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PRINCIPLES OF ALCHEMY
EARTH

Part Two

Let's do a "tour" of the Periodic Table starting with hydrogen. It is an unusual element because it has only a single electron.

Why does that make it unusual?

No other atom finds itself in the unusual situation of being able to form covalent bonds, or ionic bonds of EITHER ion.

What!? But hydrogen forms covalent bonds with oxygen to form water. What's this stuff about it being "either" ion in an ionic bond?

You are right. Hydrogen forms covalent bonds, but it also forms ionic bonds. And when I say "either" ion, I mean that hydrogen can be either a cation or an anion.

Hold on there! How can hydrogen do so many different things? All it has got to work with is a single electron.

That's right, and that's the reason it can do so many things. Let's take them one at a time. You mentioned that hydrogen forms covalent bonds with oxygen in water molecules, but you also know that those bonds are slightly ionic. Right?

Yeah, yeah. Pauling's electronegativity Table shows why. Oxygen's electronegativity is 3.44 and hydrogen's is 2.20, so the difference (1.24) means that the electrons will be shared between them most of the time. But a few of them will be ionic bonds. We don't know which are ionic and which are covalent, just that most are covalent.

Right, and we can't separate the water molecules into covalent and ionic ones because they keep switching from one to the other.

Water is SO complex! And I thought we had heard that last of it with the last lesson.

No Alchemist ever heard the last of water! It's such an amazing molecule. Even in the 20th century Alchemists study it. However, the simple properties that I want to teach you are easy to understand. I think you will agree that the O-H bonds in water molecules can be either covalent or ionic.

Yeah, but mostly covalent.

Right. Now tell me what kind of bond would you expect to be holding together a molecule of hydrogen fluoride (HF). Just look at the Periodic Table and calculate the difference in electronegativity, like....

Ionic! I don't have to look at the Table because I know that fluorine is an electronegativity "monster". It steals electrons all the time.

A good guess and the right one. But remember that you MUST calculate the difference in electronegativity in order to figure out the bonds. Don't go by instinct (unless you must!).

OK. OK. The difference in electronegativity for HF is 1.78. That's probably a big enough difference to make an ionic bond. Right?

Right. Like I've said before, there is no clear line between ionic and covalent bonds. Only a gradient (a continuous transition) from one to the other. Let's say, for the sake of argument, that 1.5 is the cut off. I'm just making up that number so we can imagine it more clearly. Do NOT take 1.5 as some "all powerful" number in electronegativity differences. It is not.

But it IS convenient. That means the HF molecule is an ionic molecule.
Hydrogen releases its electron to become a cation (H+) and fluorine sucks up an electron to become an anion (F-). They are drawn together and held together by electrostatic attraction.

Well said! Very good. Later, I'll give you a question or two about fluorine and the bonds it makes. Remember, fluorine may well be an electronegativity "monster" but it is rash to quickly say it will always be ionic. Fluorine is not always ionic. Sometimes it is bonded covalently. It depends on the electronegativity of the other atom.

OK, I'll keep that in mind and use the electronegativity differences all the time when I have to predict a bond. (I'll use 1.5 as my cut off just to keep things simple.)

Good, now back to hydrogen!

Yeah, you said something about hydrogen being able to form ionic bonds as EITHER ion. What did you mean by that?

I mean that hydrogen, under certain circumstances, can GAIN an electron to become an ANION (H-). This unusually ion of hydrogen is called a hydride (pronounced "hi-dried").

Why would it do that? How could it do that?

It does that to fill its outer shell. Remember?
Hydrogen has one electron in its only shell, the K-shell. But the K-shell can hold two electrons.

Oh, I see. It grabs an electron to fill its outer shell! To take on the shell structure of a noble element and that is what makes it a hydride ion!

Right. To complete its outer shell it takes on an extra electron. It's that desire for a complete shell that drives atoms' bonding behaviour. A hydride ion (H-) forms when hydrogen tries to complete its outer shell. Just like the formation of a fluoride ion (F-), except hydrogen doesn't do it as well as fluorine.

So when does the hydride form? It must be when the other atom has a much lower electronegativity. Right?

Absolutely right. Hydrides (H-) form when the electronegativity of the other atom is much less than hydrogen's. Take a look at Pauling's electronegativity values and pick out an element that you believe will cause hydrogen to form a hydride.

OK, Lithium (Li) has an electronegativity of 0.98 while hydrogen's electronegativity is 2.20. A difference of 1.22. Hey, that's not enough to form an ionic bond. That's less than 1.5!

Now Arthur! What did I say about that number?

Oops! It is NOT a true cut off between the two types of bonds, just a rough guide.

Right. Hydrogen and lithium form ionic bonds between them, with the hydrogen being the anion (the hydride, H-) and lithium being the cation (Li+). We call this compound (LiH) lithium hydride.

It seems so.... backward.

Aye. I see what you mean. Notice that the compound appears to be written and spoken "backward". That is just the rule we use in naming ionic compounds. Remember?

Yeah, when naming an ionic compound we always put the cation first and then the anion.

Most of the ionic compounds of hydrogen are like hydrogen fluoride (HF) or hydrogen cyanide (HCN). Hydrogen comes first because it is the cation (H+). Most of the time.

But when hydrogen becomes the anion (H-) it goes to the end of the molecule's name, like all anions. Also, it takes on an ""ide" ending (like all anions) so we call it a hydride.

Right. There's really nothing "backward" about it. Except that hydrogen is behaving "backward". We still use the rule of cation first and then the anion. Always!

So it is called lithium hydride (LiH). Not hydrogen lithide (HLi). Right?

Absolutely right.
Hydrogen's unique electron "situation" means it can be a cation (H+), an anion (hydride H-) or covalent! It all depends on the OTHER atom's electronegativity!

So where does hydrogen go in the Periodic Table? Group I because it has only one electron in its outer shell? Or in Group VII because it needs only one electron to complete its outer shell?

Do Groups have to do with the number of electrons the atom has in its outer shell or the number of electrons in its outer shell that keeps it from reaching its noble state?

Ohh... The Groups should be arranged according to the electrons it wants in its outer shell. It's that "need" to have a complete outer shell that determines the atom's Alchemy. So the Table should be arranged to show how they try to get to their noble state.

Right! If the Periodic Table is to be useful, it must reflect the Alchemy of its atoms.

So it should be placed in Group VII along with fluorine and chlorine. But that seems all wrong! Its radius and other properties make me think it should stand above lithium in Group I.
I don't know. I'm so confused.

You are in good company with the rest of the world of Alchemy! Hydrogen can be in either Group I or Group VII. It really depends on the conditions the hydrogen is in. Over the years Alchemists have argued and argued about where hydrogen should be on the Table. We still argue about it.

Why not a compromise? I would but hydrogen somewhere between those two Groups in a place of its own.

You have the wisdom of Solomon, and you are not alone. Most Alchemists display hydrogen somewhere in between. I like to move it closer to the Group VII elements because that places it near the elements that form covalent bonds too. I just remind myself that hydrogen is nothing more than a proton! Usually.

But hydrogen may also have some neutrons.

Yes. Tell me, Arthur, would the isotopes of hydrogen, deuterium (2H) and tritium (3H), have the same properties as just normal hydrogen (1H)?

Yes! The chemical properties of the elements are a function of the protons.

Correct. Only the mass is different. Isotopes (of the same element) have different mass but the same number of protons. So isotopes of the same element will have the same Alchemy.

I can see why we spend so much time talking about hydrogen. It's the most common atom in the universe and has the weirdest chemistry!

Aye, it does. But it is chemistry you can understand. Now that we have tackled the most difficult element in the Table, the rest is easy! So let's move on.

OK. Let's hear about all the other elements.

Good idea.

The Periodic Table can be divided into two types of elements. Metal and nonmetals.
Do you recall what a metal is?

Yeah, metals form metallic bonds. Their outer shell electrons are kind of loose and they form a "sea of electrons" instead of specific bonds. This "super sharing" of outer shell electrons is a metallic bond.

Yes, very good. The most important metal property is the ability to conduct electrons.

It's that "super sharing" of electrons. That allows metals to pass electrons along very quickly, producing a current of electrons.

Right, but there are other metal properties. Remember malleability and ductility?

Oh yeah, I forgot.
Metals are malleable, meaning they can be hammered or pressed into thin sheets. And they are ductile, meaning they can be drawn into a wire (by pulling it through a hole).

Right. Metals also conduct heat very well. It has to do with those same properties that help it conduct electricity.

Yeah! A cold piece of iron really pulls the heat out of your hand when you hold it. You get very cold holding cold metal!

That's right. Metals conduct heat very well and that is another property of a metal.

And they are shiny! Metals have a shiny look to them.

That's right. Pure metals have a high "luster", which is what an Alchemist says instead of saying it is "shiny".
All these are metal properties - conduct electricity and heat, are malleable and ductile, and have a high luster.

So we divide up the elements into metals and nonmetals depending on whether they have these properties?

Yes, we do. It is a major division of the Periodic Table.

But we haven't talked about metals yet. These tiny Tables you have shown me don't even have gold or iron on them. They're metals, right?

Yes, they are. We will leave those types of metals until last. Let's stick to our "simple" table.

But there are no metals in the simple Table.

Oh yes there are!
All the Group I and Group II elements are metals and so are some of the elements in the most of the other Groups! Let's start our tour of the "simple" Table and you will see what I mean.

OK. Let's start with Group I.

A very good place to start. I'll return the hydrogen to Group I (because that's how it was in our original Table and where you find it in most Tables), but the properties we are now going to talk about have to do with the Group I elements below hydrogen.

OK. (The hydrogen will just remind me we're on Group I.)

All Group I elements have a single electron in their outer shell and form cations readily.
(Even hydrogen!)
They also have excellent metallic behaviour in their pure form.
(But hydrogen only displays these metallic properties at extreme pressures, like on Jupiter!)

I've never thought of sodium as a metal. I just think of it as a salt.

An easy mistake to make. Salt (NaCl) has different properties than the PURE elements from which it is made. In our tour of the Table always keep in mind that we are talking about the behavior of the PURE elements.
Pure sodium conducts electricity and heat very well. Pure sodium is malleable and ductile. Pure sodium has a high luster.

I see. Pure Group I elements are metals because in their pure form they have all the properties of a metal. However, the compounds they make might not.

Exactly. When discussing the properties of elements you must always remember we are talking about the pure ELEMENTAL form. Not a compound.
We use these properties of the pure elements to predict how they behave and form compounds. However, it is the elemental properties that tell us about the compounds they can make. Do you see what I mean?

I think so. The Periodic Table is a Table of pure element properties. We can use it to predict how they form compounds with each other (in elemental molecules) or other elements (compound molecules).

Right! All The Group I elements are metals, in their pure form.
All of them are also "alkaline".

"Alkaline"? What's that?

Alkaline is a property of some elements. It means that when added to water, it causes the water to have an excess of hydroxide (OH-) ions.

How does sodium do that?

Well, that is a problem I'd like to leave for another time.
You should understand that Group I elements rapidly form hydroxides. That is, sodium (Na) will quickly become sodium hydroxide (NaOH) when placed in water or even in moist air!

But the air is always "moist". How can you keep sodium pure?

It isn't easy. You have to keep the sodium under a layer of oil to keep the air away.
We say that sodium has a high "reactivity" with water, meaning it reacts quickly with water.

How "reactive" is it? Does it happen in a flash?

Yes, it can happen in a flash. If you drop a piece of pure sodium into water it forms sodium hydroxide so quickly that it actually fizzes and gives off heat and light in the reaction! It is very impressive.

What if you just let the sodium sit around in moist air. Does it fizz?

No, but it does "tarnish". That means it loses its luster. It loses its luster because its outer surface, the part of the sodium in contact with the moist air, reacts to form sodium hydroxide (NaOH).

So it is no longer a metal. The outer surface I mean.

That's right. The outer surface of a piece of sodium in contact with (moist) air is no longer sodium anymore because it has formed a compound of sodium hydroxide (NaOH).

I see. A piece of sodium in (moist) air will just tarnish, but if placed into water it will actually fizz away, all of it becoming sodium hydroxide, not just the surface.

Very good, Arthur. You have pictured it well. The tarnished layer (made of NaOH) forms a barrier to the air, so water in the air can't get deep into the piece of sodium. But if you drop it into a bowl of water, the tarnish doesn't have time to form. Just too much water to keep out!

If you cut open a piece of tarnished sodium you would find pure sodium underneath. Right?

Right, but you would have to look very quickly at the pure sodium before it tarnished! That's why we Alchemists keep pure sodium under a layer of oil. The oil acts as a barrier keeping the water out and thus the sodium pure. Even in very dry air, like in a desert, sodium will quickly tarnish. There always seems to be some water in the air no matter how dry it feels to us.
All the Group I elements tarnish quickly because all the Group I elements rapidly form hydroxides (something-OH) and those hydroxides are very alkaline (have lots of OH-).

You said they do this quickly. Are all the Group I elements equally quick.

No. The speed and efficiency of this reaction changes. Their reactivity increases as you go down Group I. When exposed to air, lithium (Li) takes on a slight tarnish and does so rather slowly. Sodium is quicker. This reactivity continues to increase (speed up) as you go down Group I. Pure cesium (Cs) will react so quickly that when placed in moist air it actually fizzes.

Like sodium when placed in water! What happens if you drop a chunk of cesium into a bowl of water?

What do you think would happen?

I bet it would react VERY quickly.

You are right. Dropping cesium into water causes a massive explosion!
It is an extremely violent and dangerous experiment.

Wow, an explosion! Does it explode because it reacts so quickly?

Aye. The reaction is so rapid that the energy of the reaction overwhelms everything.
All explosions are caused when too much energy is released too quickly.

Well, I won't be dropping cesium into water!

Good. Actually, pure cesium is hard to make.

And once you made it you better keep it under oil. Away from water or (moist) air!

Right. In fact all the Group I elements are stored under oil. All Alchemists know you never put a Group I element into water.

Unless you WANT an explosion!

Aye, but nobody wants that! Lithium will causes the tiniest explosion and cesium the greatest because reactivity increases down this Group.

I bet that has to do with the fact that their electronegativities and electron affinities decrease down the Group.

Right! The larger shells mean the outer electrons can react more quickly to form bonds and make hydroxides of them.

Does anything else about them change as you go down Group I? Any other properties?

Yes. All Group I elements are very soft metals and that softness increases as you move down the Group. That has to do with the way each atom is bonded to its neighbor atoms. (All of the same element, of course, because we are talking about pure elements.) The larger shells allow the metallic bonds to slip more easily past each other.

So it would be easier to cut open a piece of cesium than a piece of lithium, but it would also be more dangerous!

Yes, that's right.

Great. Anything else I should know about Group I elements?

Yes, you should know that all the Group I elements are also called the "alkali metals".

Because they are all strongly alkaline (rapidly forming hydroxides) and they are all metals (having all the properties of metal). Right?

Right. Very good. Alchemists call the Group I elements the alkali metals, because that is exactly what they are!.

Easy. Lets move on to Group II.

Good idea. You will be happy to know that Group II really is easy to learn now that you have a good understanding of Group I. What do you know about Group II elements?

Ah, well they have two electrons in their outer shells and therefore from cations with a +2 charge.

Very good. The only thing I would add about the Group II elements is that they are like Group I but less so.

What is THAT suppose to mean?

It means that they also form excess hydroxide ions (OH-) and have good metallic behavior. But their reactivity and metallic behaviors are not as good as Group I.

So Group II elements have the same important properties as Group I, but do not produce as much hydroxide as Group I or conduct electricity as well as Group I. Right?

Yes. Group II elements are kind of "mini" alkali metals. Instead of calling them "mini-alkali-metals", we call the Group II elements "alkaline metals". The "ne" at the end reminds us they are less than a full alkali metal (Group I).

Kind of confusing isn't it. I mean the Group I are more alkaline, but we call them "alkali". However, the Group II elements are less alkaline but we call them simply "alkaline". Why have Alchemists made it so backward?

I agree that it is unfortunate that these two names got into use the way they did. It has to do with the history of Alchemy. You see Alchemists studied the Group II elements and named them before they studied the Group I.
Everyone called magnesium (Mg), calcium (Ca) and other Group II elements the "alkaline metals" because they were the first metals found to have the properties of being "alkaline". Later on, Alchemists discovered that Group I elements were even MORE alkaline, but they couldn't call them the "more alkaline metals".

Why not?

Because it would sound stupid. So they called Group I the "alkali metals", removing the ending.

So the "alkaline metals" are alkaline, but the "alkali metals" are even more alkaline! Too bad they discovered this all in reverse. We could have had an easier way to remember their names!

True. However, I'm sure you can remember that "backward" way of naming them.

Yeah, I guess so. Why did Alchemists start their work with Group II elements anyway? Makes sense to start with Group I.

Well, it might make sense looking back on it, but not in real practice.
Imagine you were starting to work out all these properties of pure elements. If you think about it enough, I bet you will understand why Alchemist first studied Group II elements. Think about it. Think about their properties and how that might influence their decision to start with Group II. From what little I have told you about Group I and II elements, you should be able to make a good guess.

Hmmm
To study these properties you must have them in pure form. Right?

Right (I think you're on the right track.).

Hmmm
The only thing you told me about Group II is that they are like Group I but less so. Less metallic and less alkaline than Group I elements.

That's right. Keep going. I think you are almost there. Think about the real life situation these Alchemists must have been in as they learned Alchemy from Mother Nature.

Hmmm.
They would have to isolate pure forms in order to study them.
OH! I got it. Group II elements are less reactive than Group I elements. They would form hydroxides more slowly. Yes, that's it. In the (moist) air it would be easier to make and keep the Group II elements pure than the Group I elements. That's it!
The early Alchemists could make and keep pure Group II elements but the Group I elements were more difficult to make pure and keep them pure!

That's right! All the Group II elements react less violently than their neighbor on the left. Beryllium (Be) is less reactive than lithium (Li). Magnesium (Mg) is less reactive than sodium (Na). This is true of all the comparisons of the first and second Groups.

So they studied the Group II elements before the Group I elements because it was easier!

Aye. That's it. The Group II elements are like their neighbors to their left (Group I) but not so extreme.

I see. Do the same trends continue down Group II? Like reactivity and softness increasing as you go down the Group II elements?

Yes, they do. Beryllium (Be) is the hardest and least reactive of the Group II elements. Magnesium (Mg) is softer and more reactive than beryllium, but magnesium is harder and less reactive than the element below it, calcium (Ca). The trend continues.

That's easy to remember. It's just like the Group I trends (but less so).

Right. Now you have a way of comparing all the Group I and II elements. At least with respect to their ability to be cut or react. Tell me, would magnesium (Mg) tarnish faster or slower than sodium (Na)?

Hmmmm
Magnesium would be slower to tarnish than sodium because sodium is to the left of magnesium. Is that right?

Yes, that is right and the right way to think about it. The reactivity of a Group I element is always greater than the next element in the Period (to the right, the Group II's).

So, in any comparison of Group I and II elements, reactivity increases down the Groups and decreases across the Periods.

Right. That all has to do with properties of the outer shell which we have been talking about since the very beginning of our lesson about the Periodic Table.

I see that the Table IS very useful. It displays these trends very well. All you have to do is remember is which direction they change.

Yes. Because you know how atomic radius, electronegativity and other properties change as you move around the Table, it is a simple matter to keep track of these changes.

Let me try this.
I would guess that calcium (Ca) is more reactive than magnesium because calcium is below magnesium. But calcium is less reactive than potassium (K) or strontium (Sr) because those elements are to the left (K) or below (Sr) the calcium. Right?

Right! By Jove I think you've got it!

It's kind of like a staircase. As you go down or to the left the reactivity increases.

That is an excellent observation, Arthur. You will see staircases a lot in the Table. As a matter of fact, we see staircases as we continue to move from right to left across the Table.

Look at this Periodic Table. I've drawn it to show the metals, nonmetals and semimetals.

The staircase if obvious. It looks like a staircase of metal with a carpet of semimetals and then a layer of all the other elements

Yes, that is exactly the way it looks to me too. You will notice that elements in Group III and beyond are either metals, nonmetals or semimetals.

You know, this staircase seems to start just after the Group II elements. It's as if something special happens there. Isn't this where those missing elements belong? You know, the gap in atomic number between calcium (atomic number 20) and gallium (atomic number 31).

Yes, you are right. Something very strange happens at this point in the Table. That "something strange" are the transition elements. I've left them out for now because they are so strange. We will discuss them last in our tour. All the other elements, from Group III to Group VIII, are called the "post-transitional elements".

Because they are "past the transition elements", in the Table?

Right!
Metals in the Groups III and onward are the "post-transitional metals". Notice how they are arranged.
Aluminum (Al) is at the top of the metal staircase in Group III, Period 3. The staircase goes down as you move right and has polonium (Po) at the bottom in the position of Group VI, Period 6.

What causes this staircase?

That is a very good question. It is clear to see that as you go down these Groups the metallic properties of the elements become stronger. That fact is as true of the post-transitional elements as it is for the "pre-transitional" ones (the alkali and alkaline metals).

Metallic properties increase as you move down the Groups.

Right. The electron shells added as you go down the Groups cause the elements to behave more like a metal. They will give up their outer electrons more easily allowing them to be "super shared". BUT the additional protons added as you go across a Period counter that metallic behaviour.

I remember. When you add more protons to the nucleus it causes the electrons to be pulled more closely. You are adding electrons too, but you are adding them to the same shell. That means the whole shell is pulled closer as you move across a Period. So the outer electrons are more difficult to remove as you move right.

Right. So those electrons are held more tightly and will not get involved in conducting electricity. At the far right of the table each outer shell is drawn so tightly to the nucleus that those elements have no metal properties at all. They can't conduct electricity or heat. They aren't malleable, ductile or shiny.

The only element among the post-transitional metals that I recognize is lead (Pb as in "plumber"). That's the metal used to make pipes! Right?

Right. In our ancient world, lead was the only post-transitional metal to be mined. It is easy to work with and easy to make into pipes.

So it's the most important post-transitional metal!

Well, yes, in our medieval times. However, aluminum (Al) is the most important metal in the 20th and 21st centuries. Aluminum is the most abundant metal on earth. Or on the moon! It is the best metal in the world for manufacturing and construction.

I've never heard of it!

No one has in the Middle Ages. Although aluminum is found in most rocks, it is VERY difficult to purify.

Why?

Aluminum is a fairly active metal. It binds to the oxygen in the air and holds it very tightly. Rocks full of aluminum are full of aluminum oxides.

What do you mean by "oxides"?

Well, you recall carbon dioxide (CO2), right?

Oh yeah.

And dihydrogen oxide (H2O)?

Huh?

Water. Dihydrogen oxide. (H2O)

Oh, yeah! Right.

Well oxide just means it that it is bonded to oxygen. Aluminum oxides are any aluminum bonded to oxygen. The exact number of oxygens is not important. What IS important is that all the aluminum you find is contaminated with oxygen. We call those rocks of aluminum oxide "bauxite". Bauxite is "rusty" aluminum.

"Rusty" just means it has oxygens bound to it?

Actually, only iron (and steel) "rust". Alchemists can be pretty particular about what they mean by "rust". So don't take what I said about "rusty" aluminum too seriously. It's slang. The aluminum isn't really "rusty", it just has oxygens bonded to it.

Because only iron can rust.

Right

So how did people in the future make pure aluminum from bauxite?

At first they used a variety of chemicals, but that produced only small amounts of pure aluminum. Then, in 1886, a fellow named Charles Hall found a way to use electricity to push electrons into aluminum's outer shell, forcing it to release its oxygen.

Hall used electricity to do Alchemy?

Aye! It was a major break through and a major industry was born. Pure aluminum can be pressed into any shape with just a little effort. Aluminum can be hardened into a fine material for construction by using a few processing "tricks". Aluminum is very light (in mass) and a perfect metal for building things.

Great. But wait a minute!
If aluminum is so reactive and picks up oxygens, then wouldn't these aluminum buildings and stuff "rust" away?

That is an excellent point. Yes! Aluminum "rusts" in the air! It doesn't happen as fast as the immediate tarnish that develops on Group I elements. But what does that tarnish do for the underlying metal?

The tarnish protects the metal beneath from reacting with the air. It forms a shield from the air.

Correct. A similar protection occurs with aluminum. Aluminum is reactive but it forms a protective layer of aluminum oxide on its surface. That thin film of aluminum oxide keeps all the aluminum underneath it in the pure state. The film is so thin that you don't normally notice it.

Aluminum is a "super metal". It's light and can be made very strong. But it's too difficult to make pure without lots of electricity. I see why it had to wait until the 20th century before it caught on. You once told me that electricity didn't really catch on until the 20th century.

That's right. Let's get back to the Table and continue our tour.
Slightly above the metals staircase is a carpet of elements we call the "semimetals", because they have semi-metallic properties. They all conduct electricity but not nearly as well as the "real" metals.

I suppose that is the effect of this staircase structure. They are not fully metals or nonmetals. They are in between.

Right. The conductivity of the semimetals is very much influenced by specific molecular arrangements they might make or small amount of other elements added to them.

Can you give me an example?

Sure. Carbon is a semimetal, although most of the time we think of it as a nonmetal. But one form of carbon, called graphite, conducts electricity. Graphite is carbon arranged in flat, molecular sheets. It carries electrons in directions parallel to the carbon sheet but not in other directions.

So carbon is an example of an element that is usually not a conductor, but which can conduct electrons in certain ways if arranged in a particular molecular structure. So carbon is a semimetal.

Aye.
Silicon, right below carbon, is also a semimetal. When other metals or semimetals are added to silicon it begins to conduct electricity. We call that a semiconductor.

Because it conducts only when mixed with another metal or semimetal?

Aye. The exact reason why semiconductors conduct electricity is a very complex topic. At this level in your Alchemy education I think you should understand that the electrical properties of semimetals can be increased by adding wee amounts of metals or other semimetals. Let's not bother with how or why.

OK. I guess that's all that can be said about semimetals. It all has to do with their "in between" electrical properties.

Yes. However most of the semimetals also display slightly the chemical properties of a metal. Like the ability to produce alkaline solutions. That's a chemical property, not a physical one. Some Alchemists call those semimetals "metalloids" because they display some chemical properties of a metal.

I'm confused.

I'm not surprised.
Many Alchemists are confused too!
The difference between a semimetal and a metalloid is not really a property of the Table or the element's position in it. It has to do with the properties of EACH element. Metalloid refers to the chemical properties. Semimetal refers to the physical properties.

So the elements called semimetals in the Table have the semi-electrical properties we talked about, but not necessarily the chemical properties?

Right. The "not necessarily" part is important. Most semimetals are metalloids. But not all of them and certainly not under all conditions.

Let me get this straight. The Table lists semimetals very clearly. But that is based on their physical properties, not their chemical ones?

Right. I agree that it is odd to use physical properties to define elements in a "Chemistry" Table. I didn't make up the rules. I just teach them!
This part of the Periodic Table can trip you up because it is a fuzzy zone where elements display the chemical (metalloid) or physical (semimetal) properties of metals under certain conditions.

I'll learn not to trip on the semimetal carpet.

Good! So let's move on to my favorite part of the Table. The nonmetals.

I bet they are bad at conducting heat or electricity, and are not malleable, ductile or shiny.

Right. Except carbon. Remember what I said about graphite?

Oh yeah. Carbon's a semimetal.
Hey! You just said it was a nonmetal! Are you making this up as you go?

No, I'm not! Carbon is an amazing element. It sits in a very important part of the Table, where all the important properties of elements seem to come together. I keep coming back to the "fuzziness" of this part of the Table because it gets complex.

Sure does!

Carbon, like all the Group IV elements, has four electrons in its outer shell. Both carbon and silicon can share these outer electrons to form covalent bonds.

Hey, I though silicon is a semimetal. It should form semimetal bonds. Shouldn't it?

Well, silicon DOES form metal bonds. However, there's no such thing as "semimetal bonds". Either you "super share" the outer electrons or you don't. Remember silicon is in that fuzzy area of the Table where the definitions can be confusing. Sometimes silicon "super shares" its outer electrons, in which case it is behaving like a metal. At other times it might share its outer electrons in a covalent bond. It depends on specific conditions.

Oh, I hate this fuzzy area of the Table! It seems anything is possible in it!

Yes, it seems that way. And that's what makes it so important. Lots of amazing Alchemy occurs here. Let's spend some time talking about carbon. It is a very important element.

Why?

Because carbon is VERY good at making covalent bonds. By doing so, carbon can make a huge number of very large, very complex molecules.

More complexity! I don't think I can take much more!

Don't worry. The Alchemy of carbon is such a complex and important subject that it deserves a course in its own, Organic Alchemy. We won't go into that now.

Phew!

All you need to know about carbon is that it is the foundation of all organic molecules and it has three allotropes.

Allotropes? Remind me again, what's an allotrope?

An allotrope (pronounced "al-oh-trope") is an element in a different form. It has a different arrangement of atoms. Carbon has three allotropes; graphite, diamonds and fullerenes. They are all made of carbon - nothing but carbon. Yet they can be arranged in different ways.

You said that graphite is carbon arranged in sheets. That's why it can conduct electricity (in the same direction as the sheets).

Correct. Graphite is just sheets of carbon stacked on top of each other. Like sheets of paper in a book. Speaking of books, graphite is what most people call "lead". It's the stuff in pencils.

You know, I always wondered about that. I use to chew pencils but was told the lead would poison me! So you can't get lead poisoning from chewing pencils?

That's right. You can't. Pencils don't have real lead (Pb) in them. They have graphite (C). The minerals of lead and graphite look alike to the untrained eye. I think that's why people got confused and named graphite "lead".
(But it's still not a good idea to chew pencils because the paint on them might be a poison.)
Layers of graphite slide past each other very easily and some layers of graphite are left behind when you drag the graphite across paper.

Is that how a pencil works? By leaving a trail of graphite on the paper?

Aye.
Graphite is an allotrope of carbon. It is carbon bonded to other carbons in sheets.

How?

How? Well, each carbon shares three molecular orbitals with three other carbons. Imagine a carbon bonded to three carbons in a perfect triangle. and each of those carbons is bonded to other carbons the same way. And so on.

I think I can imagine it. Triangles are flat, so the whole structure is flat. Like a sheet.

Yes, exactly.

Hmm. What about that fourth electron? Carbon has four electrons in its outer shell, so it should try to make four covalent bonds not just three.

That's right. The fourth electron doesn't get involved in a bond. At least not in graphite. Remember, just because it wants a complete outer shell doesn't mean it will always get one. A lot depends on the conditions it is in, and a lot depends on how it will get there. But that's a topic best left for later when I teach you about chemical reactions.

OK, so graphite is a sheet of carbons in flat triangle patterns with an extra electron which is not involved in the bond.

Right. That electron not involved in the bonds is said to be "delocalized". It just floats above the carbon plane. As a matter of fact, it helps keep the sheets a wee bit apart.

Sounds complex.

It is. But, here's something you may understand. The carbons in graphite are bonded to three others in a flat triangular plane. Carbon uses its s and p orbitals to form a special hybrid orbital called an sp2 orbital. All sp2 orbitals make flat triangles (because of VSEPR theory).

Oh, no! Not back to VSEPR theory!

Relax, Arthur. I'm not going into it any further. I just waned to point out that sp2 orbitals are used to make flat sheets of carbon.

OK. What about the other carbon allotropes?

Diamond is another carbon allotrope. Diamonds are NOT flat sheets of carbon. They are beautiful crystals of carbon linked together in ridged three dimensional structures.

Would that be a tetrahedron?

Aye, you're recalling your VSEPR theory.

Well, yeah I guess it has its uses. Are the carbons in diamonds using sp3 orbitals?

Yes, they are! You see, you are getting a grasp of molecular shapes.
The carbons in diamonds are linked together in all three dimensions to other carbons, which are linked in all three dimensions to other carbons, and so on and so on. This produces a large crystal of equally bonded carbons.
It is this three dimensional crystal structure, equally bonded in all three dimensions using sp3 orbitals, that makes diamonds the hardest substance on Earth.

I see. These two allotropes of carbon, graphite and diamond, have different properties because of the way they are arranged.

Aye.

What was the third carbon allotrope you mentioned, "fuller-than-something"?

Fullerenes. They are another three dimensional arrangement of carbon but not the flat-sided crystals of a diamond. Fullerenes are carbons arranged in the shapes of cages, shells and tubes. Fullerenes are more "distinct" as a molecule.

What do you mean by distinct?

Graphite and diamonds are bonded to all their neighbors in a long and continuous network of covalent bonds. They have no specific end to them. They come to an end at any point in the chain. But fullerenes have distinct sizes and shapes. The most common fullerene is a sphere of 60 carbons which some folks call "buckyballs".

Hmmm....
Let's not get into the VSEPR theory of them. I think you're sneaking a bit of Organic Alchemy into this course!

Oh, goodness no! I wouldn't do that. So let's move on to the rest of the nonmetals.

Good, and away from this fuzzy part of the Table. It makes my head hurt!

OK. OK.
All the Group V elements have five electrons in their outer shell. What does that tell you about their bonding?

With only five electrons in their outer shell they need three more to complete it. They would probably make three covalent bonds.

Ah, but....

EXCEPT (you didn't let me finish!),
Except for the two metals further down (antimony, Sb and bismuth, Bi) which would form metal bonds by "super sharing". And then we must not forget that awful semimetal (arsenic, As) which makes me dizzy just thinking about it!

Yes, you are right. Sorry about your dizziness, Arthur.
You know, we live in a complex universe and some parts of it require more brain power to understand it. The staircase in the Table is every Alchemist's nightmare and dream!
Take your time and keep in mind what is happening in this part of the Table, and you'll do just fine.

I'll try my best.

Then you'll do well.
You'll be happy to know that nitrogen (N) and phosphorous (P) are the only Group V elements you are likely to deal with.
Nitrogen (N2) makes up 80% of the air we breathe and is a VERY stable molecule. It's practically inert. Can you guess why N2 is such a stable molecule? Think about its bonds.

N2 has three covalent bonds. That's a lot of bonds between just two atoms. It's N2's triple bond that makes it so stable.

Right! The air is full of nitrogen gas (N2).

What about phosphorous?

Ah, now phosphorous is very different! Pure phosphorous has two principal allotropes, white phosphorous and red phosphorous.

Meaning they are made of nothing but phosphorous atoms, but arranged differently. That's an allotrope.

Aye. White phosphorous is a VERY dangerous allotrope of phosphorous. It is extremely poisonous and when exposed to the air it will catch fire!

What!? How?

Set a piece of white phosphorous out in the air and in a few minutes it will start to burn, sucking in the oxygen to form....

Phosphorous oxides!

Right! As it forms these oxides it gives off a great deal of heat. It actually catches on fire without you lighting it. We Alchemists say it "spontaneously ignites".

Wow! How can you store it?

Good question. I bet you can think of the answer. We've talked about ways to keep highly reactive elements pure.

Oh, like the big alkali metals. Put them under oil.

Yes. Oil will protect the white phosphorous from the oxygen in the air. So will water.

Why don't we use water to protect the alkali metals from reacting?

Think about it, Arthur.
What do the alkali metals react with? What does the phosphorous react with? They're different.

Hmmm... Oh, I remember!
Alkali metals react with water to make metal hydroxides (metal-OH). If you put them in water they would react all the more quickly! The big alkali metals would actually explode!

Right! So instead we use oil to isolate the alkali metals from the moist air. With the alkali metals, water is the problem. Oil keeps the water out.

But with white phosphorous it is the OXYGEN you want to keep out. So water is OK for storing white phosphorous. But not for storing alkali metals!

Yes. Very good. White phosphorous in a bucket of water is isolated from the oxygen in the air. So phosphorous oxides can't form. No spontaneous ignition.

And the water would stop a fire if it did start!

Why, yes! I hadn't thought of that. Good point, Arthur.
Now, red phosphorous is another matter. It is an allotrope of phosphorous but it is perfectly safe. It isn't poisonous and it won't spontaneously ignite.

What makes the two allotropes different? It must me the way they are arranged.

You've answered your own question. White phosphorous is a crystal - all the atoms arranged in a regular pattern. That pattern helps move oxygen quickly into the piece of phosphorous.

Like the fuse on firecracker!

Yes. The crystal structure of white phosphorous moves the oxygen in very quickly.
But red phosphorous is amorphic.

Remind me again what is "amorphic"?

"Amorphic" means "without form". We use that word to describe a solid that has no repeating patterns.

Oh, right. I remember. If the pattern of the atoms or molecules repeated then that solid would be a crystal. Like quartz. But amorphic solids have no regular arrangement. Like glass.

Right. Red phosphorous is amorphic, a "glass". There is no pattern to the arrangement.

I see. Without the crystal structure the fast chemical reactions can't occur.

Right.

These allotropes are really amazing things. I expected elemental molecules to be very dull and just one type because they are all made of the same element. Allotropes are exciting because the way they are arranged causes their properties to be very different! I didn't expect that.

Yes. Allotropes are important variations on the same elements and they are prevalent in Groups IV, V and VI.
Now let's continue our tour. We still have a lot of elements to cover.

Like oxygen (O) and sulfur (S) in Group VI. I bet they have a covalency of two because they need to share only two electrons to complete their shell.

Absolutely right! Ah, well, ah...

Are you going to say "ah well, sort of" and then tell me something to confuse me?

I'm really trying hard not to, Arthur.

Let's have it. I can take it.

OK. Just a minor point really. The four elements at the top of Groups V and VI have some peculiar tricks about their bonds.

You mean nitrogen (N) and phosphorous (P) don't have a covalency of three, and oxygen (O) and sulfur (S) don't have a covalency of two?

Yes, and no.

Arrrgh! Out with it wizard!!

OK. OK! It's a minor point really, but all four of those elements, under certain conditions, can have "odd" bonding. Under certain conditions, they seem to make 3, 5 or even 6 bonds!

What? When? How?

Now that would be going into deep water. You might sink. Instead I just want you to know that sometimes these four elements can make more bonds than you might expect. It has to do with d and f orbitals and a quality called "hypervalency". Let's not get into it, OK?

OK.
I see that Group VI is the home of oxygen. That's an important element, isn't it?

Indeed it is! Oxygen is the most abundant element on Earth. (Aluminum is the most abundant metal element.) Oxygen makes up 20% of the air (as elemental molecules of O2), 89% of the oceans (as the compound molecule H2O) and 49% of the crust of the Earth (as many different compound oxides).

Wow. Earth is mostly oxygen!

Yes, it is. No other planet in this solar system has as much oxygen in its atmosphere. It occurs as two types. The one you are most familiar with is O2, but it also forms molecules of O3, called ozone.

Are they allotropes?

What do you think?

Yeah, I think they are allotropes. They have to be! They are pure, elemental oxygen but arranged differently.

Right. You see now that allotropes can be different in different ways. We can use that word to describe the difference in size between atoms, like we did when we talked about atomic size. We can use that word to describe differences in their (solid) structures, like the red and white phosphorus. We can even use that word to describe differences in the elemental molecules they make, like O2 and O3.

I bet O2 and O3 have very different properties.

Yes. Diatomic oxygen (O2) dominates the atmosphere and is needed by all living things (ignoring some strange microbes).

Diatomic oxygen (O2) -you can't live without it!

Right. Ozone (O3) is also a gas in the air, but it is in very low concentrations.

Is ozone good for you or bad for you?

That depends on where it is. You don't want to breathe ozone. It is a very reactive allotrope of oxygen. Even more so than O2 and that's still very reactive.

Yeah, O2 reacts with aluminum and phosphorous. Does ozone?

Oh, yes! Ozone also reacts with your lungs, eyes and skin. Ozone can destroy them!

So we need to breathe O2 but we shouldn't breathe O3.

Right. However, the Earth's atmosphere has a layer of ozone, high above our heads, which protects the Earth from harmful sunlight.

Harmful sunlight? Like what?

One form of sunlight, called ultraviolet light, can damage living things. It hits the skin and causes nasty chemical reactions. It can even cause cancer.

So we want O2 down here where we breathe it to keep us alive. But we want O3 up above to shield us from those dangerous rays.

That's right. In the 20th century our species nearly destroyed life on Earth by messing around with ozone. Man's misuse and misunderstanding of atmospheric Alchemy caused him to make some stupid mistakes. Many of their machines created ozone at ground level.

But they would breathe it! Why did they do that?

They didn't do it on purpose. It was a side effect of chemical reactions they used to run their machines. Some of their machines produced a nasty mixture of chemicals called "smog". The smog turned O2 into O3. That caused damage to people's lungs.

What about the ozone we need high above? Did they mess with that too?

Yes, they did. Again, they didn't know what they were doing. Some of their machines and industries released chemicals into the ozone layer that broke up the ozone.

So the ozone (O3), way up there, decreased in the 20th century?

Aye. As the ozone shielding disappeared, more harmful ultraviolet sunlight got through. That increase in ultraviolet threatened to destroy any living thing that stood in the sunlight.

Including plants?

Aye, including plants. The decreasing ozone layer threatened to destroy the Earth's entire ecosystem by killing all the plants! If ozone had continued to decrease it would have been the end of life on Earth (as we know it)!

But it didn't. Right?

Right. By the end of the 20th century Man understood what he was doing and banned (most of) the chemicals that harmed the ozone layer. It was a close call.

It's a good thing they learned Alchemy in time!

Yes indeed. The full story is a course in its own. However, we have other elements to talk about.
Like sulfur. Sulfur is one of the least reactive nonmetals.

Kind of like nitrogen.
There seems to be a "flip-flop" among these four elements at the top of Groups V and VI.

How do you mean?

Nitrogen is very inert, but the element right below it, phosphorous, is very reactive. At least as white sulfur. But in Group VI the reactive element is oxygen, at the top, and sulfur is the nonreactive element. Reactivity "flip-flops" among these four elements.

Why, yes it does! Good observation, Arthur. I told you strange things occur among these four elements!
sulfur can be found in huge amounts as raw sulfur.

Does sulfur have allotropes?

Yes, it has several allotropes, but I think we will skip them for now.

OK. Can we move on to Group VII now?

Yes, we can. The Group VII elements are also called the halogens

Why?

Because they form salts. "Halogen" is from the Greek meaning "salt-producer". The most common salt on Earth is sodium chloride (NaCl). Geologists call rocks of NaCl, "halite".

The halogens sound important.

They are. Except astatine (At). It's too rare. But all the others are fairly common, and they can all form covalently bonded, diatomic, elemental molecules.

Which means they make F2, Cl2, Br2 and I2 by sharing electrons in their outer shell.

Right. Like O2 except they make only one covalent bond between them....

.... because they need to share only one electron to complete their outer shell.

Right! I'm glad you've been paying attention.

But I thought they made ionic compounds like HF and NaCl!

That's also right. These covalently bonded, diatomic, elemental molecules of Group VII are VERY reactive. They have very high electronegativities. (Just look at the table!) That means they will readily grab another electron in order to complete their outer shell and become an anion to do it.

Wait! Wait! I'm lost.

I'm not surprised. The point to understand is that the Group VII elements can form covalent elemental molecules. But these elemental molecules are so reactive that (given the chance) they will quickly react with what ever they can in order to become an ionic compound.
Do you see what I mean?

Maybe. When Group VII elements are all alone, nothing but fluorine for example, then the only thing they can do is form covalent bonds by sharing their outer electron.

Right. Fluorine's high electronegativity means it can pull electrons away from any other atom EXCEPT another fluorine! Two fluorines have the exact same electronegativity, of course. One fluorine can't take an electron from the other.

I see. Because they are equally matched in electronegativity, neither one can win the tug-of-war, so both share an outer electron pair. Thus, fluorine all by itself forms F2 by covalent bonds.

Right. But when you allow F2 to touch another atom of a different element, the fluorines get their chance to grab electrons. Their electronegativity becomes important when they have a weaker atom from which to steal an electron!

I get it. F2 is stable only as long as no other kinds of atoms come in contact with it. But when a different element comes along it switches from covalent to ionic behaviour.

Right. If you add F2 to water the fluorines will rapidly steal an electron from the hydrogen to form the ionic compound, HF.

I see. What about Cl2?

Same thing. If you add it to water you get HCl because the chlorine takes the electron. If Cl2 were to come into contact with sodium (Na)...

... it would grab an electron from the sodium and convert itself from covalently linked elemental molecules of Cl2 to ionic compound molecules of NaCl.

By Jove I think you got it!
The electronegativities make Group VII elements very reactive because they form ionic compounds whenever they can. But the only time when they can is when there is another weaker atom around.

I understand. Does this reactivity decrease down the Group?

What do you think?

I think reactivity would decrease because electronegativity decreases down any Group so fluorine is the most electronegative, then chlorine, then bromine..

Right. Electronegativity decreases down the Group and because it is the electronegativity that causes these ionic compounds to form. Reactivity will decrease as you move down the Group.

Just like the metals!

Ah, yes and no.

What do you mean?

The electronegativity decreases down both Group I and Group VII but the effects on reactivity are the opposite. Recall the reactivity of the Group I elements. Think what happens as you add each to water. Remember?

Oh, right!
The Group I elements increase in reactivity as you go down the Group, but the Group VII elements decrease in reactivity.
Hey, the electronegativity of both Groups decrease down the Groups! What gives!?

Think about what the electronegativity means to elements in those two Groups. The decreasing electronegativity of Group I means they are MORE likely to give up an electron to become a cation. The decreasing electronegativity of the Group VII means they are LESS likely to become anions.

I see.
The effects are opposite because they are creating opposite ions.

Right! Take a good look at the Table and make sure you know what that means. You can always refer back to the Table of electronegativities in order to predict the relative behaviour of the elements and the bonds they make with each other.

Great. Can we move on to Group VIII elements now?

Aye. The Group VIII elements are the noble elements.

So they don't form molecules at all. They are always monatomic.

Yes, and they are all gases, so we often call Group VIII the "noble gases".

What else can you teach me about the Group VIII elements?

Nothing else. They are so inert that the only thing to know about their chemistry is that they are inert!

And that's it! We finished the Table! We're done.

No. We haven't talked about the other part of the Table where the transition elements hide.

Oh, I almost forgot they were there!

Easy to forget them. I haven't been focusing on them because they are so unusual. But I would be a poor teacher if I didn't teach you about them.

Why have you hidden them? What is so weird about them?

The huge shells in this part of the Table play complex tricks with their orbitals and their second outer-most shell.

What's a "second outermost shell"?

That's the shell immediately below the outermost shell. In this part of the Table some complex interactions occur between shells and orbitals.

Does that mean you aren't really going to explain it? Just tell me it's weird?

Oh, I'm not going to let you off that easy!
We are now entering the most complex area of the Periodic Table, the home of the transition elements. These atoms are not easy but you can understand them if you take it slow and think about it. So, here goes.
Large atoms have large shells. When shells become big two important things can happen.
  1. The energy levels between the two outermost shells are not very different so assigning electrons becomes more complicated. We have to think about two shells at once. (Sometimes three shells at once!)
  2. These atoms can start to use their d and f orbitals.
Do you understand what I mean by the outer shells getting larger?

Yeah, they are larger than M-shells.

Yes, they are N, O, P and Q-shells. Atoms with these large shells juggle their electrons in a weird way.

In the middle of the Table the atom can put electrons into the d and f orbitals of its second outermost shell. They kind of back up and use d and f orbitals that they skipped.

But you taught me that doesn't happen!

It doesn't happen for the typical elements, but the transition elements are...

Not typical! Why not?

The transition elements can fill the d orbitals of the shell that is just below their surface. These d orbitals are hidden below the outer shell of an s orbital.

More complications! I don't now if I can take it!

Just remember that the elements in the fourth Period can place electrons into the d orbitals of their M-shell but first they fill their s orbital in the N-shell. The first two fourth Period elements, potassium (K) and calcium (Ca) are Group I and Group II so they fill their s orbital first and there is nothing special about them. After all, they are typical elements. But scandium (Sc) puts its third electron in a d orbital in its M-shell!

But that's down a shell!

Right, and that's what makes the transition elements so difficult to understand.
The transition elements continue to put their electrons into the d orbitals of their second outermost shell until those 5 d orbitals are full. Zinc (Zn) has all 5 of its d orbitals in its M-shell filled so the next element, gallium (Ga) puts its next electron into a p orbital of its N-shell.

So, as you move across the Period of the transition elements the electrons go into d orbitals in the shell below! That is weird.

Yes, it is. Notice that once all the d orbitals of that inner shell are full the next element will behave correctly by placing its next electron in a p orbital of its outer shell. So that element is a typical element.

Yeah, I see what you mean. Gallium (Ga) is a typical element. Once the d orbitals are filled the elements become typical again.

That's right. Gallium is a typical Group III element and its outer shell is like those of all the Group III elements. Its outer shell (its N-shell) has two electrons in its s orbital (just like calcium) and one electron in its p orbital. Gallium also has a full set of 10 electrons in the d orbitals of its M-shell (like zinc) but those are now buried so deeply that they do not affect the properties of gallium (much). So gallium behaves like a typical Group III element.

I bet similar things happen with the transition elements in the fifth, sixth and seventh Periods.

You bet right.
As you move across the Periods of these larger atoms you find there is a transition after Group II because these atoms start to put the electrons into the d orbitals of the previous shell. Once those 5 d orbitals are full (it takes 10 electrons to do that) the next electrons go into the p shells of the truly outer shell. It is at that point that the properties of the elements again become typical.

I think I get it. You really have to look at this part of the Table to see what you mean.
So what can you tell me about the transition elements?

Because the transition elements play tricks with their outer shells, the electrons in their SECOND outermost shell can take part in bonding. That means their chemical behavior is influenced by the second outermost shell as well as the outermost one.

That's weird!

Yes. A few of the transition metals do strange things. That's because their lower d orbitals are so close in energy to the s orbital in the outer shell.

I guess the electrons get confused. And so do I!

Let's try to remember that MOST of the things you have learned apply to all elements but some of these big atoms surprise Alchemists. Do you see what I mean?

I think so. We usually think of the properties being caused by the outer shell only, but in the transition elements the next inner shell (the second outermost shell) also affects them.

Right!
Tell me, Arthur, how would this strange shell behaviour affect their conductivity? Take a guess.

Well, conductivity has to do with the ability to move electrons in a current. I would guess that transition elements would be good conductors because they are so sloppy with their outer electrons. My guess is that the transition elements are good conductors of electricity.

Correct. A very educated guess. What do we call an element that is a good electrical conductor?

A metal.

Correct again! All the transition elements are metals so they are sometimes called the transition metals.

So they are just like the alkali and alkaline metals?

Well, only in their ability to conduct electrons. It just so happens that the transition metals have the properties most people think about when they think of a metal.

What do you mean?

Up until now all the metals we have been talking about, all the metals in the typical elements, were soft.

So the metals in the typical metals are not typical metals? I'm lost!

Now, don't get yourself confused here.
All the typical elements are called "typical" because they fill their shell in a typical manner. That includes the metals in Group I, Group II and even the metals in the other Groups.

Like aluminum (Al).

Right. All the typical metals, from lithium (Li) to polonium (Po) are "unusual" among metals in that they are fairly soft. You can cut them with a knife. And they are also "unusual" because they have fairly low melting temperatures.

But they still conduct electricity so they are metals.

Right. All the transition elements are metals too, because they also conduct electricity. But the transition metals are mostly hard and have high melting points, not at all like the metals among the typical elements. The transition metals, although not typical elements (because of their shell filling rules) have the properties that most people think of as being "typical" of a metal.

Like being hard and having a high melting point.

Right.

Hey, what about mercury (Hg, atomic number 80)?!
It's a liquid! It isn't as strong as iron. It isn't strong at all!

Mercury is an exception among the transition metals. It is the only one that is a liquid at room temperature. Perhaps it has to do with the fact that it is the last transition metal of that Period. Strange properties might arise because of mercury's well-filled shells.

Hmm, maybe. I see what you mean. The next element, thallium (Tl), is a Group III element so maybe mercury is extra weird.

Maybe. Regardless, mercury is a liquid but it is still a metal because it conducts electricity very well. If you were to cool mercury it would form a hard solid like iron, not a soft solid like sodium.

Strange. I guess that's the transition elements for you. They're strange. Weird.

Yes. The weird outer shells of the transition elements cause them to have strange properties. Some of them appear to have unusually small atomic radii because of the way they can shuffle the electrons in their outer shells. That is especially true when they form compounds or ions.

I see. I guess they form metal bonds.

Yes, but some of them can also form covalent bonds.

What? You mean they share their electrons?

That's right. Again, it all has to do with the weird way they fill their shells. Some of them can use their weird shells to share their outermost and second outermost electrons.

Weird.

Aye, it is weird, and it gets weirder. Some transition metals will share their electrons with ions! That produces a bond that is a covalent link between a metal and an ion.

That's very weird! Kind of hard to imagine.

Yes. It sure is. It also makes for some very complex bonding behavior.
Also, some of the transition metals can have a variety of stable ions.

What do you mean?

Well, take iron (Fe). It can lose two electrons to form a cation with a charge of +2 (Fe+2). But, under different conditions, it will give up another electron to form Fe+3.

So transition metals can have different ions.

Right. Some of them can behave as if they were two different elements!

This all has to do with their weird shell filling behavior?

Aye. Now, don't get the wrong impression about the transition elements. Each element has its own special properties. Even the transition elements follow certain rules about their shells, bonds and ions they can form.

So it isn't true that "anything goes" with a transition element.

Right. Each element has its own rules, its own properties. Take copper (Cu). It can form at least two types of cations and make some of those funny covalent bonds. But a different transition metal, titanium (Ti) for example, will do something different.

I see what you mean. Each element, transition or otherwise, follows certain rules.

Right. In the case of transition elements these rules get very complex and there appears to be no order or pattern to them.

But if I knew all the details about their strange shells, I would understand all their weird chemical properties.

Probably.
Let's do a little tour among the transition elements.
Iron (Fe), copper (Cu), silver (Ag) and gold (Au) have been important elements in creating our civilization. You've probably heard of them before.

Oh, yeah. Especially silver and gold!

Yes, those two elements get a lot of attention because they are so rare. But iron and copper are more useful because they are more abundant. And they have better properties for making things.

Like swords.

Yes.
The transition elements (transition metals) start in the fourth Period with scandium (Sc). The last fourth Period transition element is zinc (Zn). The fifth Period starts with the two typical elements of rubidium (Rb) and strontium (Sr) but the transition elements for that Period run from yttrium (Y) to cadmium (Cd). The sixth Period transition elements start at lanthanum (La) and go to mercury (Hg). The seventh Period has only one transition element, actinium (Ac).

Hey, there's a gap in the transition metals! Just like the gap I found in the typical metals. Look. Lawrencium (La) has 57 protons but the next element, hafnium (Hf), has 72!

Yes, very good observation. What do you think that means?

That you've hidden another Table, and my guess is that it contains even weirder elements!

Right on both points. In that gap are the "inner" transition elements. Cerium (Ce) is the first inner transition element and the first element to use an f orbital. It places an electron in the f orbital of its N-shell.

Its N-shell! But cerium is in the sixth Period. Its outer shell should be its P-shell and its next outermost shell is its M-shell. What gives!

Well, cerium does have an outer shell made of its P-shell. It has two electrons in the s orbital of its P-shell, just like the outer shell of the sixth Period Group II element barium (Ba). And it also has an electron in the d orbital of its next outermost shell, its M-shell, just like lanthanum (La). But the last electron of cerium gets buried still deeper in an f orbital in its N-shell.

I see. (I think.) The inner transition elements place electrons into the f orbitals of their second outermost shell. So they are like the transition elements in that they put electrons into lower shells. They just bury them deeper.

Right. The inner transition elements place their new electrons in f orbitals a shell lower than the shell into which they place their d orbital electrons.

This is awfully confusing.

I agree. And it isn't always as simple as this might suggest.

Simple?! You call this simple?

OK, it isn't simple. What I mean to say is these orbital patterns get very complex. For example, praseodymium (Pr), the element that follows cerium, places its next electron into an f orbital in the N-shell.

Yeah, I'd expect that.

Me too. But we would not expect praseodymium to also move the one electron from its d orbital M-shell to the N-shell as another f orbital electron.

What!? This is ridiculous.

I agree.
Actually all the rest of the inner transition elements set up their shells like praseodymium. They leave their N-shell empty and fill up their M-shells with f orbital electrons. The last inner transition element of Period six, lutetium (Lu), then places its next electron back into the d orbital of the N-shell so it looks more like cerium again but with all its f orbitals in the M-shell full.

This makes no sense at all!

I agree. I don't expect you to memorise this material. But I think you should know that the weird things that happen among the transition and inner transition elements become even weirder at the parts of the Table where you move from one type of electronic structure to another.

OK. What should I know besides that?

Lanthanum (La) is the transition element of the sixth Period that "introduces" the inner transition elements of that Period.

Indeed, the inner transition elements from cerium (Ce) to lutetium (Lu) are sometimes called the lanthanoid series. (Some folks call it the lanthanide series.) Lanthanum introduces the lanthanoid series.

But lanthanum itself is not an inner transition element. It's an ordinary transition element.

Right.
There is a second Period of inner transition elements starting with thorium (Th). That series is sometimes called the actinoid series (some folks call it the actinide series) because actinium (Ac) introduces it. Actinium is a transition element that introduces another inner transition series. These two series fill deeply buried f orbitals with electrons.

Phew. I'll write that down as a definition.

Good. You may have noticed the actinoid series doesn't have an end.

What do you mean by that?

The largest element found in nature is uranium (U). It has 92 protons and one of its isotopes is very unstable.

You mean it is radioactive. It decays into another element. Right?

Right. In the 20th century Man learned to slam atoms together and make elements even heavier than uranium. They are unstable and of little use to Alchemists.

So Man can make new elements?

Yes, in the 20th century. That bit of Alchemy is in the realm of nuclear physics and we won't go into it. What I want you to know is that all those large, artificial elements are very unstable. Because they follow uranium, they are often called the transuranic elements.

But we won't discuss them because they are of no use to Alchemists. Too unstable.

Right. Uranium is the largest natural element. It has 92 protons (and 92 electrons, unless it's an ion).

Are any of the inner transition elements useful?

Uranium (U) is. One of its isotopes (235U) is very radioactive and can be used to make nuclear power. But that's in the 20th century.

Is it hard to make nuclear power? Is it worth it?

That is a very good question. Uranium power can be harnessed to make large amounts of electricity. (The exact way it does that is best left to a class on nuclear physics.) However, in the process of making that power, lots of radioactive atoms are made. One of them is plutonium (Pu). It has 94 protons.

So plutonium is a transuranic element and very unstable.

Right. Plutonium is a VERY dangerous element. It, along with other radioactive atoms produced by nuclear power, must be stored or disposed of safely. That's not easy to do.

So, NOW we are done with the Periodic Table. Right?

Right. Any questions?

Yeah. Why did you show me the Table in bits and pieces? Why not show it to me as one complete Table?

OK. Here is the complete Periodic Table of the elements.

Wow. The Table is very long.

Yes, it is. All those transition and inner transition elements take up a lot of room. In order for all of them to be seen you have to shrink the Table so small you can't read it (very well). Notice that the inner transition elements start immediately after one transition element (either La or Ac) and that causes the Table to get spread out very wide. Remember, those elements immediately after Group II are transition elements but it's hard to label them due to the split caused by the inner transition elements.
Most of the time we work only with the typical elements so we can use an "abbreviated" Table to see all those elements at once. Some Alchemists will include a Table that shows the transition elements as well. But that is when things start to get too big and complicated to see very well.

So Alchemists just focus on the parts of the Table they are interested in and only skip to other Tables to see other elements when they must.

Right. Now we have finished the Table.
Any questions or comments?

Yeah. I think the shape of the Table has to do with orbitals. Is that right?

Yes indeed. Now tell me what you mean by the orbitals affecting the shape of the Periodic Table.

Look at the Table again. It looks like big gaps and changes occur when the new orbitals are added.

That is an excellent observation! You're right! Try to explain it.

The first two Groups (Groups I and II) have s orbitals as their outer orbital. Right?

Right.

All the atoms with p orbitals as their outer orbitals are on the far right of the Table in Groups III to VIII.

Yes, that is true. Good observation.

I remember that some weirdness occurred when we went from Group II to Group III elements. There was a hiccup in the first ionization energies. Remember?

Yes, I do. Congratulations, Arthur. You have discovered a wee bit about the Table I was not going to cover. You are discovering how the orbitals affect properties. That "hiccup" in the ionization energies as you went from the Group II to Group III elements is due to the addition of a new type of orbital. The p orbital. Can you guess why?

Hmmmm...
It should be harder to ionize a Group III element than a Group II element, because the electron in the outer shell should be attracted closer to the nucleus of a Group III element and held tighter to it. But it looks like that isn't true on this ionization Table. At least for the first three pairs (Be-B, Mg-Al, and Ca-Ga). It's as if the first electron to be put into a p orbital is held loosely.

You are absolutely right! The lone electron in the first p orbital is "looser" than it should be. That has to do with the quantum mechanical world. Further down the Table those weird properties are hidden and that lone electron (in the p orbital) behaves correctly.
Is there anything else you've noticed now that you've seen the whole Table laid out? What about the other dips in the Table.

Those other dips are where all the "real" metals show up. They all conduct electricity and "super share" electrons. That has to do with those other orbitals we never drew.

Yes, it does. The d orbitals are introduced starting with the fourth Period, but the M-shell is where the d orbital electrons go first.

Notice that all the fourth Period elements have an N-shell as their outer shell, but the M-shell is used by the transition elements of the fourth Period as a place to store their electrons (in d orbitals).

The fourth Period is where the transition metals start with scandium (Sc). Scandium is the first atom to have a d orbital. It's the first transition metal. All the transition elements start right after Group II and end right before Group III.

Right! All the transition metals have d orbitals as their second outermost orbitals. (They like to fill their s orbital in their normal shell first.) That's what makes them complex.

D orbitals have funny, complex shapes and there are 5 different types. Right?

Right. Tell me, how many electrons can fit into all those d orbitals?

Well, Pauli's exclusion principle always applies and he say that you can have no more than two electrons per orbital, so I guess these new orbitals can hold up to 10 new electrons. That makes sense because there are exactly ten elements across the transition elements. Scandium (Sc) has atomic number 21 and zinc (Zn) has atomic number 30.

Very good observation. Now you know how they fit into the Table. After zinc (Zn) comes gallium (Ga) and that's where the first p orbital of the N-shell begins. Gallium is a Group III element because its outer shell (N-shell) is like that of any Group III element. It's made of a complete s orbital holding two electrons and a single electron in a p orbital.

The remaining Periods work in a similar way. Yttrium (Y) starts the transition elements in the fifth Period and cadmium (Cd) finishes that Period of transition elements. Lanthanum (La) starts the transition elements in Period six and the last transition element of Period six is mercury (Hg). Actinium (Ac) is the only transition element in the seventh Period.

Very good. What about the inner transition elements?

The inner transition elements start in the sixth Period. The lanthanoid series starts with cerium (Ce), the first element to use f orbitals. It ends with lutetium (Lu). The actinoid series is in the seventh Period and starts with thorium (Th).
There are 7 different types of the f orbitals and they can hold two electrons each so there is room for 14 electrons. That means you can fit in 14 elements as inner transition elements using the f orbitals to hold the electrons.

Yes.
We would have to get into some complex physics to understand all of the Table and learn all the exceptions. But you know that the Table can be a useful tool to understand atomic behavior.

We've covered a lot of information here. I think these last Tables prove it! Read through it all again and check your notes. Take a good look at the Tables.

I'm going to draw my own Periodic Table.

That's a great idea because by drawing your own Table you will learn a lot of its details. As you draw your Table, keep track of the Periods and Groups. Keep track of where each new type of element starts. And draw on your Table what the outer shell of each Group should look like. You can draw a Lewis structure for each Group. Every Alchemy student makes his or her own Table and has learned a great deal as part of that project. (You'll need a big piece of paper, ruler, pencil and eraser.)
Try the questions I'll give you later.
(Don't worry. I won't give you a difficult time with transition or inner transition metals.)

(Thanks!)

This work was created by Dr Jamie Love and licensed under a Creative Commons Attribution-ShareAlike 4.0 International License Creative Commons Licence.