Cycloalkanes

I just found an error in my textbook. Seriously. The book even bolds its own error. The offending sentence reads...

Cycloalkanes have molecular formula C_nH_2n and contain carbon atoms arranged in a ring.

That is only true for cycloalkanes with just one ring. Cycloalkanes can have more than one ring, and each additional ring means two fewer hydrogens. And there are a lot of those. Try to keep up, textbook. Anyway, the examples the book then uses for cycloalkanes are all ones with just one ring. In fact, the examples in this section of the book have all of the carbons in the ring, but this is not necessary or even particularly significant.

The smallest cycloalkane ever is cyclopropane, with a molecular formula of C₃H₈ and a skeletal structure that looks like this...
Yes, it's a triangle. This really should not surprise you if you've been paying attention, which you haven't. I could show it in 3-D, but so far I haven't figured out a way to post my wonderful 3-D images here in this blog thing without them looking like crap (because I am pasting them into MS Paint and saving them there. If you want pretty pictures, go read a pretty pictures blog or something. I hear that they have those. Such things may be more suited to your intellect.

Cyclobutane's skeletal structure looks like a square. This should be easy to visualize. Same with cyclopentane and a pentagon. I have already shown the skeletal structure for cyclohexane here and here.

Also, the book claims that the largest known cycloalkane with a single ring has 288 carbon atoms. But this is in a problem asking for molecular formula (and obviously the molecular formula is C₂₈₈H₅₇₆) and I cannot tell if it is giving me authentic trivia or merely posing a hypothetical for the purposes of asking such a question and reinforcing the concept.

One last thing, which the book apparently omits in this section (although it will probably come up later) is ring strain. The carbon atoms are most stable at a certain bond angles. In the case of alkanes (and lots of other things, really), the ideal bond angle is 109.5° and all four groups attached to the carbon atom are equally far away from each other, forming a tetrahedron with the carbon atom in the center and each attached group in one of the corners. But when carbon atoms form rings, the bond angles become strained. This ring strain causes the molecule to be more reactive. Cyclopropane, with 60° angles between the carbons, has the most ring strain. After that, it becomes important to note that these rings are three-dimensional objects. They can be denoted with two-dimensional skeletal structures on paper, but are under no obligation to lie flat. So cyclobutane does not actually have 90° angles between its carbons, as "puckering" reduces strain and creates larger angles. Later in the chapter, this is explored for cyclohexane in particular, which has the most stable ring among cycloalkanes.

A Note on Complexity and Isomerism

My textbook has a table with information that I did not include in my last post, but that may improve understanding of isomerism. In case it is not obvious, the number of isomers grows with the size of a molecule. In my last post, I showed the two isomers of butane. Larger alkanes have even more, because with more atoms, there are more ways to rearrange them. Small alkanes are easy to understand in this regard. A hydrocarbon with one carbon has no isomerism. The same is true for two or three carbons. When we get to four, as already demonstrated, there are two possibilities: a straight chain and one with a branch. Five carbons means three isomers. With seven carbons, we get nine isomers, which is still manageable, but then add a single carbon and there are eighteen isomers. The table ends with icosane (C₂₀H₄₂), which has 366,319 constitutional isomers.

And that is just acyclic alkanes. There are so many other things to consider, that the complexity is staggering. And that is why we have a systematic method of naming molecules. Anything else would get pretty impractical.

Constitutional Isomers Redux

I suppose that my textbook introduces constitutional isomers in the alkanes chapter because alkanes are pretty straightforward and can ease one into the concept. Constitutional isomers can and do occur in other molecules. Isomerism is when two or more different compounds have the same molecular formulae. In other words, they have the same kinds of atoms and the same numbers of those atoms, but something makes them chemically distinct. Later on, we will explore stereoisomers, and it will be very exciting. But for now, we're looking at constitutional isomers, which differ in the way the atoms are connected to each other. Let's take a look at two molecules that are constitutional isomers of each other...

Name: n-butane (or just butane)
Molecular formula: C₄H₁₀
Condensed structure: CH₃CH₂CH₂CH₃
In stunning 3-D:
Yes, I just figured out that I could render butane three-dimensionally with my nifty software. Anyway...

Name: isobutane (or 2-methylpropane)
Molecular formula: C₄H₁₀
Condensed structure: CH(CH₃)₃
In glorious 3-D:
Both molecules have the same quantities of the same atoms. But the bonds are not identical here. A carbon bonded to two other carbons and two hydrogens is electromagnetically different from one bonded to three other carbons and one hydrogen. Also, the three-dimensional forms are quite different, and when the molecules interact with other bodies (including other molecules just like themselves) the results will be at least slightly different. Although very similar, these two compounds have different chemical and physical properties. They are more like each other than other compounds that have different atoms and other, more striking differences. Because of these facts, we use the term "constitutional isomers" to denote the relationship between these similar molecules.

But when it comes to properties, constitutional isomers are not always so similar to each other as those two. Some constitutional isomers contain different functional groups from each other and, if you remember the importance of functional groups like you should, this means they can have dramatically different chemical and physical properties...

Name: ethanol

Molecular formula: C₂H₆O

Condensed structure: CH₃CH₂OH

In brilliant 3-D:

It's an old friend: ethanol. I don't know how many times I've shown ethanol before, but you had better know that this is what it looks like. And if you managed to actually have some brain capacity, maybe you even remember that this compound is an alcohol, as it has a hydroxyl functional group. Easy, but here's a constitutional isomer of ethanol.

Name: dimethyl ether (or methoxymethane)

Molecular formula: C₂H₆O

Condensed structure: CH₃OCH₃

In spectacular 3-D:
Since the name has "ether" in it, you have deduced, unless you are a total idiot, that this is an ether (the name of the functional group is methoxy in this case). But the molecular formula is the same. The functional groups here are so unlike each other that reactions possible for one would be impossible for the other. Oh, and remember hydrogen bonding? Ethanol has it. Dimethyl ether cannot have hydrogen bonding because there is no hydrogen attached to the oxygen, so these two even have different intermolecular forces. In this way, two constitutional isomers can be quite dissimilar. What kind of atoms a molecule has and how many are very important, but the configuration of the bonds holding the atoms together in a molecule matters a lot too.

Edit: After posting this, I started going back to tag my posts. I noticed that way back in February, I wrote a post about constitutional isomers. I think this new post is better, but here is the old one. If you do not get the concept after reading this post, read the old one. If you still don't get it, tell me, I guess. It seems fairly simple to me and I think I did an adequate job of explaining it both times, but maybe I am wrong...

Cyclic and Acyclic Alkanes

As I mentioned in my Functional Groups post, alkanes are hydrocarbon molecules with no π-bonds. They can be straight chains of carbons with attached hydrogens, or there can be branches or rings or both. All of the fourth chapter in my textbook is dedicated to alkanes. But the first part is just about getting acquainted with them. Alkanes are something of a baseline in organic chemistry. It's when functional groups are added that the chemical properties behind so much of our world come into play. Lacking functional groups, alkanes are not particularly reactive. They can react, though. And I know we'll come to that eventualy. There's a lot to learn from alkanes, though.

Firstly, let's distinguish between acyclic alkanes and cyclic alkanes. If it has a ring, it's cyclic. If it does not have a ring, it is acyclic. Simple, right? It better be. No, two rings is still cyclic. What counts as a ring? Oh, good question. A ring is pretty much what it sounds like. Three or more atoms bonded to each other with a loop that can be formed from the bonds between them. Carbon #1 is attached to Carbon #2 and Carbon #2 is attached to Carbon #3, which is itself attached to Carbon #1. Three atoms is the minimum, but larger rings are more common.

For an acyclic alkane, the number of hydrogens will always be two plus double the number of carbons. H = 2C+2. Actually, a little logic should demonstrate this point. No amount of branching chains changes the formula. But a single ring does. I shall illustrate with some examples. First, here is hexane...

Name: n-hexane
Molecular formula: C₆H₁₄
Skeletal structure:
Well, that's a nice, simple acyclic one. How about an acyclic alkane?

Name: Cyclohexane
Molecular formula: C₆H₁₂
Skeletal structure:
I Know I've shown this one at least once here, once upon a time. Hexagons should hopefully be pretty recognizable. And notice that it has two fewer hydrogens than the last one? That's because of the ring. What? You want to know how the ring makes it so that there are two fewer hydrogens in the molecule? Really? Look, just pretend we sever the bond between two carbons. Any two. Now those two carbons need a new bond to something else because, remember, carbon forms four bonds. So we stick a hydrogen onto each of them, and look at that, it's n-hexane, the same molecule I already showed you just before this one. Amazing. And that is why the ring makes it so that there are two fewer hydrogens than in an acyclic alkane. Simple.

Acid Strength

Acid strength refers to the potential of the acid to lose a proton. Acids that give up a proton easily are stronger. Acids that don't readily give up a proton are weaker. Here, I'll show you a reaction mechanism for an acid dissolving in water. I know you don't know what those are, but if you paid attention the first time I showed you, then you would. It's not my fault you weren't paying attention. Grow up. So yeah, mechanism...
For a little review, I'll note that the first reactant (A—H) is the acid. Water, the second reactant, is the base. And the products are the conjugate base of the reactant acid and hydronium, the conjugate acid of water, respectively. This reaction is reversible, so the hydronium could protonate the conjugate base of the original acid and leave us with the reactants again. But the equilibrium favors whichever side of the equation has the weaker acid. This might seem intuitive, actually. The acid that more readily gives up a proton (the stronger acid) will do so more and therefore will show up less. But unless we know which acid is weaker, that is, unless we have a way to measure acid strength, the knowledge doesn't really help us. Fortunately, quantifying the acid strength is possible. But this goes a bit beyond the scope of the textbook I'm using, which assumes the student remembers certain things from general chemistry. Besides, it involves math. So I'll just tell you that in organic chemistry, the figure that is typically used is pK_a. That's the opposite of the common logarithm of the acid dissociation constant. The acid dissociation constant is determined by multiplying the equilibrium concentrations of the products and dividing this by the equilibrium concentration of the acid. Easy, right?

So I totally don't remember how those equilibrium concentrations are determined (with instruments, I guess). I just look up the pK_a of the acid in question using a table of pKa values. That way, all I need to know is that the lower the pK_a, the stronger the acid. My textbook states that typical pK_a values for organic acids range from 5 to 50. I'm not sure where those numbers came from, but whatever. Be aware that some organic acids are not typical. Also, in lab, I dealt with inorganic acids all the time. Acids with negative pK_a values are considered "strong" acids. Actually, Wikipedia says that strong acids are those with pK_a values less than -2. Whatever. What it means for an acid to be "strong" is that essentially all of it will lose its protons to water. In other words, the equilibrium completely favors the products and none of the original acid is present in any concentration. So really, the strongest acid that exists in water is hydronium. Anything stronger just protonates the water to form hydronium.

Acid-Base Definitions

Do you know what acids and bases are? Seriously? I know I used to think I did and I totally didn't. I mean, I knew some examples of acids and bases, but the chemistry behind them was completely unknown to me even after I took chemistry in high school and even when I started taking it in college. But finally, in one class I was introduced to three acid-base definitions. I knew there were more, but I had no idea how many. This post will introduce some of them. Really, I've only ever used two of them and those two, which happen to coincide with each other a lot, are the only two that will be important here, but I think the historical definitions are interesting and this is my blog or whatever and so I get to make a post including them.

Lavoisier definition
In case you didn't know who Lavoisier was, he was one of the founders of chemistry and you are not worthy. Through a series of experiments, he determined that oxygen combining with other elements could lead to "acidic" (the word comes from the Greek word for "sharp") properties. He extrapolated from this that oxygen was the element that contributed the acidity. This is where oxygen got its name, which means "acid-generating." There was a small problem with this: Lavoisier was wrong. Considering that he basically invented chemistry, I think he's allowed to be wrong every once in a while.

Liebig definition
You know who Liebig was too, right? Because he was another great chemist. Anyway, this was sort of the first real definition. A Liebig acid is a molecule that contains at least one hydrogen atom that can be replaced by a metal. This was only a definition for acids. Back then, bases were loosely defined as the opposites of acids. Known acids were generally liquids, so a base was whatever compound would react with an acid to neutralize it into a solid salt. Such reactions are known as acid-base reactions and we'll deal with them other posts pretty soon.

Arrhenius definition
This was one of the three definitions I originally learned and it was the main definition everyone used for a long time. But that was in the past. The distant past. Like before I was born, even. Probably before you were born too. Arrhenius acids are molecules that, in water, lose hydrogen atoms, generating hydrogen ions in the solution. Well, we now know that they're actually hydronium ions. Oh, hydronium is important. You should know what it looks like. Here, I'll paint one for you.
If you weren't stupid, you'd remember what a water molecule looks like. Man, I'm not posting water again here just for you. Go back and find the post where I did show it or something. Anyway, this is like water, but with another hydrogen. Oxygen normally only forms two bonds. I guess I never talked about formal charges or whatever, but the oxygen is positively charged now. Really. We'll talk about it later if you want. Hydronium is properly written as H₃O⁺. But you should be aware that a common shorthand is just to just write H⁺.

So those are Arrhenius acids, but there are also Arrhenius bases. They are molecules that, in water, generate hydroxide ions. You want a hydroxide ion? Here you go.
Note that when it has oxygen has one too many bonds, it is positively charged, but when it has one too few bonds, it is negatively charged. Unlike H⁺, hydroxide ions actually can and do exist in water. The shorthand for hydroxide is ⁻OH. That's how I would notate it if I were writing stuff down by hand, but superscripts and subscripts, although necessary for chemical notation, can be annoying to do on Blogger, so I'll probably stick to just typing "hydroxide" most of the time.

Now might be a good time to mention that in aqueous systems (systems with water as the solvent—I'm not going to elaborate on this for now), hydronium and hydroxide act as a sort of currency of acidity and bacisity. When a reaction takes place, the actual atoms from the acid/base aren't the ones that are participating. I'll illustrate this with a classic acid-base reaction...

HCl + NaOH → NaCl + H₂O.

So, hydrochloric acid and sodium hydroxide yield sodium chloride and water. Sodium chloride is the salt in this case. In other reactions, other salts would be formed. A salt and water are the products of Arrhenius acid-base reactions. But keep in mind that water is the solvent in which this whole thing is taking place. HCl is a gas and NaOH is a solid. When we do this reaction in a lab, we're likely to just mix samples of water that have the compounds dissolved in them. The bond between the chlorine and the hydrogen breaks and the hydrogen reacts with a water molecule to form hydronium while the chlorine atom becomes a chloride ion and sits there dissolved in the water. The bond between sodium and oxygen likewise breaks and we're left with hydroxide and a sodium ion, which also sits there dissolved in the water.

The hydronium ion produced by dissolving the acid in water can and will react with another water molecule. But the product of that reaction leaves the original hydronium ion turned into an ordinary water molecule and the water molecule attacked by hydronium as the new hydronium ion. This process occurs repeatedly, but without really changing anything because the number of hydronium ions remains constant. The same is true with the base. If hydroxide reacts with a water molecule, the hydroxide gains a proton and is now a water molecule while the water molecule it reacted with lost a proton and is now a hydroxide ion. The charges are rapidly transferred from one water molecule to the next, but they remain there. This tranfer would continue for a long time, but when we mix the two solutions, in addition to reacting with water, the hydroniums and hydroxides can react with each other. Whenever this reaction takes place, the two ions neutralize each other and we're left with only water (H₃O⁺ + ⁻OH → 2H₂O). This reaction also produces heat, so if you do it at home for some reason, keep that in mind so that you don't die or whatever. If you were wondering, the sodium and chloride ions stay dissolved. You have seen what happens when you add sodium chloride to water, right? Seriously, you'd better have. If not, go do it right now.

So that's the Arrhenius definition. There are several other definitions, some of them relevant to certain fields, but two definitions are by far the most popular and important, so I'll introduce them now.

Brønsted-Lowry definition
This is the definition that seems to be used the most in my textbook and introductory chemistry courses. It is more inclusive than the Arrhenius definition. A Brønsted-Lowry acid is a proton donor. A Brønsted-Lowry base is a proton acceptor. One big difference between this and the older Arrhenius definition is that water is not necessarily present as a solvent. It can be and often is, but it's not necessary. The most important difference to keep in mind might be that not every Brønsted-Lowry base will lose hydroxide. Every Brønsted-Lowry acid must have a hydrogen atom that can be lost in order to give off a proton, but the bases need not look anything like Arrhenius bases. What a Brønsted-Lowry base does need is the ability to form a bond to a proton. That means it needs free valence electrons that can be recruited to form the bond. Lone pairs of electrons work best, but electron pairs in π-bonds also have basic potential.

My next few posts will be dealing with Brønsted-Lowry acids and reactions with them, so this is the definition to pay the most attention to. Hopefully, you'll get a feel for what chemicals act as good acids and bases and how acid-base reactions work. For now, definitely remember that an acid is a proton donor and a base is a proton acceptor.

You might have deduced that once an acid has donated a proton, what's left is capable, under the right circumstances, of accepting a proton and recreating the original acid. Likewise, a base that accepts a proton now has a proton that could be donated to something else, recreating the original base. These are conjugate acid/bases. That is, whenever an acid-base reaction occurs, the acid becomes its conjugate base and the base becomes its conjugate acid. The reaction could potentially reverse. Here's one reaction...

H—A (acid) + :B (base) → :A⁻ (conjugate base) + H—B⁺ (conjugate acid)

Now we'll reverse the reaction...

:A⁻ (conjugate base) + H—B⁺ (conjugate acid) → H—A (acid) + :B (base)

The system can bounce back and forth between these two. So where does that leave us? The rule I learned in general chemistry and used all the time in my organic chemistry classes was that the equilibrium favors the side with the weaker acid. Unhelpful if you don't know which acid is weaker, but we have ways of figuring that out.

Under the Brønsted-Lowry definition, acidity and basicity are relative to what something is reacting with. Water, for example, is an acid when it's reacting with hydroxide (it donates a proton) and a base when reacting with hydronium (it accepts a proton). Also, you might have deduced that the conjugate acid of water is hydronium and the conjugate base of water is hydroxide. If you deduced that, good job. You get a gold star.

Lewis definition
This definition was formulated by Gilbert N. Lewis. You do know who he was, right? A Lewis acid is an electron pair acceptor. A Lewis base is an electron pair donor. This definition is more inclusive than the Brønsted-Lowry definition. A base still needs free valence electrons to donate, but the acid in the reaction no longer needs to provide hydrogen. So while all Lewis bases are also potential Brønsted-Lowry bases, not all Lewis acids are Brønsted-Lowry acids.

I'll say more about Lewis acids later. This material is in the last section of the acids and bases chapter in my textbook and it's time to move on to some specific Brønsted-Lowry acid-base reactions.

Electronegativity

That last post was pretty long, huh? I'll bet it was too long for you. Well, you're in luck. This post is going to be a short one because I'm tired and electronegativity seems pretty easy to explain. It should be easy to explain, even to someone like you. That's why I'm skipping some stuff in this book that looks like it's about molecular orbital theory. Do you think you can handle molecular orbital theory? Yeah, maybe you can. I don't know. But for now, we're not going to worry about it.

Electronegativity is a property atoms have. Atoms that are highly electronegative have a high affinity for electrons. Atoms that are not very electronegative have a low affinity for electrons. Is that simple enough for you? Here. I'll personify the atoms for you, just like every high school science teacher on the planet has ever done. Atoms that are highly electronegative want electrons. They will take them. Atoms that are not very electronegative are more likely to lose their electrons to those mean, electronegative bullies. Now I feel all dirty.

Anyway, electronegativity is inversely proportional to atomic radius. So smaller atoms tend to be more electronegative than bigger ones. There are various explanations for this, but I took general chemistry a long time ago. I usually look at it in terms of "shielding" although technically, there's more to it. For an example, let's compare oxygen to sulfur. They're in the same group and have similar chemical properties, after all. Now, oxygen has eight protons, but sulfur has sixteen. And if you were good and read the second post I made here, you know that protons reside in the nucleus of the atom. So all those protons in sulfur are presumably twice as attractive to any electrons as the only half as many protons in oxygen. If that were the whole story, we'd expect sulfur to be more attractive to electrons. But sulfur also has twice as many electrons sitting in its orbitals. And they take up more space. Hence the mention of atomic radius. The sulfur atoms doesn't just have more subatomic particles. It is spatially bigger. And that large, negatively charged space has a repelling effect on other electrons. It's not strong enough to override the attractive effect the nucleus has on electrons. It doesn't shield the nucleus that much. But the effect is enough to make sulfur's electronegativity lower than oxygen's.

Of course, not all elements are in the same group as oxygen. And that matters because the number of valence electrons is also very important when it comes to electronegativity. One very simple way to look at electronegativity is that it increases as one moves to the right on the periodic table (ignoring the very last column for reasons that are obvious if you take general chemistry) and decreases as you move down the periodic table. The most electronegative element of all is fluorine. Second is oxygen, followed closely by chlorine. Fourth is nitrogen followed closely by bromine, with iodine being sixth. Whenever one of these elements forms a covalent bond with some other element I didn't just name, the electronegative one will have more of the electron density from the covalent bond.

By the way, carbon is more electronegative than hydrogen, but only by a little bit. It's not usually enough to concern us. But when oxygen or nitrogen (or any of the halogens) is in a molecule, electronegativity becomes important.

Lewis Structures

I am not pleased with the pace at which I've been updating this blog so far. I am still getting the hang of it and I think with more practice I'll do better, but right now it's frustrating. Checking out my textbook again, I'm not sure what to do about the whole "homework" thing. The first chapter has 84 problems. I could do them all. I do have the solutions manual if I get stuck and I generally remember this stuff. That's not the issue for me. I don't want to do these problems. They look too easy. At least from what I've skimmed through, this is really basic stuff. I don't want to plod through it before moving on to the problems I really need to work on. I apparently thought that I would work on problems and write here about the concepts I'm studying. But I'm just not motivated to do 84 problems that will mostly be really easy anyway. Maybe I'll make some sort of split where I work on the problems later in the book and write about the simpler stuff. I don't know if that could work.

Enough about me. The first chapter of this book sure seems to think that Lewis structures are important. So we'll talk about Lewis structures. This would all be so much easier if I could write stuff by hand. The QWERTY keyboard does not, to my knowledge, really do Lewis structures. I know I could easily go over Lewis structures in some sort of classroom setting, but here on the web, I don't have even have a whiteboard. Or do I?
Well, it's a toy I'll need to practice with (practice, practice, practice--it's all about practice).

So with Lewis structures, my book cites three general rules

1. Draw only the valence electrons.
2. Give every second-row element an octet of electrons, if possible.
3. Give each hydrogen two electrons.

Also, a line represents a covalent bond. As aslways, each covalent bond is made up of two electrons. Here's a Lewis structure for methane.


I could make all the bonds the same length, as they are in the real molecule, but I'm really bad at drawing and crap, so you'd better get used to it now. Also note that this is not what the molecule actually looks like. This is just the Lewis structure. The hydrogens repel one another and so the configuration they'll be in is the one where each hydrogen is as far away as possible from each other hydrogen. Last time I checked, the world is not two-dimensional, so the actual molecule ends up with a tetrahedral shape. Each hydrogen is in one corner of the tetrahedron and carbon is in the center. I'll save you some time and tell you that the angle between any two of those bonds is 109.5°. But on paper, it's a whole lot easier to just ignore the whole third dimension thing and draw the Lewis structure. Next, let's try the Lewis structure for water.

The two non-bonding pairs of electrons are represented as dots. They're in pairs because that's how electrons roll. They use the buddy system. Actually, it's because each atomic orbital holds two electrons, but that's not what we're worried about right now. The molecule has a bent shape because non-bonding pairs take up more room than bonding pairs due to greater repulsion force and compress bonding electrons closer together. You did know that, right? Please say you knew that. Anyway, the point of Lewis structures isn't really to represent bond angles. But it is possible to have the atoms line up in a straight line. Here's carbon dioxide.
See? The carbon has no non-bonding electrons, so they can't compress anything and the whole molecule has a linear shape. Both oxygen atoms have non-bonding pairs, but think about it. What would they compress? Yeah, the linear shape keeps those electrons as far from each other as they can possibly get. And that's Lewis structures. How about one more? I'll make it one of your favorite organic compounds.

Here's some of what you'll need to know in order to have an idea of what I'm talking about...

I can't relate all of the fundamentals of chemistry in a single post. A lot of what I'd like to get across would best be done with drawings, images, graphs, and such. I'm too used to using words, but I'll try to provide some images, starting with this very post. Well, let's introduce the periodic table.

I hope you've seen it before. And if you haven't, what is wrong with you? There a lot to say about this table. It's one of the most important developments in science. It's fascinating and I recommend learning about it, but if you find it intimidating now, you can take solace in the fact that I'll only be using a small portion of it for this blog. Here, I'll show you.

See? It's the same table, but I erased everything we don't really care about. Some of the elements I erased from the table will actually be important in the future, but not for a while. We'll cross that bridge when we come to it. I also shaded in the elements we care about most. And out of the ones I didn't shade in, each one is pretty similar to the ones it shares a column with. That means we're really looking at about eleven different things, rather than over a hundred.

Of course, there's the issue of what these elements actually mean and why they're placed on the table the way they are. Each element is a substance that has its own unique atom. Atoms have three building blocks: protons, neutrons, and electrons. Neutrons are the heaviest. They have some important properties, but for our purposes they don't do much besides sit there in the middle of the atom, a region called the nucleus. Right beside them are the protons. Every atom has a specific number of protons, which is its atomic number. The number of protons is what makes an atom have the chemical properties that it does. And the table organizes them according to this number. You can see that hydrogen (atomic symbol H) is the very first one. Hydrogen atoms each have one proton. Helium (atomic symbol He), which I erased in the second table, is second and has two protons. Lithium (atomic symbol Li) is third and has three protons. This continues for the whole table. Every proton has a positive charge, often represented as +1.

Atoms also have electrons. But electrons don't sit in the nucleus. They buzz around in an area surrounding the nucleus. Electrons are much smaller than protons, but each one has a negative charge that cancels out the positive charge from a proton, often represented as -1. This means, of course, that an atom with the same number of electrons as it has protons will have no charge (it's neutral). Electrons interact in three-dimensional geometric structures called "orbitals" that I happen to find very annoying to draw. I'll cover those at some point in the future. What's important is that these orbitals take up a lot of space. Even though the nucleus is much heavier than all of the electrons combined and would take up more space than all of them, the negatively charged area of the orbitals engulfs the nucleus. I've heard the size difference likened to a golf ball sitting in the middle of a stadium, with the golf ball being the nucleus (with almost all of the mass) and the walls of the stadium being the outermost electrons.

The electrons on the highest or outside orbital are valence electrons. They're the ones that form bonds to other atoms. Chemistry deals with the countless ways in which bonds interact, where they can form, what it takes to break them, etc. Electrons in this "valence shell" typically exist in pairs (there's a reason for this, but it's another thing that I'm skipping for now). With this in mind, I'll go over the interactions we expect to see in bonds with our important atoms.

The alkali metals:

The alkali metals we're worried about are lithium (Li), sodium (Na), and potassium (K). And really, they're all pretty similar to each other, which simplifies things for us. Alkali metals are very reactive. They'll typically form ions. Ions are like atoms, but with either fewer or more electrons than the number of protons in the nucleus. This happens because the atoms either lose or gain valence electrons. An atom that has lost valence electrons is called a cation and has a positive charge (because it has more protons than electrons). An atom that has gained valence electrons is called an anion and has a negative charge (it has more electrons than protons). Alkali metals all have one valence electron. When they react, they become +1 charged cations. The most famous example of an ionic compound is sodium chloride, common table salt. The sodium cations and the chloride anions (anions of chlorine) are attracted to each other because they're oppositely charged (you did know that opposites attract, I hope). We won't be dealing with metals much, because this is organic chemistry. But it is important to understand ions.

The alkaline-earth metals:

The alkaline-earth metals we're worried about are beryllium (Be), magnesium (Mg), and calcium (Ca). When I was taking organic chemistry, I dealt with them even less than the alkali metals (and I didn't deal with berylium at all). But like the alkali metals and some other metals, these can come up when we get into organometallic chemistry (which won't be for a while). The only important thing to keep in mind about these metals is that they have two valence electrons and will react to form cations by losing both. I used sodium chloride as an example above. The equivalent to that here would be magnesium chloride, which is less well-known, but it is used to make tofu and is sometimes used as a dessicant (it absorbs water from the air). Because it needs to lose both of its valence electrons to become a stable cation, magnesium chloride crystals actually have pairs of chloride ions attached to the magnesium ions so the ratio of anions to cations is 2:1, whereas in the case of sodium chloride, the ratio was 1:1. Enough about metals, though. We'll be dealing almost entirely with non-metals.

Boron:

Boron is of some importance. When I reach material that deals with boron, I'll give more details on boron. For now, just note that boron has three valence electrons (it has two other electrons in a lower orbital, and only valence electrons are used in bonding). So boron will form three bonds. When these bonds are to other non-metals, they're generally covalent bonds. Covalent bonds are links between atoms generated by sharing of valence electrons. Boron forms three.

Hydrogen:

Hydrogen is a very simple case. Hydrogen atoms can bond to only one atom each. That is, each hydrogen atom can form only one bond. If two hydrogen atoms are bound to each other, it's a molecule of hydrogen gas. But most of the bonds we'll be seeing between hydrogen and other elements will be with hydrogen and carbon, nitrogen, or oxygen. Again, the rule for hydrogen is simple: hydrogen can only form one bond. Hydrogen can also form ions, but we'll deal with those later.

Carbon:

Here's where things get more complicated. Carbon has four valence electrons and forms four bonds. That means any given carbon atom could easily be bonded to four other atoms. But that's using only sigma bonding. There's also pi bonding. Rather than worry about that for now, just remember that double and triple bonds are also possible. So a carbon atom can be bonded to two, three, or four other atoms. Two bonds would mean that is forming either a triple bond and a single bond or two double bonds. Three bonds would mean that the carbon is forming a double bond and two single bonds. Four bonds would mean that it has single bonds to four different atoms.

A double or triple bond to hydrogen is impossible. What I said earlier about hydrogen only forming one bond also excludes double and triple bonds (those require electrons too). But carbon is capable of forming bonds to other carbon atoms, single, double or triple. Long chains or rings are possible. It's this versatility that makes organic chemistry unique.

Nitrogen:

Nitrogen has five valence electrons. One pair will always stay with the atom, so it can form up to three bonds. A nitrogen atom with a triple bond to another nitrogen atom forms a molecule of nitrogen gas, and you should know that this gas makes up most of the air you breathe. Triple bonds between nitrogen and carbon are also possible, as are double and single bonds.

Oxygen:

Oxygen has six valence electrons. Two pairs will always stay with the atom, so it can form two bonds, either two single bonds or one double bond. No triple bonds are possible with oxygen. Oxygen double-bonded to oxygen forms a molecule of oxygen gas, which makes up most of the rest of the air you breathe (the part that isn't nitrogen).

The halogens:

The halogens (the only ones that matter, anyway) are fluorine (F), chlorine (Cl), bromine (Br), and iodine (I). The differences will eventually come up, but as far as the fundamentals go, they're all interchangeable and when using atomic symbols to represent atoms, "X" in organic chemistry means one of these elements. They have seven valence electrons. Three pairs will always stay with the atom. One electron can participate in covalent bonding.

Silicon:

We won't be using silicon much. For now, all you need to remember is that, like carbon, silicon is tetravalent. Again, this means that it forms up to four bonds.

Phosphorus, and sulfur:

These two get a bit tricky. Phosphorus behaves a lot like nitrogen and sulfur behaves a lot like oxygen, but they can do other things too. These elements can form more than four bonds. Even though phosphorus will usually form three like nitrogen and sulfur will usually form two like oxygen, both can, in certain situations, form several bonds. So while you should expect not to see a carbon with five bonds to it, this can totally happen with these two elements.

That's all for now. I should probably cover some more, but I want to post this now.