depends on the number of protons not their total nucleon count (mass). Protons define the elements. Neutrons just add weight. I mean MASS. You taught me that.

Yes, but I didn't teach that to Mendeleev. He and I never met! You, Arthur, are fortunate to have my 21st century knowledge to guide you. Mendeleev didn't. Mendeleev wasn't stupid, he just misunderstood. As it so happens, the number of neutrons in an atom usually increase in the same manner as the number of protons. At least for most of the elements that Mendeleev was working with.

How do you mean?

Well, for example, helium's most common isotope (4He) has two protons (that's why it's helium) and two neutrons. Carbon's most common isotope (12C) has six protons (that's why it's carbon) and six neutrons.

Does the number of protons always equal the number of neutrons? Is that some kind of rule?

NO!!! It is a "trend" but it isn't a rule. For the smaller elements it is a good guess that the most common isotope has the same number of protons as neutrons. Unfortunately, that trend doesn't fit all the elements and it lead Mendeleev down the wrong path.

So Mendeleev's Periodic Table is all wrong? Why am I learning this?

His Table was not all wrong. It was just a wee bit off. Yet he was able to use it to predict the existence and properties of elements that were unknown in his time! More importantly, he showed (as did Newlands) that there was some order to all this madness - all these elements - even though he was a little off on some.

Who set Mendeleev straight?

A fellow named Moseley. He worked in an area of Alchemy that we haven't really touched on. Moseley found a way to cause an element to produce X-rays - powerful light (but not quit as powerful as gamma rays). I won't go into the details of how or why he made X-rays from elements because that's a subject for Advanced Alchemists. The important thing is that Moseley discovered that the larger the elements the more X-rays were produced. But he could not get an exact relationship between the atomic masses and the energy, so he simply assigned a number to each element. He called it an "atomic number", so as not to get people confused with atomic mass. Hydrogen produced the least amount of X-rays so he numbered it 1. Helium produced more so it was number 2. Lithium produced more than helium, so it was assigned the atomic number 3. It went on and on. Eventually he had a list of all the elements arranged according to how much X-rays they made.

It sounds like Moseley's X-rays let him count the protons and arrange the elements in increasing proton number. Atomic number!

That's right, but at the time he didn't know it. When Alchemists arranged the elements according to Moseley's atomic NUMBER instead of atomic mass, suddenly Mendeleev's Periodic Table made absolute sense! The elements which had been causing trouble before, because they didn't fit the "patterns", now fit the Table. Moseley had discovered a better way to arrange the elements. The correct way. The world of Alchemy turned away from its fascination with atomic mass and started to pay attention to Moseley's atomic numbers.

So it should really be called "Moseley's Periodic Table of the Elements".

Well, we still think of it as Mendeleev's Periodic Table. Actually, it was a bunch of people who helped put the Table together.

OK. Let's see this magic Table!

All in due time, my lad. If I were to show you the whole table all at once you would get carried away with all its detail. You might think that it caused you more difficulty than help. So, I will teach you all about the Periodic Table of the Elements one step at a time.

Let's start with the really important elements, the "typical elements".

What's typical about them?

They are called typical elements because they behave "typically". What I mean is the typical elements all follow the rules you have learned. They fill their electron shells from the inside out. They take part in forming covalent or electrovalent bonds, just like we've discussed. You can use Lewis structures to study them. They are also the most common elements in the universe!

I see. I guess these typical elements account for most of Alchemy.

Yes, that's right. Therefore, we will spend a great deal of time learning about them and how we use the Periodic Table to predict and understand their behaviour.

I guess the "untypical elements" will have to wait.

Yes. (By the way, we have better names for them than "untypical", but that's later.)
We will start with a Periodic Table of (just) the Typical Elements. Typical elements include the first 20 elements (those with atomic numbers between 1 and 20) and another two dozen heavier elements whose atomic numbers do not follow a simple series of numbers.

Why don't they?

Why don't the two dozen heavier elements continue the series?
Well, it has to do with the fact that many of the larger elements have unusual properties. With bigger shells things can get more complicated. We'll discuss all that later. The important point is that the typical elements are the first 20 plus another 24 heavier ones. OK?

So Let's see this Periodic Table of (just) the Typical Elements!

OK. Here it is for the first 20 elements. These are the most important elements in the universe, and you're already familiar with most of them.
Let me point out the important features. The elements are arranged so that each has an atomic number one unit (one number) greater that the preceding element. Notice that the "periodicity" runs from left to right and each row repeats when the outer shell is full.

But the two elements in the first row (hydrogen and helium) kind of stand out because they can only fit 2 electrons into their outer shell. All the others can fit up to 8 electrons into theirs.

That's right. The first row, or "Period" as we call them, represents the K-shell. You know that the K-shell can only hold two electrons. Once it is full, electrons are added to the L-shell. The second Period (the second row) represents the L-shell.
With each newly added proton, you move to the right, to the next element. As you do that you continue to add an electron for each proton. Each new electron is added to the outer shell until that shell is full. When the outer shell is full you move down to the next Period (row) and start over.

Each Period (row) is a shell. Easy!

Yes, and useful. Because when lined up this way, each column (running from top to bottom) contains elements with the same type of outer shell. See?

Yeah, I see what you mean. All the elements in the first column (hydrogen, lithium, sodium and potassium) have one electron in their outer shell. The elements in the second column (beryllium, magnesium and calcium) have two electrons in their outer shell.
Hey, wait a minute. Helium also has two electrons in its outer shell. Why isn't it placed next to hydrogen, right above beryllium? Why is it over with neon? Neon's got 8 electrons in its outer shell. This table is wrong!

No, the Table's right, but you're looking at it wrong. You've made a common mistake. (Every student does it.)
The columns are not meant to line up the number of electrons in the outer shell (although they usually end up that way). The columns are arranged to show HOW FULL IS THE OUTER SHELL. Elements in the same column have the same TYPE of outer shell, not necessarily the same number of electrons.

Oh, I see what you mean. Because the first shell is a K-shell, holding only two electrons at the most, helium is put into the same column as the other noble elements. All those elements in the 8th column have a complete outer shell. Right?

Right!

I see. (I think.)
Helium, neon and argon all have complete shells so they are all in the last column. It's their complete shell which makes them noble elements.

Yes indeed. That brings us to the most important property of the Periodic Table. The columns, or "Groups" as we Alchemists call them, each contain elements with the same type of outer shell. Therefore, elements in the same Group will have similar chemical properties.

So the elements in the second column, I mean second Group, need to gain six electrons, or lose two electrons, to complete their outer shell.

Yes.
You know that it's the outer shell that gives each element its ability to bond in certain ways. Beryllium, magnesium and calcium will behave the same way to try to complete their shell. All three of them will donate their 2 outer electrons to form cations with charges of +2. So they will take part in electrovalent bonds to anions.

I think I understand. All the elements in the 8th Group have complete outer shells so they are inert. All the elements in the 7th Group need only one electron to complete their shell, so they are likely to steal it away and become an anion, with a charge of -1. So all the elements in the 7th group are likely to form electrovalent bonds with any cations they can find.

Very good. I think you are getting the hang of it. One thing to know, however, is we Alchemists like to speak about the 7th group as Group seven or simply Group VII (using Roman numerals). We name all the Groups that way.

Do I have to memorize this stuff?

Well, you should know how to USE the Periodic Table, but there is no need to memorize the Table itself. You can always look it up in a book. However, if you understand how the Table is made, you should be able to build the Table from "first principles". All you need to remember is how to line them up to produce the periodicity.

Yeah, and memorize the order of all the elements! You said there are about a hundred of them. I can't memorize a hundred elements and their atomic numbers.

You don't have too. That would be quit a challenge. Even I haven't bothered to learn them all. That's what the Table is for. It's a reference and guide to the behavior of the elements.
You know, I have memorized the order of the first 20 elements. It helps me at times when I just need a quick answer about them. After all, the most important elements are the first 20!

Ugh! How dull to sit around memorizing 20 elements! You need to get out more.

I'll ignore that. I'll even let you in on my little secret about how I keep track of the first 20 elements. Instead of memorizing each element, I memorized a silly poem which reminds me about them and their order. It goes like this:
"Hi. Here Little Beggar Boys Catch Newts Or Fishes. Never Nab Maggots Alone Since Poisonous Substances Clog Arteries. Keeping Calm."

That is stupid. It doesn't even rhyme! And they don't line up!

It's not meant to rhyme. Each word reminds me of an element, and the words are in the same order as the elements. True, they aren't lined up into Groups or Periods, but I can easily do that because I've memorized their order. From that I can draw them in their proper position in the Table. It's the order that is important.

It's a stupid poem.

Then make up one of your own. I find it useful.
For example, if I can remember that carbon has 4 electrons in its outer shell, then I know oxygen has 6 because it is just two words away - from "Catch" to "Or". If I want to know the atomic number of any of the 20 elements I just count on my fingers as I recite my poem until I get to the word that relates to the element I'm after.

You'll need to use your toes if you're counting up to calcium. Or, should I say "Calm"?

Go ahead and make fun, but my "stupid poem" is a useful device.

OK. OK. I'll try to come up with a better poem.

Good. You'll learn a lot by trying.

Now, I mentioned that there were another 24 typical elements. Let's add them to the Table. I'll place them into their proper Groups and write in the atomic numbers too. What do you notice about this "complete" Periodic Table of the Typical Elements?

Ah, that there are a lot of new abbreviations to learn?

You'll pick up the abbreviations as we go along. The important ones are the ones that will stick in your head without really thinking about it. The rest you can look up later in a list.

OK. You don't have to convince me not to study.

I didn't say that! I said don't memorize a long list of abbreviations. Now, what else do you notice about the Table.

Lots of atomic numbers are missing. Ten are missing between Ca (calcium) and Ga (gallium) and another 10 are missing between Sr (strontium) and In (indium). Look! There's even more missing between Ba (barium) and Tl (thallium). I bet the ones that are missing are the "untypical elements".

A good bet! The "untypical elements" (as you call them) do not have the same electronic configuration as the typical elements, so we leave them out. For now.

I see you've labeled the Periods and the Groups.

Yes. This table is more complete and useful. Notice the Groups now have many new elements. Because they are within Groups whose elements we studied earlier, we can predict the behavior of these new elements by referring to the fact that they have outer shells like the smaller elements we've been studying. Do you understand?

I think so. I suspect that Rb (rubidium), Cs (cesium) and Fr (francium), will all form electrovalent bonds just like Li (lithium), Na (sodium), or K (potassium), because they are in the same Group (Group I).

That's right! All the Group I elements have a single electron in their outer shell and would be happy to lose it in order to achieve a complete outer shell. So all the Group I elements will form cations (with a charge of +1).

Does that mean atoms of rubidium, cesium and francium will make electrovalent bonds to chlorine in the same way as sodium chloride (NaCl)? Making molecules of RbCl, CsCl and FrCl?

Yes, that's absolutely right!
You see. You've predicted the bonding of elements whose names you never heard before, because you know how to use the Periodic Table to relate them to elements you know very well. That's the power of the Periodic Table!

I see what you mean by the Table being a useful reference to follow. It is a map of element behaviors.

Aye, that it is. You seem to be grasping it well.
Also, you may have noticed that you are very naturally putting the cation first and the anion second, in your description of these kinds of molecules. That is the correct way to write ALL ionic compounds - cation then anion. It is also a very natural way to read the Table - from left to right. The correct way to speak an ionic compound is to say the cation first and the anion second. Any questions about naming ionic compounds?

Why do I call it sodium "chloride"? Shouldn't it be sodium "chlorine"?

Good question. We Alchemists like to change the ending on the anion by adding the suffix "-ide" to the root of the word. So chop off a few of the letters at the end of the anion and add "-ide" in their place.

Sounds simple enough.

It is. Now. Tell me about the elements in Group VII.

They are like chlorine (or fluorine). They will all have seven electrons in their outer shell. So they will be happy to take any extra electron they can find in order to complete their outer shell (to have the electronic configuration of a noble element). So they will form anions (charge -1) and bond to cations by electrovalent bonds.

Yes. Can you "guess" a few compound molecules you could make using group I and group VII elements?

Sure. Sodium chloride (NaCl) is the one we've talked about a lot already. I suppose you could also make sodium bromide (NaBr), sodium iodide (NaI) and sodium astatinide (NaAt).
I suppose you can use lithium as if it were sodium to make lithium fluoride (LiF), lithium chloride (LiCl), lithium bromide (LiBr), lithium iodide (LiI) and lithium astatinide (LiAt).
Hey, I can see even more, like potassium fluoride (KF) and....

Yes. Yes! I think you've got it. There are many different compounds one can make from combinations of the Group I and Group VII elements. That, again, is the power of the Periodic Table. You can match any Group I element with any Group VII element to make an ionic compound containing one atom of each.

What about compounds made from Group II elements?

Like magnesium (Mg) and calcium (Ca)?

Aye. Predict the compounds you can make from combinations of Group II and VII elements.

Simple. All the Group VII elements will behave as before, forming anions with charges of -1. But the Group II elements will form cations with a +2 charge.
I bet you thought you could trick me into forgetting that the Group II elements must loose two electrons to have a complete shell.

There is just no tricking you, is there Arthur?

That's right. Anyway, those cations (of the Group II elements) will attract the anions (of the Group VII elements) by electrostatic attraction. So they will form electrovalent (ionic) bonds between them. Each Group II element will bond to TWO elements from Group VII, because the Group II ions have a charge of +2 but the Group VII ions have a charge of -1. So you need two Group VII ions to neutralize each Group II ion.

Very good and clear thinking Arthur.
Name a few compounds made from Group II and VII elements.

OK. We talked about magnesium chloride (MgCl2) a lot last time. All the Group II /VII compounds will be made the same way. Each Group II atom will form electrovalent bonds to two Group VII atoms. So you can have magnesium chloride (MgCl2), magnesium bromide (MgBr2), magnesium iodide (MgI2), mag.......

Very good Arthur. I think you've got the hang of it.

But I've only done the magnesiums. I can substitute any of the group II elements to get other compounds like calcium chloride (CaCl2), strontium chloride (SrCl2)...

Yes very good. Now...

and barium iodide (BaI2), and radium fluoride (RaF2), and ....

That's enough Arthur! You've got it! (Phew!)
Notice. All the Group I atoms make diatomic, electrovalent compounds with one Group VII atom. All the Group II elements form triatomic, electrovalent compounds with two Group VII atoms.
Tell me, will a Group III or V atom form electrovalent bonds with three or five Group VII atoms?

No, that's not the way it works! Those Groups don't form electrovalent compounds. If they did they wouldn't simply bond to as many atoms as their Group number would suggest. The bonding has to do with the electrons in the outer shell.

That's right. The Group I and II elements are quick to donate electrons (from their outer shell) in order to obtain the electronic configuration of a noble element. The Group VII elements are quick to pick up an electron in order to do the same. The other Groups are more likely to share their electrons in COVALENT bonds in order to fill their outer shells.

So the Periodic Table lets me keep track of the outer shells and that tells me how it will bond. Anything else to be learned from the Table?

Oh, yes indeed. The Periodic Table allows us to compare more than just the elements' bonding properties. The physical properties of the elements follow trends that are best understood by referring to the Periodic Table.

What do you mean by "physical properties of the elements"? Name one.

OK. Atomic size.

But atoms are all very, very small. Do you mean to say that some atoms are bigger than others?

Yes. That's right.

These distances must be very, very, small. Are they measured in tiny fractions of an inch?

No, no. Nobody uses inches in the future. In the 21st century the last nation on Earth to use inches FINALLY converted to metric. After nearly a century of trying!

Why did it take them so long? A century to change to the system the rest of the world uses!

There's an old saying, "Some people would rather die than change. And eventually they do."

So the old ways of measuring died out with the old folks who refused to change.

Aye. All sciences (even in the old-fashion nation) use metric. It's all metric.
Time for a quick metric measure - the meter. A meter is just a little over a yard, about 39.4 inches, when compared to the old fashion way of measuring things.

But atoms are not nearly so big as a meter! They must be measured in some tiny fraction of a meter. Right?

Yes, you're right. Sort of.
The meter can be divided into smaller and smaller units. Instead of doing it by fractions (1/2 then, 1/4, then 1/8 and so on, the old fashion way), we Alchemists use decimals. So a meter is first divided into ten smaller "decimal fractions" of 1/10th of a meter. Each of these can then be divided further into ten more decimal fractions, each 1/100th of a meter. These are called centimeters from the Greek "one hundredth of a meter". Each centimeter (or "cm" to use the abbreviation) is about half an inch.

That's still not small enough for atoms!

You're right. We have to keep on dividing by ten. Next comes millimeters ("mm" is the abbreviation). Millimeters are 1/10th of a centimeter, so they are 0.001 of a meter. Again, we use a Greek word "milli" for a thousand. A penny (coin from the 20th century USA) is about two millimeters thick. So that coin is (about) 0.002 of a meter.

What if I don't have a penny?

Well, fingernails are about half a millimeter thick, or 0.0005 of a meter (1/2 of 0.0010 is 0.0005).

I thought "milli" would mean "millionth"!

When I first learned Greek, I thought so too. A millionth of a meter is called a micrometer, or sometimes simply "micron". That would be 0.000001 of a meter. It's far too small to see.

Yeah and difficult to write all those zeros. Are atoms measured in microns?

No. Even microns are too big to describe atoms (without using a lot of zeros). Alchemists measure atoms in angstroms. An angstrom (pronounced "ang-strum") is 10 billionth of a meter, or 0.0000000001 of a meter.

Wow, that is small. How do Alchemists measure atoms? You can't see them!

Alchemists use X-rays to figure out the distance between two atoms, using a technique I don't want to get into right now. For example, using X-rays an Alchemist measures the distance between the two hydrogens in H2. From that "interatomic distance", we can calculate the radius of each hydrogen atom.
The atomic radius is half the distance between the centers of adjacent atoms. See?

Yeah. Isn't that another way of saying the atomic DIAMETER is equal to the distance between the centers of two atoms?

Yes. That's true. Some Alchemists express their numbers that way. Let's use ATOMIC RADIUS for this discussion, shall we?

OK. So exactly how big is an atom?!

That depends on the atom and it also depends upon how they are bonded (if they are bonded at all).
Imagine two atoms of the same element pushed together until their electron shells "touch". Then they start to repel each other.

Because of the electrostatic repulsion of all those negative charges from the electrons. Right?

Right.
We are not thinking about bonding them together. No electrons are shared, so they are not covalently linked. Let's say the centers of these atoms are separated by 1.82 angstroms (measured by X-rays). What is the atomic radius of these atoms?

Well, if I understand your diagram correctly, that would be half the inter-atomic distance. ("Inter-atomic" just means "between atoms".) So each atom has a radius of 0.91 of an angstrom.

That's right. The atoms in that example have an atomic radius of 0.91 angstroms. However, atoms of the same element are not always the same size.

What!? Why not?

Carbons involved in covalent bonds will have smaller radii (plural of radius) than a pile of unbonded carbons. Can you tell me why?

Could it be that when atoms share their outer electrons, to make covalent bonds, they are drawn closer together?

Yes. Covalently bonded atoms not only share their outer electrons, they also share their entire outer shell.

You know, when I draw Lewis structures I think that only a specific pair of electrons is shared, forming the bond.

Yes, that is a common way to draw the structures but in reality any electron in the outer shell might be involved in the covalent bond, at any one time. As a matter of fact, all the electrons in the outer shell take part in a covalent bond. They take turns. For each covalent bond, only one electron from each atom (the "covalent pair") is involved at any one time.

Does that mean the molecular orbitals I learned, don't work?

Oh, they "work"! At any one time the electrons are distributed among the orbitals exactly as we described them. However, they might switch places (orbitals) with each other (within a shell). It's another quantum mechanical property. It's hard to follow a single electron in its orbitals, so we think of it as an averaged effect.

Don't worry about it. Right?

Right. Let's get back to the idea about covalent radius.

OK. Two atoms held by a covalent bond will be drawn closer together because a bit of their outer shells can overlap.

That's right.

So how much smaller are covalently bonded carbons compared with unbonded ones?

That depends on exactly what material the carbons are in. For example, in a diamond all the carbon atoms are spaced 1.54 angstroms apart. In graphite (the correct name for "pencil lead") the distance between the two covalently bonded carbons is 1.42 angstroms.

So the carbons in diamonds have a radius of 0.77 angstroms but in graphite they are 0.72 angstroms in radius. (I just divided by 2 to get the radius.)

Right. Diamonds and graphite are allotropes of carbon. Forms of an element having different physical properties (like size) are called allotropes (pronounced "al-oh-tropes") or allotropic forms of that element. All the carbons in graphite are arranged in flat sheets. In diamonds those same atoms are arranged in three dimensional patterns which allow the covalent bonds to stretch a wee bit.

Sounds very confusing. Am I suppose to memorize all these different sizes for carbon.

No. No. That would be a waste of time. What I want you to realize is that the way you measure an atom's size depends on if, and how, it is bonded.
Carbon as an atom, not in a molecule, has a radius of 0.79 angstroms. We call that the atomic radius, because it is the radius of atoms not involved in molecules. The atomic radius is the distance from the atom's nucleus to its outermost shell. The outermost shell is where that atom will run up against another atom's shell.

When carbons are involved in covalent bonds they are able to get much closer together by sharing electrons in their outer shell.

Yes. The atom's covalent radius will be smaller than its atomic radius because the covalent radius is measured using covalently linked molecules of the atoms.

If atoms of the same element can have different distances from each other because of the ways they are "formed", what good is the idea of atomic size?! It isn't constant.

You're right it isn't a simple constant number, but the important point is that atoms CAN be measured and compared with others. We can express the size of carbon as an average for diamonds or graphite or some other carbon to carbon molecule. (There's lots of them.) That would be carbon's average covalent radius.
Just keep in mind that the size varies because of the way the atoms are packed or the way they interact. There are different ways to pack the atoms and make the measurements.

What's the best way?

There is no best way. Just different ways. I like to measure atoms by their atomic radius. If the atoms are not involved in bonds, their outer shell will (usually) be constant.

OK, let's stick with atomic radius for now. You were going to tell me how the Periodic Table helps you to understand the elements and the size of their atoms.

Precisely. So, let's get back on track by accepting that Alchemists can work out an "average" covalent radius by measuring the distance between atoms, as an average, in a molecule. But atomic radius is much better to work with because they are constant for each atom (or element), for the most part.

OK. I can live with that.

Good. Now tell me, Arthur. As you go down a Group, say from hydrogen (H) to lithium (Li), would you expect the atom's radius to get bigger or smaller.

Bigger.

Why?

Because lithium (Li) has an extra shell. Hydrogen has only a K-shell, but lithium has a full K-shell and a start on an L-shell. Lithium has an extra shell so it will be a little bit bigger than hydrogen.

That's right and it is right for all atoms within a Group. As you go down a Group the atoms get bigger and bigger because they have more and more shells. Understand?

I think so. Hydrogen is the smallest atom. The next largest in the Group is lithium, then sodium, then potassium, than rubidium, then cesium and finally francium is the largest! Right?

Yes.
As you go DOWN the GROUP the atoms get BIGGER. Each new Period adds a new shell.
But as you go ACROSS a PERIOD, from left to right, the atoms get SMALLER!

What?! You mean beryllium (Be) is smaller than lithium (Li)?

Yes, that's right.

Why? I would think that beryllium would be a little bit bigger than lithium because it has an extra electron.

I can understand how you would think that. Most students of Alchemy think that way, but it isn't true. You see, lithium and beryllium both have an L-shell, the number of electrons in the shell doesn't make the shell any bigger because it is the same (L) shell.

So lithium and beryllium should be the same size, because they have the same L-shell! Why did you say that beryllium is smaller than lithium? Why aren't they the same size?

Good questions. Your thinking is good and it would be easy to make that mistake. As it turns out, the atomic radius decreases as you move across a Period (from left to right) even though they all have the same shell.

Why? What causes the shells to shrink as you move across a period?

The protons! As you move across the Period (from left to right), you add a proton to the nucleus. Right?

Right. So what? You also add an electron to balance the charge.

Yes. However, each extra proton draws the shell a tiny bit closer. It's that electrostatic attraction again.

You mean, because beryllium has one more proton than lithium, it attracts the electrons in the L-shell closer to it?

Yes, exactly! As you move across the Group, you add protons but the shell is still the same (even though you are adding electrons to it, to balance the charge from each proton). Each extra proton pulls the shell a wee bit closer, causing the atom to shrink.

Is that true of only the second Period? Or all Periods?

It applies to ALL Periods in the table, not just the second Period, and not just the typical elements. All of them!
That's a very important thing to understand when you use the Periodic Table, so let me repeat how the size of the atoms changes as you move around the Periodic Table.
In ANY PERIOD the increase in nuclear charge of successive elements (due to the addition of each proton) causes an increase in the electrostatic attraction for the electrons in that shell. So the shell shrinks as you move from left to right across a period.

But in ANY GROUP the atomic radius will increase as you move down the group because another level of electron shell is added each time you move down.

Very good. It is this double effect of adding shells (going down a group) and adding protons (going across a Period) which causes the opposing growing and shrinking as you move around the Periodic Table. If you keep that in mind you will be able to predict the relative size of each atom by its position in the Table.

What do you mean "relative size"?

I just mean, how they compare to each of their neighbor elements. Let's look at the size of the typical elements and you'll see what I mean.
Notice how the radius changes as you go from one element to another. Take magnesium (Mg) for example. It has an atomic radius of 1.60 angstroms. Its neighbor to the left if it, sodium (Na) is larger (with an atomic radius of 1.91 angstroms) and its neighbor to the right is smaller. (Aluminum has an atomic radius of 1.43 angstroms.)

That's because of the effect of different numbers of protons attracting the same shell differently.

That's right. Also notice that the element directly above magnesium (beryllium) is smaller (1.12 angstroms) and the one below it (calcium) is larger (1.97 angstroms).

That's because they have different numbers of shells.

Yes. Exactly. This explains the "relative" atomic size of all the elements, not just the first 20 and not just the typical elements. All elements follow this rule:
In any Period the atomic size decreases from left to right (across a Period).
In any Group the atoms increase in size from top to bottom (down a Group).
Any questions? Or comments?

Can you use the Table to predict other things about the elements?

You sure can.

The Periodic Table can be used to explain and predict such things as ionization energy. That's the energy needed to remove electrons.

So ionization energy is the energy needed to make a cation. Right?

Yes! Very good deduction. The first ionization energy is the energy needed to remove the least tightly bound electron from a neutral atom, creating a cation. There are also second, third, fourth and more ionization energies for removing further electrons, but that is getting too complex for your introduction to the Periodic Table.

Let's stick to the first ionization energy. Do you have a Periodic Table showing the first ionization energies?

Yes, I do and here it is. Notice that the energies are lowest at the lower left of the Table, near Cs, and largest at the upper right, near He. The numbers change in a way similar, but OPPOSITE to, the way they change for atomic radius.
Moving down a Group the energies get smaller. That means it is easier to ionize an element below an element than to ionize that element itself. Do you know what I mean? (That came out a bit twisted.)

Yeah, I think so. As you go down a Group it is easier to ionize the next element. Is that because the atoms get bigger?

You tell me. How would you imagine it?

I would imagine that a big atom, with more electrons and more shells, might have more difficulty keeping all its electrons. A small atom would have fewer shells and fewer electrons to lose, so it holds them tightly. Does that make sense?

Yes, it does! You are absolutely right. The extra shells added as you move down a Group mean the extra electrons are at greater distances from the nucleus. That's why they are bigger! It also means those bigger atoms have more difficulty holding those electrons. The electrostatic forces which hold the electrons in their shells are weaker the further the electrons are from the nucleus. Each additional shell adds electrons further from the nucleus. So they are held more weakly, making them easier to remove.

I see. As I move across a Period does the same idea apply?

Sort of. The atoms get smaller as you move from left to right across a Period, because the electrostatic attraction draws them closer. That's what we said earlier was the reason atoms get smaller as you move across a Period. Remember? So, as you move across a Period (from left to right) the atoms draw in their electrons, making the atoms smaller and making it more difficult to remove the electrons.

I see. As you move across the Table the electrons are drawn closer to the nucleus and that "nearness" causes them to have greater electrostatic attraction. So they do not give up their electrons easily.

Right. Therefore it takes more energy to remove an electron from a smaller atom within the same Period. The atoms grow SMALLER as you move left to right across a Period, so the ionization energies INCREASE from left to right.

Hey! Not so fast wizard.
Look, beryllium (Be) has a HIGHER ionization energy than boron (B)! And look, look! Magnesium (Mg) and calcium (Ca) do that too! Their neighbors to the right (Al and Ga) have lower ionization energies! Some of the Group II elements stick out. They are wrong.

Well, they aren't "wrong", they just don't follow the "rule". (So it isn't a rule at all. It's a "trend".) There are some subtle factors involved which I don't want to go into because they require more advanced knowledge about quantum mechanics and electron behavior. The exceptions you've found are the only exceptions to this trend in first ionization energies. OK?

Yeah, I can live with that, but I can't live with these numbers without knowing what they really mean. How is this ionization energy measured? What's this "electron volts" written on the Table top? Are they the units of ionization energy? What do they mean? How are they determined?

You ask a lot of questions. I like that in a student. It means you want to learn and not merely memorize.
The "how are they determined?" is really an advanced topic, because there is no one method that works well for all elements. This Table is a summary of the work of many folks using many methods to determine the first ionization energy.

OK. Then what is an electron volt?

Ah, an electron volt. One electron volt (eV) is the energy gained by an electron when it "falls" through a potential (difference) of 1 volt (V).

And that means....

It means you need to know a bit of physics to really appreciate it. Here goes.
Everything that has a charge, including electrons, is influenced by electrical fields. These electrical fields are the electrical equivalent to gravity. If you drop a rock from a distance of (say) the top of your head to the ground, it picks up speed and slams into the ground.

Or onto your foot!

Yes. That is energy created by the falling rock. We can measure it and calculate its impact. Electrical fields are similar to gravitational fields in some ways, but not all ways. For example, there is no "up" or "down" in electrical fields. However, there is a force produced by these fields on anything with a charge. If you move an electron through an electric field it will gain or lose energy (depending upon which way you move it in the field). The physics folks do that all the time and use electron volts (eV) as their "yard stick".

Don't you mean "meter stick"?

Oh, yes! "Meter stick". Of course it isn't really measured in meters or any other units having to do with length. It's measured in "volts".
If you move an electron through an electric field it picks up "volts". Imagine volts as the "height". A rock at the top of a building will have more energy when it hits the ground than a rock dropped from a table top. Right?

Yeah. I wouldn't want to be under a rock dropped from the top of the house!

That's right. Well, volts are like height, but don't get the two confused! Volts don't have anything to do with distance. They are just a way to measure the energy using electrical methods. OK?

OK, I guess. Is it important that I clearly understand volts?

Not for our purposes. Just understand that an electron volt (eV) is the energy an electron gains as it "falls" through a field over an ELECTRICAL "DISTANCE" of a volt.

OK. So electron volts are just a way to measure electrical energy.

Exactly. For example, it takes 13.6 eV to remove the electron from a hydrogen atom. That means, you must put 13.6 eV into each hydrogen atom to ionize it. So hydrogen's first (and only!) ionization energy is 13.6 volts. Get it?

I think so, but I might be confused. Would this 13.6 eV be the energy one electron would get if it passed thorough an electrical field "height" of 13.6 volts? Or is it the energy that 13.6 electrons would pick up if they passed through a field of just one volt?

Both! Either way you get 13.6 eV, but for this discussion it helps to think of just one electron at a time. Besides, 0.6 of an electron is hard to come by. Actually it's impossible!

I see. If I understand this right, the Periodic Table shows that I could use 13.6 eV to ionize a hydrogen atom. Could I use it to ionize a lithium (Li) atom?

What do you think?

I think I could. According to the Table, lithium needs only 5.39 eV to ionize it. That is much less than 13.6 eV, so I would even have energy left over.

That's right and that is because lithium lies below hydrogen in the first Group. It is always easier to ionize an atom lower on the periodic table.

Hey, you know, if I had 13.6 eV and it takes only 5.39 eV to ionize lithium, I would have enough energy to ionize two electrons from the lithium. It would take twice the ionization energy.
Look it takes only 10.78 eV to ionize two electrons from the lithium. (2 X 5.39 = 10.78 eV). I would even have 2.82 eV left over. Right?

Wrong, but don't feel bad. You've made a common mistake.
Remember this is a Table of the FIRST ionization energies. It takes more and more energy to remove the other electrons.
Think about it. With one electron removed, due to the first ionization (using only 5.39 eV), the atom is now a cation (Li +). That cation will hold on to the remaining electrons very tight. Right?

Yeah, the cation will even have an extra positive charge to hold them.

Exactly! To remove the second electron from lithium you need to consult a different Table which gives the SECOND ionization energies. I won't show you that Table now. (It can be found in a large handbook of chemistry tables.) I looked it up and found it takes 75.62 eV to remove lithium's SECOND electron (from the lithium cation, Li+).

Wow, that's a lot of electron volts! That doesn't even come close to the highest value on the Table of first ionization energies. That's about three times as high as the highest eV on the Table. (Look at helium.)

Yes. You should not confuse the first ionization energy with the ionization of the other electrons. It gets harder and harder to remove more electrons.

So what happens to any extra energy? The stuff left over? Does it just go away?

Yes, sort of. If all you have is a single lithium atom it will lose its first electron - the electron that is easiest to lose. The remaining energy might be picked up by the atom as "heat", causing it to wiggle a bit more. But I don't want us getting into that right now. I have a different thought.
What would happen if you had TWO lithium atoms to begin with and you added 13.6 eV to them?

Ah, I see. They would both ionize! They would both lose their first electron and I would have two cations of lithium (2Li+).

That's right. Two lithium atoms would each absorb 5.39 eV for a total of 10.78 eV.
(2Li X 5.39 eV per Li = 10.78 eV). You would have 2.82 eV left over, probably as heat.

I get it now. Why are we using 13.6 eV all the time?

Well, you started it! We were talking about hydrogen's first ionization energy (13.6 eV), but we could use any given energy in our calculations.
Now Arthur, just to see if you really understand this idea, tell me what you could do to Group I atoms with only 10 eV.

OK. Well, first thing I notice is that it is not enough energy to ionize hydrogen. Is that right?

Right. 10 eV will not ionize hydrogen.

But it would ionize all the other Group I elements because they all have first ionization energies below 10 eV. Right?

Absolutely. Go on.

That's it. What do you mean "Go on."?

Go on and tell me what would happen if you had more than just one atom of each Group I element.

Oh, I see. (I thought you were trying to trick me into mistakenly using the first energies twice on the same atom. That would be wrong.)
Well, to ionize TWO lithium atoms takes 10.78 eV (2Li X 5.39 eV per Li = 10.78 eV). So only one lithium atom can be ionized.
I need 10.28 eV to ionize two sodium atoms (2Na X 5.14 eV per Na = 10.28 eV), so again I am left with only one ionized atom of sodium (Na+).

You're right.
You could have gone through that more quickly if you had thought that with only 10 eV to work with you would need atoms with first ionization energies of less than 5 eV in order to ionize more than one at a time. See? (10 eV / 2 atoms = 5 eV per atom).
What you did is right! I'm just pointing out another way to think about it.
Please continue.

OK. Potassium (K) is a different story because its first ionization is only 4.34 eV, much less than 5 eV. I need only 8.68 eV to ionize two of them (2K X 4.34 eV = 8.68 eV). So that 10 eV will make two ions of potassium (2K+) with a little bit of energy left over (1.32 eV). Right?

Right!

Rubidium (Rb) needs only 4.18 eV to ionize its first electron, so 10 eV could ionize two rubidiums (2Rb X 4.18 eV = 8.36 eV) with some energy left over (10 eV - 8.36 eV = 1.64 eV).
Finally, that 10 eV could ionize two cesium atoms (2Cs X 3.89 eV = 7.78 eV) with some energy left over (10 eV - 7.78 eV = 2.22 eV). If I had a little more energy I could ionize a third cesium atom.
Your Table doesn't have a value for the first ionization of francium (Fr), but I would guess that it has a lower value then rubidium. So I might be able to ionize three francium atoms with 10 eV

Very good, Arthur. You made a good prediction. (And you're right!)
I think you have a good grasp of ionization energies and that will become very important as we progress through Alchemy.

Great! It's just simple math.

Yes, it is. It's also important to understand that ionization energy is caused by how well an atom can hold its electron and that has to do with the atom's size.

But that is only a trend not a law! Some of the Group II elements are a hiccup in the pattern.

That's right. Understand that the energy needed to ionize an atom (the ionization energy) decreases as you go down a Group, because the extra shells make it difficult to hold all the electrons. Also, ionization energy (usually) increases as you move across a Period, because the atoms hold them more closely.

Easy! What else can I learn from the Periodic Table?

Well, there's electron affinity - a measure of how well an atom ACCEPTS electrons.

Kind of the opposite of ionization energy.

Yes, as a general idea, but electron affinity is caused by an atom's ability to pull extra electrons into its outer shell using its nuclear charge.

So it is a measure of how easily an atom becomes an anion!

Yes, that's right! Elements with high electron affinity readily take electrons, when they can get them, to become anions. Those would be the elements close to fluorine in the Table.

It's that greedy fluorine again stealing electrons every chance it gets!

Yes, it is! Look at fluorine (F) on this Table! It has a value of 3.4 eV. Look here at chlorine (Cl), 3.62. What a monster!

I sort of see some trends in this Table but it isn't real obvious.

Don't waste your time trying. Generally speaking, elements with the highest electron affinities are on the right side near the top of the Table. But the pattern is far from constant around the Table.

Does it have to do with complex stuff like quantum mechanics?

Yes, it does. It certainly requires a great deal more attention to their details, but we need not worry about it. The Periodic Table is something to be used even when the details behind it are unknown. The values in the Tables have been determined by many Alchemists doing many experiments. We can use the information gained by their hard work to understand Alchemy better, but we don't need to go too deep to use it.

OK, but what do these values mean? Some are even negative!

Aye. Confusing isn't it? We could be very detailed about this kind of stuff, like we did for ionization energy. I really see no reason to do that. However, we can make some general observations.

Like chlorine and fluorine have large positive values so that means they are good at grabbing electrons.

Yes. As a matter of fact, all the atoms in Group VI and Group VII are good at grabbing electrons. Some better than others, but all of them are pretty good. What does that tell you about the ions they can make?

Ah! It means Group VI and Group VII elements are likely to become anions. Or at least be very electronegative parts of a covalent bond.

Right. Now look here at the Group II elements. All negative values. What do you think that means?

That they are all bad at attracting electrons but good at releasing them. The negative elements become cations. Right?

Well, you are half right. Negative values mean these elements are bad at accepting an electron. That much is right, but it is wrong to think that just because an atom has a negative electron affinity it will be good at releasing them.

But they all form cations, don't they? I recall that magnesium forms the cation with two positive charges. Magnesium chloride is MgCl2 and the magnesium loses two electrons to become a cation. Right?

Right. But you are focusing your thinking in the wrong way.
Magnesium becomes a cation because of its low ionization energy, not because of its negative electron affinity.

Oh, I see. Even tough they have opposite effects, ionization energies do one thing and electron affinity another.

Right. Don't feel bad. It's a common mistake.
Remember, use the ionization energies when talking about producing a CATION and use electron affinities to figure out if an ANION is produced.

It seems as if they should overlap in some opposite way. You know?

I know. It is a "gut feeling" but it isn't true and I'll show you where it really falls down. Look at the electron affinities of the Group VIII elements, the noble elements. Tell me what you see.

Negative values for all of them. Very negative. Hey, Group VIII elements never form anions! They never form any kind of ion. Not naturally.

Right. They're noble. They have no affinity for extra electrons because they are perfectly happy to have their outer shell exactly as it is. Full. An extra electron would ruin it!

Oh I see! The noble elements dislike extra electrons so much that they absolutely hate the idea of taking one on board. So they have the most negative of electron affinities. But that doesn't mean that they would be quick to give up their own electrons. Because if they did they would be ruining their complete shell!

That's right. The noble elements don't form anions or cations because they would then have incomplete shells. The Periodic Tables we've been looking at would support that idea. Notice that the noble elements have very high ionization energies (so they won't form cations) and very negative electron affinities (so they won't form anions). Look back at the Tables.

I see what you mean. So the negative electron affinity values for the Group II elements don't have anything to do with their ability to form cations. It has to do with their low affinity for extra electrons.

Right. You would have to look at the (first and second) ionization energies to see how easily magnesium loses electrons.

You know, even nitrogen (N) has a negative value, just barely.

So what does that tell you? What does nitrogen's electron affinity tell you in comparison to other elements in the Table?

Let's see. Nitrogen has an electron affinity of -0.1. So it has more electron affinity than a noble element (which are much more negative). Also, it means that all those elements with positive values are more likely to accept an electron than nitrogen.

Yes. Notice what you have been working with. Comparisons. That's the power of the Table. Pick an element and compare it to the others.

OK, oxygen. Oxygen (O) has an electron affinity of 1.46. Very high. I see that any atom with a value less than 1.46 (say hydrogen at 0.75) will not have a chance to grab a spare electron if oxygen is around. Any element with an electron affinity lower than 1.46 will have no chance at grabbing an electron before oxygen gets it. (If oxygen is around.)

Yes! Fantastic! That is the point of the Table.

An element with a higher electron affinity than oxygen (say iodine at 3.06) would get to that free electron before even oxygen could.

Yes! Precisely. I think you got it!

I can see why these two Tables are valuable. They tell you how atoms treat electrons. Their own electrons as well as spare ones. It's electrons that make bonds.

Right you are! You can use the properties of ionization energy and electron affinity to understand bonds better. Compounds too.
The strongest diatomic compounds made of ionic (electrovalent) bonds are those made from one atom with a very low ionization energy (like Na) and the other atom with a very high electron affinity (like Cl). Do you know why?

Sure. The atom with low ionization energy (Na) will readily lose its electron, becoming the cation (Na+). The atom with the high electron affinity (Cl) will readily accept an electron, becoming the anion (Cl-).

That's right. So the production of a sodium cation and a chlorine anion is "easy", according to these two Tables.

It would be nice if you had a Table showing the power of electron affinity and ionization energy all at the same time. It would be more convenient. It might be more useful.

Yes, it would and it is. It uses the idea of electronegativity.

Electronegativity! In the last lesson (WATER) we talked about how the electronegativity of an element affects its behavior. Electronegativity affects bond formation. Right?

Right. Electronegativity is a measure of how tightly an atom holds onto electrons (its own electrons or spare electrons it finds). It's a great way to learn about the electron behavior in molecules. You can use ionization energy and electron affinity to figure out a great deal about electron behaviour. But electronegativity is a value specifically designed to illustrate electron behavior in molecules and bonds.

Fantastic! Throw out the other Tables and let's learn something useful!

Those other Tables ARE useful!
You use them to study things like how salts dissolve and how plasmas work. When you start to focus on molecules it is good to find the values of each atom's electronegativity.

OK. So what is electronegativity really? How are the values made and how do you use them?

A fellow named Linus Pauling came up with the idea of representing all this "to and fro" of electrons in molecules. He analyzed the energies in many molecules and bonds to come up with a new value we call electronegativity.
The greater an atom's electronegativity, the greater the tendency for that atom to attract and hold electrons.

I see. Chlorine (Cl) is very electronegative, so it quickly picks up an electron becoming an anion (Cl-). But sodium (Na) is not very electronegative, so it readily gives up its electron becoming a cation (Na+).
Both of them do that in order to obtain the electronic configuration of a noble gas (a full outer shell).

That's right. Notice that sodium (Na) and chlorine (Cl) are on opposite sides of all the Tables. That's because they have opposite electronegativity and that displays a trend across the Periods.

As you move from left to right across a Period, the elements become more electronegative.

Hey, wait a minute.
The noble elements in Group VIII are even further to the right than the elements in Group VII. That would mean that neon should be more electronegative than chlorine, but neon doesn't steal electrons away from sodium. Neon isn't an anion. What gives?

If neon were to grab an electron, where would neon place it? Into what shell would neon have to place its new electron?

Ah, let's see. I get it!
Neon has a complete outer shell (the L-shell is full). If it were to take on another electron, neon would have to create a new shell (an M-shell) and place that one electron in it. That won't work.

That's right. It won't work because it violates every rule in the book about electron shells. The noble elements do not have a value for electronegativity because it wouldn't make sense. We still have them in the Table, because they belong to the Table. If we were to leave them out, we'd have an incomplete Table.

So, ignoring the noble elements, electronegativity increases as you move across the Periods (from left to right) and it decreases as you move down a Group.

Yes. Cesium (Cs) has the lowest electronegativity (0.79) and fluorine (F) has the highest (3.98).

Fluorine is a monster! No wonder it sucks up electrons every chance it gets. Other elements don't stand a chance!

Yes, fluorine is an "electron monster"!
This Table of electronegativity can be used to predict which compounds will be mostly covalent or mostly ionic. Two elements with very different electronegativities form ionic bonds. For example the compound cesium fluoride (CsF) is very ionic because those two elements have the biggest difference between them on the Table.

I see. Cesium has an electronegativity of 0.79 and fluorine's is 3.98, so the difference between them is 3.19.

Yes. Such a large difference in electronegativity (3.19) would mean that CsF is a VERY ionic compound. And it is! That huge difference means that one atom will be "happy" to give up an electron and the other atom "happy" to take it.

That forms two ions which are drawn together by ionic bonds! Great, but what about compounds made of atoms with similar electronegativities?

What do you think?

Hmmm. I guess that if two atoms had similar, roughly equal, electronegativities, they would be expected to form covalent bonds to satisfy Lewis' rule (to complete their outer shells). Yeah, that would make sense. Both atoms would have about the same amount of attraction for the electrons, so a covalent bond would form. Neither atom could win the tug-of-war.

Absolutely right! Imagine a molecule of methane (CH4). It is held together by C-H bonds. Use the Table of electronegativities to figure out how its bonds are formed.

OK. Carbon has an electronegativity of 2.55 and hydrogen has an electronegativity of 2.20. So the difference in electronegativity is 0.35. That's not much of a difference! It's about a tenth (1/10) the difference in CsF. So I think the C-H bonds in methane are covalent. The difference in their electronegativity is so low that neither atom can give or take an electron from the other. So, they share. They form covalent bonds.

Very good. You worked that out very well. Here's another one to try.
Consider molecules of water. How would you use this Table of electronegativity to understand the O-H bonds in water?

Well, hydrogen has an electronegativity of 2.20 and oxygen has an electronegativity of 3.44. The difference between them is 1.24. Hmmm. That's kind of in between. What is the cut off line between covalent and ionic bonds when you use electronegativity differences?

Good question. There is no "cut off" line between the two types. Instead there is an "overall" bond that you get. A kind of "average" bond. Water, for the most part, is a covalent molecule, the O-H bonds being formed by electron sharing. But these same bonds are very slightly ionic too! Sometimes a hydrogen atom will actually donate its electron to the oxygen.

That would create an ionic bond. The O-H bond would behave like an ionic bond. Right?

Right. But remember, these ionic bonds in water are pretty rare. At any one time most water molecules are covalent, but there are always a few ionic ones in the bunch.

Which ones are the ionic molecules and which are the covalent molecules? Can you separate them?

"Can you separate them?"! Goodness, Arthur you sound like a real Alchemist when you ask good questions like that. To answer your question, no, you cannot separate them.

Why not?

Because it is not as if some water molecules are always covalent and some always ionic. Each and every water molecule can switch between the two kinds very rapidly. Even if you could separate them, they would quickly turn into the other form.

I see. So this value of 1.24 (the difference in electronegativity between oxygen and hydrogen) means that most of the time the O-H bonds are covalent, but sometimes they are ionic.

That's right. Sometimes those ionic O-H bonds break, forming two ions. The hydrogen leaves its electron with the oxygen. That hydrogen ion (H+) might drift away leaving behind an anion of (OH-). We call (OH-) a "hydroxyl" ion. But they can get back together again and reform a covalent bond. These two ions (H+ and OH-) are called an "ionic pair".

You mentioned this in the last lesson didn't you?

Did I? Well, it's worth repeating. This very slight ionization of water happens all the time, but only to a small number of water molecules at any one time. Pure rain water has about one ion pair for every ten million covalently linked molecules.

Hardly worth thinking about.

That number seems very small but it has important consequences. I'll teach you about that in Advanced Alchemy. For now let's stick to the easy stuff.

OK. Is there anything else I can learn from electronegativities? Besides covalent and ionic bonds.

Aye, there is! You should notice that this Table of electronegativities helps you to understand polarized molecules and the polarized bonds that make them.

We talked about polarized bonds in the last class. The O-H bond in water is polarized. The electrons are shared between the two atoms, but they spend more time with the oxygen than with the hydrogen. So the oxygen has a slight negative charge and the hydrogen a slight positive charge.

Right. Polarized bonds leave the atoms with very slight charges which we call "delta minus" or "delta plus". They are not complete ions! These polarized bonds are due to the electronegativity of the two different elements which make the bond.
You already showed that the difference in electronegativity between oxygen and hydrogen (1.24) produces a polarized covalent bond.
Now use these electronegativity values to tell me about bonds between carbon and hydrogen in methane.

I did that one earlier. Remember? Carbon has an electronegativity of 2.55 and hydrogen has an electronegativity of 2.20 so the difference in electronegativity between them is 0.35.

Yes, you are right. What type of bond would that be?

Oh, it's covalent. Very covalent. It's even more covalent than the O-H bonds in water!

Correct. While water's bonds have a very slight ionic "character", the bonds in methane are definitely NOT ionic at all. The lower difference in electronegativity of C-H bonds (0.35) compared to O-H bonds (1.24) means the C-H bonds are "more covalent" than O-H bonds.
Now tell me, are the C-H bonds polarized?

Ah, yeah they are, but not as much as the O-H bonds in water.

Why do you say that?

Well, the O-H bonds are very polarized because the difference in electronegativity between the two atoms is 1.24. The oxygen in water will have a slight negative charge because the electron pairs that make the covalent bond hang around the oxygen more than the hydrogen. On the other hand, the hydrogen will be slightly positive because the electrons would rather be near the oxygen.
Methane's C-H bonds are only slightly polarized. The lower difference in electronegativity of C-H bonds (0.35) would cause less polarization. Right?

Absolutely right! You explained that well. Which atom in the C-H bond has which partial charge? And why?

Ah, the carbon is the more negative of the two because it is more electronegative than the hydrogen. So electrons will hang out with carbon more than hydrogen. The carbon is delta minus and hydrogen delta plus, in a C-H bond.

That's right.
I think you see how useful this Table of electronegativities can be. Electronegativity determines what kind of bonds are formed, including polarization of the covalent bonds. And polarization involving hydrogen atoms can produce hydrogen bonds. Remember?

Yeah. The polarized bonds draw the electrons away from the hydrogens, so the hydrogen could be used to make hydrogen bonds. A polarized bond will cause the hydrogen to be slightly positive (delta plus) because its shared electron spends most of the time with the other, more electronegative, atom. That slightly positive charge on that hydrogen can attract and hold (slightly) any negatively charged atoms. That's a hydrogen bond!

Right. The O-H bond (in the water molecule, for example) is SO polarized that it causes the hydrogen bond to form.
Thanks to Linus Pauling we have these convenient electronegativity values to helps use understand all kinds of bonds!

Linus Pauling. What a great guy!

Yes. I think Linus was the greatest Alchemist of the 20th century.

I now understand how to use this Table to predict the properties of atoms, but I was hopping for a kind of "tour" of the Table. Exactly what kind of properties do these Groups share? Do all the elements in a Group behave exactly the same?

Good questions.

If you want to continue choose the next hyperlink.

PRINCIPLES OF ALCHEMY
EARTH

Part Two


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