Constitutional Isomers

Earlier today I said that I was going to do a post on isomers tomorrow. Well, I'm so anxious to talk about isomers that I am starting it early. Only half an hour left until tomorrow anyway as I'm typing this sentence, so maybe I won't finish it until it's tomorrow. We'll see.

Constitutional isomers are compounds that have the same atoms, but arranged in a different structure. This is completely different from resonance structures, because both isomers are real molecules and can have very different properties. So don't get the two mixed up. Remember, resonance structures have double-headed arrows between them. Isomers don't. Because I'm lazy, I'll just use the first example I find in my book...
Both molecules have the same atoms. They each have three carbons, one oxygen, and six hydrogens. But in molecule A, the double bond is between two of the carbons and the oxygen is bonded to a hydrogen. In molecule B, the double bond is between the oxygen and the carbon it's attached to, while the other carbons each have three hydrogens.

Remember how it's bonding that really makes molecules what they are? These compounds have different bonding structures, so I would expect them to have very different chemical and physical properties.

Addendum: People should know what common or important molecules look like. I just realized, as I was about to close my textbook, that "molecule B" is, in fact, acetone, a chemical you might be familiar with. Acetone is an important solvent. I used it in the laboratory all the time in my chemistry classes. It's also in paints and stuff. And it's used to make acrylic glass. As for the other molecule (A), its name is 2-propenol and I'm pretty sure it's just an enol form of acetone, so basically it is highly unstable and will typically turn into acetone by itself. But that's a topic for another day.

Stephen Bahl Has Triple Bonds

So I totally realize that I can use a symbol for a triple bond. Yep. It makes things easier. Like here's ethyne, also known as acetylene...

H—C≡C—H

I chose that molecule because I didn't even need MS Paint to do it. Awesome. I'm so excited to have that symbol at my disposal...

≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡≡

Awesome. The symbol is actually available because of its meaning in logic and mathematics, not because of anything to do with chemistry. But I'll take what I can get.

Resonance Stabilization

Once again, it's been a while. I had so much fun with the Lewis structures post that I wanted to make a new post the very next day, but other things kept getting in the way. Well, I'm not putting it off anymore. My introduces resonance after introducing Lewis structures, so I suppose I'll follow its lead...

Many molecules, both organic and inorganic, are resonance stabilized. Such molecules cannot be represented by a single Lewis structure. For that, we can use resonance structures—if we want to. I don't really like resonance structures myself. But you know, whatever works. Before I go any further with resonance structures, I must explain why they're used at all. They exist because of the delocalization of electrons. Delocalized electrons aren't sitting on any one atom and they're not locked into any one bond. That's why we use the word "delocalized." It's like they're spread out over multiple atoms as one big forcefield of negative awesomeness. I don't know. Let's just move on to the example.

The example my book uses is this anion. Don't worry about its name because I don't know its name either. It's the conjugate base of formamide and it's an anion, but beyond that, I have no idea. Here's are the resonance structures...
Yes, the circled "-" sign is a negative charge. That was obvious, right? That double-headed arrow indicates that these are resonance structures. The important thing to realize is that these structures aren't two molecules. They're representing the same molecule in two different ways. The electrical charge is in two different places, but all of the atoms are in the same places. The first structure makes the oxygen the center of negative charge. The second structure makes the nitrogen the center of negative charge. In actuality, the negative charge is distributed between those two (and in this case, it will be slightly more centered on the oxygen because oxygen is more electronegative than nitrogen).

Resonance structures are not real. My professor compared it to describing a rhinoceros, to someone who had never seen one, as a cross between a unicorn and a dragon. Yeah, that doesn't make any sense. He realized that after he said it. But he was trying to think of a more everyday example of using two fictitious things to describe one real thing. Resonance structures aren't always in pairs, though. Many molecules have three resonance structures.

So yeah, not real. Resonance structures. The bond between the carbon and the oxygen isn't actually a double bond. And the bond between the carbon and the nitrogen isn't a double bond either. It's more like the bond between carbon and oxygen is a little bit more than a 1.5 bond and the bond between the carbon and the nitrogen is a little bit less than a 1.5 bond. I don't know the actual numbers. It's possible to take a measurement, but I don't have the equipment or the expertise, so shut up. We'll just pretend that it's exactly 1.5 on both, even though I know that it isn't. Actually, why didn't this stupid book just use an example where that was the case? I mean, there are cases like that. This molecule isn't one of them. Anyway, it's more like those bonds are each 1.5 bonds. But Lewis structures don't have a way of representing fractional bonds, so we need a way to represent what's going on here. Some day, I'll fix the notation in chemistry. Until then, we're stuck with resonance structures.

This post was probably kind of confusing. I know I get resonance structures. But maybe you don't. Remind me to help make this post more clear. Right now I need sleep. I'm totally going to write about isomers tomorrow.

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.

Vitalism and the Origins of Organic Chemistry

Time flies. I can't believe how long it's been since my last post here. I am slacking a little, but part of the reason that posts here have been infrequent is that I've been spending more time with friends, which is something I had a goal of doing. So I'm not too upset. And I am here, right now, updating this blog. So here we go...

In order to appreciate the distinction between organic chemistry and the rest of chemistry, it take some appreciation for the history of the science. Organic chemistry is frequently defined as being the chemistry of carbon or the chemistry of compounds containing carbon.

My textbook puts it rather simply:

Organic chemistry is the chemistry of compounds that contain the element carbon.

Wikipedia is a bit more specific:

Organic chemistry is a discipline within chemistry which involves the scientific study of the structure, properties, composition, reactions, and preparation (by synthesis or by other means) of chemical compounds that contain carbon.

What neither of these simple definitions tell us is that not all compounds containing carbon are considered organic. My textbook doesn't seem to mention it, but looking up "organic compounds" on Wikipedia reveals this:

For historical reasons discussed below, a few types of compounds such as carbonates, simple oxides of carbon and cyanides, as well as the allotropes of carbon, are considered inorganic.

This is because of the now defunct concept of vitalism. People believed that organic matter and inorganic matter were fundamentally different (as an aside, some attribute this to Aristotle, but I haven't looked into it). Carbonates, oxides of carbon, and cyanides (and carbides, another class of inorganic carbon-containing compounds) are all found naturally outside living systems (in minerals, for example). Other carbon-containing compounds were only known to be associated with life. Then this whole view of things got wrecked in 1828 by Friedrich Wöhler doing this reaction:
Pretty cool, huh? Wöhler combined ammonia with a solution of cyanic acid ammonium chloride (dissolved in water) with silver cyanate and got, as a product, urea, which was previously only known to be produced by the kidneys of animals (mammals and some other animals produce urea as a waste product). Although this didn't immediately strike the deathblow for vitalism (which has, in some form or another, survived to this day, although thankfully not among chemists), it laid the foundation for organic chemistry as a field (these compounds were now something that could potentially be synthesized in laboratories and there wasn't necessarily any essential "life force" that was generating them).

The distinction between organic chemistry and inorganic chemistry is important today because organic compounds have some specific properties of their own that can be studied in detail and because organic chemistry is so important in biological systems. Being biological systems ourselves, we have an interest in organic chemistry as it relates to our health. Pharmacy is one branch of applied organic chemistry.

Addendum:

I forgot to mention this, but before taking organic chemistry, I was under the impression, from general chemistry, that organic compounds had both carbon and hydrogen. My organic chemistry professor pointed out early on that technically, not all organic compounds have hydrogen. For example, carbon tetrachloride (CCl₄) is considered an organic compound, but has no hydrogen. However, most organic compounds do contain hydrogen.

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.

Welcome

This is, obviously, my first post on this blog. And what a marvelous start we have: two instances of self-reference in the first sentence. One week ago, I expressed my intention of starting a blog like this one.

I still don't know exactly what this blog is going to look like, but my intention is for it to be didactic. This will require me to know things, but knowing things is exactly what I've always wanted to do. We'll see if it works. I am very serious about this.

For now, I'll be writing about organic chemistry. I want to focus on making this accessible to people. If you know basic chemistry, you should be able to follow this blog easily. If you don't know any chemistry, at all, well, you should really learn some, because it's good stuff. If something is unclear, please ask me about it and I'll try to expound on the topic. Like I said, I want this to be accessible to people. I'll also try to write some "review" posts that cover basics. I don't really have any such topics planned yet.

For those of you who don't know much chemistry or have forgotten most of what you did know, later this week, I'll make a post that will cover some fundamentals. That is all for now. I hope this blog provides some value, either for you or for myself, or maybe both of us.

Edit: I was out of the house more than I anticipated this week. I was actually planning on making two more posts here, but that number turned out to be zero. The one about fundamentals will be arriving sometime next week, though. I could have tried to do it tonight, but I'd rather make it late than inadequate.