Examples of Naming Acyclic Alkanes

As promised, here are some right out of the textbook.

The first one is in condensed notation: CH₃CH₂CH(CH₃)CH₂CH₃

Since I am so good, I immediately recognize that the third carbon has a one-carbon branch. Other than that, this is a straight chain. But let's not get ahead of ourselves. We are doing this the right way. We start with the last part of the name. With no heteroatoms, this is a hydrocarbon. With no multiple bonds, it's an alkane. With no rings, it's an acyclic alkane. We know the name must end in "-ane." Next, what's the longest carbon chain? Five. So the parent name is pentane. Branches? Yes, at the third carbon (counting either way). And the branch is a methyl group. Therefore, the name of this compound is...

...3-methylpentane. And you know what else? I checked the answer in the study guide and I was right! Woo hoo, Stephen got something right. Anyway...

(CH₃)₃CCH₂CH(CH₂CH₃)₂

This one is harder. First we have three methyl groups attached to one carbon. That carbon links to another that links to another, which is attached to two ethyl groups. Which methyl group and which ethyl group is considered part of the longest carbon chain does not matter because the groups are identical (that is, the methyl groups are identical to each other and the ethyl groups are identical to each other). So adding those three carbons to the rest of the chain, we find that the longest carbon chain is six carbons long, so this is a hexane.

Which group gets priority? In this case, we go in alphabetical order. "E" comes before "M." So this should be...

...3-ethyl-5,5-dimethylhexane. Or not. Oops. I started at the wrong end. It's actually 4-ethyl-2,2-dimethylhexane. It's that instead of the one I thought it was because 2 is lower than 5. It doesn't matter that 3 is lower than 4 because the method that gives the lowest number period is the one that gets priority, not the one that gives the lowest sum or anything like that. I hope you learned your lesson. Moving on.

CH₃(CH₂)₃CH(CH₂CH₂CH₃)CH(CH₃)₂

A propyl group? No, that's part of the longest carbon chain. They're trying to trick us. Starting from the left we have a carbon and then a string of three more, so that's four in a row. Then there's another (five) with that propyl group branching off. If we count going up the propyl group we get three more (eight). If we treat the propyl group as a branch, we get another carbon with two methyl groups, one of which would be a branch, making the total length seven. Sneaky textbook. This is actually an octane.

If we start from the end of what's being labeled as a propyl group (but is actually part of the chain) we get a branch at the fourth carbon. Starting from the left makes it at the fifth, so we start from the end of the propyl group instead. The branch consists of three carbons and two of them are attached to the other, which is where the branch connects, so it's an isopropyl group, meaning the compound is...

...4-isopropyloctane. And I'm right. I rule.

Enough of these condensed structures!
I used MS Paint because it was a small one and it's kind of hard to make them look less awful on ChemSketch. Anyway, this one seems easy to me. Five carbons long means pentane. Two methyl branches at the second carbon and two at the fourth. Therefore...

...2,2,4,4-tetramethylpentane. And I am right again. Excellent.
It's seven carbons long, but there are a couple of different ways to arrive at that. The one that give the lowest number to a branch is the one that simply starts on the far left, for a methyl group at the second carbon. There's another one at the fifth carbon and an ethyl group at the third, so this is...

...3-ethyl-2,5-dimethylheptane. And I'm right yet again. Three in a row! Let's do one more.
I moved back to MS Paint again when I perhaps should not have. But ChemSketch was being annoying (it kept trying to put rings into this). Obviously this one is larger than the other ones so far, but the principle is the same. The longest carbon chain is ten. The fastest we can get to a branch with it is on the second carbon, again counting from the far left. From there we label the other branches and put them in the proper order. About that, the branch on the fifth carbon is a sec-butyl group. When alphabetizing the branch names, this is treated as a "B" and not as an "S." The same would be true for tert-butyl but not for isobutyl. Unnecessarily confusing, I know. But in this case it does slightly affect the name, which is...

...5-sec-butyl-3-ethyl-2,7-dimethyldecane. And that's pretty much all there is to it. The study guide I used to check my answers breaks the process into three steps.

1. Name the parent chain by finding the longest C chain.
2. Number the chain so that the first substituent gets the lower number. Then name and number all substituents, giving like substituents a prefix (di, tri, etc.).
3. Combine all parts, alphabetizing the substituents, ignoring all prefixes except iso.

It takes some getting used to, but this is the basis for how other compounds, even ones with multiple functional groups, are named.

Nomenclature of Acyclic Alkanes: Prefix

I hope you have the other component of naming alkanes down, because I am never reviewing it again (just kidding). Now for the prefix. While the parent name identifies the longest carbon chain, the prefix tells us where on that chain branches occur and what the branches look like. Depending on how much branching (and what kind) is going on, the prefix may be anywhere from nonexistent (no branches, which we sometimes denote by using "n" as a prefix) to ridiculously long.

Firstly, the location of a branch is denoted using Arabic numerals. A branch at the second carbon in the longest carbon chain gets a "2" and a branch at the third carbon gets a "3" and so on. Some carbons in the longest carbon chain might have two branches. When that happens, its number gets used twice.

Often, there are multiple possible places to start from. With alkanes, the correct starting carbon is the one which, when started from, yields the lowest possible number being named first. If we start counting on one end of a chain and the first number that comes up is for a branch at the fourth carbon, but counting from the other end of the chain would make our first branch be at the second carbon, then it is the end that would make the first branch be at the second carbon that is the correct starting point.

Also, numerals are separated from each other by commas and from the rest of the name by hyphens. That's not just for alkanes. That's a universal rule. Commit it to memory, slave.

Anyway, to specify how long a branch is, we use the wonderful numerical prefixes I introduced in my last post. You know, the ones that are mostly Greek, but not really. A branch that is only one carbon is a "methyl" group. Two carbons is an "ethyl" group, etc. A branch that is seven carbons long is a "heptyl" group (and since it's not part of the longest carbon chain, that means the longest carbon chain must be really long). This all works nicely for branches that are themselves straight. But what about branches that have branches of their own? That's the hard part. Kind of. In order for considerable branching to occur, the molecule itself has to be pretty big. I've never had to deal with such compounds myself. The textbook is covering substituents with up to four carbons and that's always been good enough for what I've had to do. There are not very many. Here we go...

Methyl group: R—CH₃
Ethyl group: R—CH₂CH₃
Propyl group: R—CH₂CH₂CH₃
Isopropyl group: R—CH(CH₃)₂
Butyl group: R—CH₂CH₂CH₃
sec-Butyl group: R—CH(CH₃)CH₂CH₃
Isobutyl group: R—CH₂CH(CH₃)₂
tert-Butyl group: R—C(CH₃)₃

If you find the condensed structures confusing for those four-carbon groups, here are some links to pictures (off-site) for the butyl variations...

Butyl, sec-butyl, isobutyl, and tert-butyl.

And that's all. Now you know how to name acyclic alkanes. Oh, one more thing. If two or more of the same type of branch exists in a molecule, those branches get named together and get a Greek numerical prefix just to confuse you even more. But really, that's it. Stay tuned for next time, where I'll do a follow-up post with some examples of naming alkanes using problems from the textbook. Oh wait, this isn't a radio. You can't tune anything. Whatever.

Nomenclature of Acyclic Alkanes: Parent Name

I mentioned the IUPAC systematic nomenclature system before. I think I did, anyway. This project has been on hiatus for a while and I can't remember. But I'm back now! Really. I hope. Anyway, today we are going to learn how to name some alkanes. It's easy to do, and you need to know it to name other compounds. So learn it. I command you.

Let's start at the end. That's a good place to start, right? The last part of the name of any alkane is, get ready for this...

...it's "-ane." That should be quite easy to remember, even for you, because "alkane" itself ends in "-ane." If a compound is an alkane, its name ends in "-ane" and, conveniently enough, if a compound is not an alkane, its name will not end in "-ane." I know. Chemistry is so hard.

Next, we find the longest carbon chain. This is actually very easy, but teachers love trying to trick beginning students with odd drawings where they make part of the longest carbon chain look like a branch to people who are not paying attention. If this were a real chemistry class and I were the teacher (that would be bad), I would totally do this to you because I think it's hilarious. For now, I'll just give you the benefit of the doubt and assume that you are paying attention and can tell what the longest carbon chain in a molecule is.

Really? I shouldn't do that? Fine.
How long is the longest carbon chain? If you answered eight, congratulations, you did not fall for the dumbest trick in chemistry class. If you answered some other number, you were not paying attention or you cannot count or you're just a moron or something. I don't know. Shame on you anyway. You're bad (unless you got the right answer).

Once we know how long the longest chain is, we convert that into a numerical prefix, then attach it to our "-ane" suffix. Convert it into a numerical prefix? Yes, it's easy. No really. It is easy, just so long as you already know the Greek numerical prefixes—and use the Latin one for "nine" just to mess things up—and forget the first four prefixes and make up new special ones that are specific to chemistry. It was easy for me though! Here, I'll give you the first ten and we'll worry about going higher later.

1 = "meth"
2 = "eth"
3 = "prop"
4 = "but" (pronounced like the word "butte" just to confuse you even more)
5 = "pent"
6 = "hex"
7 = "hept"
8 = "oct"
9 = "non" (pronounced so that it rhymes with "tone" and not some other way)
10 = "dec"

Memorize them now. I command you. Done? Good. See, that wasn't so bad. Now, there's just one more tiny thing. Then we'll be all done and you'll know how to name acyclic alkanes. We have straight chains covered (unless they're longer than ten carbons long, but shut up). So a hydrocarbon that is a straight chain with five carbons would be "pentane" and one with nine would be "nonane" and so forth. Everything is fine, and then branches come and mess it all up. Not to worry: the IUPAC has an elaborate set of rules for us to denote where on a chain the branches lie and what the branches look like using prefixes and attaching them to the parent name (which simply describes the longest carbon chain. Well, it's elaborate enough that I'll save it for my next post, anyway. For now, just have the whole parent name part down.

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.

Ribosome Rant

I realize this is a departure from the content I normally post here, but I just started writing a rant on a different site and I think it really belongs here. The Nobel Prizes are being announced this week. The prize in chemistry went to Venkatraman Ramakrishnan, Thomas A. Steitz, and Ada Yonath for their work on the the structure of ribosomes. There's a sentiment that I've been seeing somewhat and it got me annoyed enough to actually write this. Here are some examples of the sentiment I am talking about...

I don’t care. For some reason, this year I’m not getting into Wednesday Madness nearly as much as I have in previous years. I’ll be happy if they give it to, uh, a chemist.

≡≡≡

Oh well. Here’s an idea. In lieu of giving out Nobel Prizes in Chemistry to achievements in chemistry (since they only seem to give it to actual chemists every other year anyway, it won’t be much of a stretch), let’s start handing them out to the authors with the best paper titles ever.

≡≡≡

As already announced biologists walked away with this year’s Nobel prize in chemistry once again, this time for work in determining the structure of Ribosomes.

≡≡≡

As chemists we would like to see the Nobel chemistry prize go to a chemist. Our Nobel hopefuls may be a measurable magnitude more chemically interesting, as measured by ChemFeeds, but there is more work for them to do until these topics become world renowned (which seems to be the dominant prerequisite these days).

≡≡≡

And again the Nobel for Chemistry goes to "bio-chemists"....
Congratulations...but as a strictly synthetic organic chemist...I am a bit ticked off.
With all the biology and the nanoscience development in recent years, it'll be eons before an organic chemist wins the prize again.

Alright guys, the applicant to get into school to work on an undergraduate degree (I do already have my A.S. at least) has some news for you: biochemistry is chemistry. I find this reaction deplorable. Chemistry is all about atoms and the bonds between them, what things are made of and how they interact with each other. That is exactly what this prize was awarded for. Perhaps word has not yet reached the innermost confines of your biology-free ivory towers, but ribosomes are made out of atoms and ribosomes have bonds—lots of them. Ribosomes participate in chemical reactions. This really should go without saying.

I could be way off here, but I don't think I would see this in other branches of science. If an annual physics prize went to scientists who did work in astrophysics, would physicists complain that the astronomers are taking physics prizes? I think not, but maybe some of them would. Maybe some of them sequester themselves in ivory towers devoid of any science that is not their own particular specialization, just as apparently some chemists do. My impression is that many, if not most, physicists have a passion for the universe and its fascinating nature. They want to see the physics in everything. I want to see the chemistry in everything. And I'd like to think I'm in good company, but the reactions I've seen to this Nobel Prize have cast some doubts on that.

How arrogant must one be to think, "Only research in the area of chemistry that I focus on should win prizes"? Some might protest that this is an unfair characterization, but if one is willing to dismiss the entirety of biochemistry, I am more than willing to err on the side of assuming that one would go on to dismiss other purportedly unworthy subjects in a similar manner. This exclusive approach is the exact opposite of what I want to stand for. I want chemistry to be inclusive. If we excise some of it because it deals with biological molecules and can therefore be considered biology, we might as well excise the parts that deal with minerals and make that geology and so on until we have divided everything up and there are no more chemists, just former chemists working in other fields of science.

The ribosome people did not win because the biologists are taking over and they did not win because ribosomes are famous and other work was too obscure. They won because they did good chemistry that is of abundant benefit to humanity.

Strength of Intermolecular Forces

This chapter in the textbook is quite long, but not all of it is well-suited to posts like these. A lot of this reviews concepts from general chemistry and has lots of pretty pictures and I don't want to spend too much time on things like melting point and solubility and soap. The soap thing is something I originally learned in high school and got to see repeated in two general chemistry classes in college and organic chemistry too. I might do a post on it, but for me, it's gotten kind of old. There is some really great material here. I especially like the explanations of biomolecules, but perhaps that's best reserved for later.

In short, I do want to write at least one more post on the odds and ends in the third chapter of my textbook. They will come soon if at all, because I am long overdue on starting the fourth chapter. Before I do either of those things, let's wrap up intermolecular forces.

The strength of intermolecular forces is, in ascending order...

• London Dispersion: caused by fluctuations in charge density across the surfaces of molecules.
• Dipole-Dipole: caused by permanent dipoles.
• Hydrogen Bonding: Caused by extreme loss of electron density on hydrogen when bonded to oxygen or nitrogen (or fluorine, technically).

Note that ion-ion forces, which hold ions together in ionic compounds, could be compared to these forces, although ions are technically not considered molecules. Ion-ion forces are much stronger than any of these intermolecular forces.

Intermolecular forces have the following effects on physical properties...

• Stronger intermolecular forces increase boiling point.
• Stronger intermolecular forces increase melting point.

There's more, but I am skipping it because that's the important stuff and you'd forget the rest anyway. If something I omitted here becomes important later, I'll just blame the problem on you. It's either that or explain the thing when the issue comes up.

Hydrogen Bonding

Hydrogen bonding is a special (and awesome) intermolecular force that can happen when hydrogen is attached to nitrogen, oxygen, or fluorine. All three of those are pretty small atoms and are also highly electronegative. They pull so much electron density away from the tiny hydrogen that the nucleus (a single proton) of the hydrogen is highly exposed and attracted to negatively charged things, letting it sort of stick to negatively charged bits of other molecules.

Fluorine is a halogen. I wrote about this, but it was so long ago that you probably forgot all about it. As a halogen, fluorine can only have a bond to one other thing. If that thing is hydrogen, then we have lots of hydrogen bonding fun, but there's only one compound for which this is possible, and that's HF (hydrogen fluoride), which isn't organic, so we won't be paying much attention to it right now.

Nitrogen and oxygen, unlike fluorine, can both be attached to hydrogen and have at least one bond to spare. So there are lots and lots of compounds that exhibit hydrogen bonding. Just look for hydrogen attached to nitrogen or oxygen. Hydrogen bonds are stronger than dipole-dipole forces and other intermolecular forces. In fact, they're partial covalent bonds, although still not nearly as strong as regular covalent bonds.

The classic picture for showing hydrogen bonds is with water, so I might as well just use one of those rather than trying to make my own picture. It's much easier and looks way better. Here's one...
Pretty much any property of water has something to do with hydrogen bonding. And it's important in lots of other compounds too. The structures of proteins and nucleic acids use plenty of hydrogen bonds. Hydrogen bonding is one of the reasons hydrogen is my favorite element. But just remember, it's stronger than other intermolecular forces.

Dipole-Dipole Forces

Do you remember how electronegativity works? I posted about it, so you should. Here, just for you, I'll link back to that post. Electronegativity is necessary to understand polarity. And polarity is how these forces operate. An organic molecule with a functional group that contains an electronegative atom such as oxygen or nitrogen will likely have a permanent dipole. The heteroatom pulls negative charge toward itself (because it's electronegative, obviously). As described in the previous post, regions of the molecule can have differential charge. But a permanent dipole is much stronger than the fleeting changes responsible for London dispersion. Because of this, molecules that have dipole-dipole interaction experience stronger intermolecular forces than ones that have only London dispersion. Compounds with this property are said to be polar and ones that do not are non-polar. Consider these examples...

Acetone
Condensed structure: OC(CH₃)₂
It has a permanent dipole that looks like this.
Those Greek letters represent partial charge. The electronegative oxygen pulls electron density toward itself, so it is the center of negative charge. The region opposite it, lying near the central carbon, is the most positively charged region of the molecule.

Carbon dioxide
Condensed structure: CO₂
This time, when we draw the arrow through one oxygen, the other oxygen cancels it out. Despite having a highly electronegative element, carbon dioxide is non-polar.

Whether a solvent is polar or non-polar tells chemists a lot about its potential uses, and some reactions need one type of solvent or the other. Acetone is a well-known polar solvent, but water is the best known in this category. Many organic compounds such as n-hexane are commonly used as non-polar solvents. That's all for these at the moment. Just remember that in order for dipole-dipole forces to occur, the compound(s) must have permanent dipoles, and that dipole-dipole forces are usually much stronger than London dispersion. Of course, a molecule can have both. But the stronger forces are considered to override the weaker ones for all practical purposes that I've encountered.

London Dispersion Forces

My next few posts will discuss intermolecular forces. The textbook put this information in the same chapter as functional groups, because functional groups play such a huge role in the type and strength of intermolecular forces a compound has. I don't know if this is the best approach, but I'm too lazy to figure out a different one and whatever, this works.

Hopefully the word "intermolecular" tips you off to the fact that these are forces that occur between molecules. If not, what is wrong with you. Intermolecular forces are much weaker than normal covalent or ionic bonds. If they weren't, they wouldn't be intermolecular because they'd be binding atoms as tightly as the bonds within the molecules and the molecules wouldn't really have distinct identities at all.

London dispersion forces are the weakest and most ubiquitous ones we'll look at. In this textbook and in the college chemistry classes I took, they're more often referred to as van der Waals forces. But from what I can tell, that term is actually more inclusive, so I'll be calling them "London dispersion forces" or "dispersion forces" or maybe "London forces" until I have some reason not to. Do note that it seems to be common to refer to these as "van der Waals" forces/interactions. They're also known as induced dipole–dipole forces.

London dispersion is so ubiquitous in organic chemistry in part because it's the basis for intermolecular attraction in hydrocarbons. The other forces don't come into play unless there are functional groups. It would be nice to have the picture in my textbook to convey what I'm talking about, because it really does seem better than what I'm finding on the web, but this ball & stick picture that uses hydrogen molecules should serve well enough...
Because of electromagnetic repulsion, positive and negative charge gather in regions of the molecule. The whole molecule is neutral in charge, but some parts of it are more positive and others are more negative. I like the methane example used in my textbook better, because it shows more randomness, but the principle is the same: the negative portion of one molecule is attracted to the positive portion of another molecule. And this interaction, magnified over massive numbers of molecules, gives the whole thing some cohesion. And when I say some, in the case of hydrogen at normal temperatures and pressures, it's not enough. That's why pure hydrogen is usually a gas. The same is true with methane. If it gets cold enough, as happens in that atmosphere of Titan, one of Saturn's moons, methane is liquid.

Not all hydrocarbons are gases at room temperature. The bigger they are, the higher the temperature has to be in order to excite the movements of the molecules enough to overcome London dispersion forces and liberate the molecules from each other. But it's not all because the molecules are simply bigger. They're also longer. And that means more surface area. My textbook points out that n-pentane has stronger dispersion forces than neopentane. Condensed structures should be enough to show why...

n-pentane: CH₃CH₂CH₂CH₂CH₃
neopentane: C(CH₃)₄

The long, straight chain of n-pentane has more surface area than the bunched up neopentane. They have the exact same quantities of each atom, but this difference has significant effects on the physical properties of these molecules.

Finally, dispersion is also affected by the polarizability of atoms. In a large atom, the electrons are further from the nucleus and are more subject to disruption and fluctuation in electromagnetic forces than in a small atom, where the electrons are more tightly held.

Oh, and a fun trivia fact that everyone loves is that geckos hold onto surfaces using these forces (the bristles on their toes are small enough to do this and well distributed enough that enough surface area is there for the forces to be strong enough to hold the lizard's whole body even when walking upside-down). Seriously. Pretty cool, huh?

Visualizing Functional Groups

I told you that I would post more functional groups and I meant it. But I also want to make the ones I've already introduced clear. And condensed structures can confuse people. I heard you're easily confused. So we'll spend a bit more time getting acquainted with these. I think I covered the hydrocarbons well enough, so we'll focus on functional groups that contain heteroatoms.

More about alcohols
The nature of the carbon that the hydroxyl group is attached to determines the type of alcohol here. First, there's a primary alcohol...
Note that the carbon attached to oxygen is attached to only one other carbon (in the R group). In a secondary alcohol, this carbon is attached to two other carbon atoms...
And as you might have guessed, in a tertiary alcohol, that carbon is bonded to three other carbons...
If you were thinking that a quaternary alcohol would be one in which that carbon is attached to four other carbons, you sure are dumb. Carbon is tetravalent. You remember that, don't you? It can't form five bonds. There is no such thing as a quaternary alcohol.

If you were wondering, whether an alcohol is primary, secondary, or tertiary has important implications for its chemical properties, hence the distinction. This is also the case with amines, as I already alluded to, but in that case, it's how many carbons the nitrogen is attached to that determine which type of amine the molecule is. So there's some cool new information for you. But now for some clarity on material I already covered in my last post.

Visualizing aldehydes & ketones
Here's the Lewis structure of an aldehyde...



And here, for contrast, is a ketone...
Notice the big difference: with an aldehyde, the oxygen is at the end of a chain and with a ketone, the oxygen is attached to a carbon that is somewhere in the middle of a chain. These structures both have a "carbonyl" group and their chemical properties are often similar, but they can be different in important ways and this distinction is certainly worth remembering.

Carboxylic acids and friends
Carboxylic acids get several other classes of compounds grouped with them as "derivatives of carboxylic acids" quite literally because carboxylic acids can be used to make these other compounds. I won't cover all of them because there's a whole chapter on this stuff and it's way later in my textbook. But because you're slow, I worry about your ability to even deduce the general appearance of these groups from a condensed structure. So here's a carboxylic acid...
Like the aldehydes and ketones, there's a carbon double-bonded to an oxygen and single-bonded to an R-group. But the fourth bond isn't to hydrogen or another carbon. It's to oxygen, which itself is attached to hydrogen. Remember acidity? You know, that thing the last chapter was all about and such. And maybe you even remember that in my "Aspirin" post I said, of the carboxylic acid, "This arrangement of atoms makes it easy for a certain reaction to occur. That reaction is a Brønsted-Lowry acid-base reaction." Really, it's not familiar. Whatever. That proton can totally come off.

Since I like functional groups so much, here are some more in condensed structure...

Acyl halide
R—COX (like a carboxylic acid, but with the second oxygen replaced by a halogen)

Imine (imino group)
R=N—R' (these come in multiple varieties and I haven't really studied them yet)

Peroxide (peroxy group)
R—O—O—R' (the oxygens are actually attached to one another)

Nitrile (cyano group)
R—C≡N

Enough. We will now cover new functional groups as they come up.

Functional Groups

Functional groups are structures within molecules that contribute to the properties of that molecule. In his excellent book, The Same and Not the Same, Roald Hoffmann cites the concept of the functional group as something in chemistry that is not reducible to the physical laws that affect it. Functional groups as concepts seem to be uniquely chemical. I, at least, can't think of anything that's more than superficially analogous. And for organic chemistry, they're hugely important.

C—C bonds and C—H bonds are extremely common in organic molecules. If you remember my post showing how skeletal structures work, you should recall that these bonds are only noted with points and intersections for the former and are left to inference with the latter. Such structures are "skeletal" because they really do show the hydrocarbon skeleton of a molecule. All those C—C and C—H bonds are usually pretty stable. They can contribute to the chemistry of a molecule, but not nearly so prominently as functional groups.

I won't attempt to list every functional group here, because there are lots and lots of them and your puny brain would probably die or something. But I will list some of them. First, I want to introduce a new notation. Actually, I'm not sure if I already introduced it, but I'm too lazy to go back and check and you probably forgot about it anyway. The letter "R" is often used to denote the rest of a molecule apart from a functional group, especially if what remains is a plain old hydrocarbon. This can be convenient for situations when it's only the functional group we care about and drawing the rest of the molecule would be impractical or wouldn't even make sense. If we do this more than once though, and the groups being condensed are not identical, using "R" to denote both would be inappropriate, but "R" for one and "R'" (R prime) for another is fine.

So for an example, I've decided to use 2-butoxyethanol because I've used it to clean graffiti off the walls in the restroom at work.
Of course, while this is the stuff I was using, other molecules that change one or more atoms are easily possible. What if we decided only to focus on part of the molecule? We might do this...
What is "R"? Is it still a chain of four carbons with nine hydrogens? Maybe. Or maybe it's still a straight chain, but one carbon longer than that now. Maybe there's a ring. Maybe it has multiple branches. Maybe there are dragons. No one knows! It's a mystery. Exciting, I'm sure. And that's how "R" works. It's a placeholder. It saves space. "Here be dragons" might work too, but "R" is the standard. I have no idea how it became that way, actually.

Aliphatic hydrocarbons
This term comes from the Greek aleiphas, meaning "fat." Actually, many of the properties fats have come from the long hydrocarbon chains they possess. If an aliphatic hydrocarbon has no π-bonds (that is, no double or triple bonds), it is an alkane. Branches and rings might be present and do affect the properties of alkanes, but they're still called alkanes, although a compound with a ring might be said to be a cycloalkane.

But if π-bonds are present, no matter how few there are or how big the rest of the molecule is, it's not an alkane anymore. A double bond means that the molecule is an alkene. So using what we learned earlier, a molecule that contains this functional group: R₂C=CR₂ (different groups are all "R" here because noting them all with superscripts would be ridiculously clunky) is an alkene. The double bond counts as a functional group. Specifically it is the alkenyl functional group.

Similarly, a triple bond is a functional group. Something with R—C≡C—R' functional group is an alkyne (no matter how many double bonds or single bonds it has). And this is an alkynyl functional group.

Aromatic hydrocarbons
The only hydrocarbons I know of that are not classed as aliphatic are ones containing aromatic rings. You might be anticipating this from the trend with what I've said regarding alkenes and alkynes, but the presence of even a single aromatic ring in an otherwise aliphatic molecule means that the compound is considered aromatic and not aliphatic. Aromaticity is a tricky concept though, and we're not covering it just yet. So for now, the only aromatic ring we'll concern ourselves with is the benzene ring.
It is not a cycloalkene, even thought it might look like one. I mentioned in an earlier post that this type of ring is resonance stabilized. The electrons are evenly distributed around the whole ring. How this happens and whether a particular ring is aromatic or not are subjects for later posts. But this ring, the benzene ring, is aromatic. If it's treated as a functional group, it's the phenyl group. To abbreviate when I'm using condensed structures rather than skeletal structures, I'll probably use "Ph" for "phenyl." So something with this functional group would be R—Ph. Other people sometimes just abbreviate a benzene ring attached to something else as C₆H₅ but I'll try not to do that so as to avoid confusion.

Another well-known aromatic hydrocarbon with its own common name is toluene. It's just like benzene, but where benzene has six hydrogens, one attached to each carbon, toluene replaced on hydrogen with a methyl group, a carbon bonded to three hydrogens. So it's like R—Ph with "R" being "CH₃" except this is a special case and gets its own name. If toluene acts as a functional group itself, with one of the hydrogens on the carbon outside the ring being replaced by a bond to the rest of molecule, and I'll draw a picture just to be sure you're following me...
The functional group is actually not phenyl. It's a benzyl group. If this confuses you, good. I like confusing you. Try to remember that this is a benzyl group and that a benzene ring attached to something is just a pheny group. I know both "benzene" and "benzyl" start with "benz-" and you're really tempted, but don't. Just don't.

Now for some more functional groups...

Alkyl halide (halo group)
R—X (where X is a halogen). It could be an alkyl fluoride, alkyl chloride, alkyl bromide, or alkyl iodide depending on which halogen is attached.

Alcohol (hydroxyl group)
R—OH

Ether (alkoxy group)
R—O—R

Amine (amino group)
R—NH₂ (primary) or R₂NH (secondary) or R₃N (tertiary)

Thiol (mercapto group)
R—SH

Sulfide (alkylthio group)
R—S—R'

Aldehyde (carbonyl group)
R—CHO (for those of you who can't deal with condensed structures very well, the carbon is double-bonded to the oxygen, bonded to the hydrogen, and bonded to the R group, so it could also be R—CH=O)

Ketone (another carbonyl group)
R—CO—R' (or R—C=O—R'). The difference between an aldehyde and a ketone is that in an aldehyde, the carbobyl group is on the end of a chain, but in a ketone, it's in the middle of a chain.

Carboxylic acid (carboxyl group)
R—COOH (like an aldehyde with the hydrogen on the carbonyl carbon replaced by an hydroxyl group, alternatively it's like an alcohol with an oxygen double-bonded to the terminal carbon)

Ester (ester group)
R—COOR' (like a carboxylic acid, but with the hydrogen on the oxygen replaced by a hydrocarbon group).

Amide (carboxamide group)
R—CONH₂ (primary) or R—CONHR (secondary) or R—CONR₂ (tertiary). Not to be confused with amines, which don't have that oxygen double bonded to the carbon that nitrogen is attached to.

What's that? You want more? Fine. Next time, I'll post some more functional groups for you.

The Next Chapter is About Organic Molecules and Functional Groups

I am so excited for this. With the huge hiatus between the "Acid Strength" post and the "Aspirin" post, I hadn't bothered to look at what the next chapter of my text book had in store for this project. Would I skip the chapter? Go over some of it and move on as quickly as possible? Functional groups are one of my favorite things in the whole world and I can hardly wait to write about them. All chemistry is interesting, but functional groups are what truly fascinate me and I hope to study them when I go back to school. But even before then, I'll be doing a little reading up on them while writing posts here. In the meantime, have a little patience for me. I really am back and giving this project my attention, but I want this next post to be my best one so far. I love functional groups.

Aspirin

It's been a while and I need to get a feel for this. I'm also not done with the chapter on acids and bases. So I thought it would be cool to show you aspirin. My book does it, although there's some other stuff before it that I sort of went over, kind of, more or less, already. This will also be an opportunity for me to try to use a program other than MS Paint to render something. We'll see how it works...

Well, that seems to work. I had to make it on this ChemSketch program and then copy the thing into MS Paint in order to upload it here, but whatever. Much easier than drawing everything from scratch.

Now let's make sense of this. There are a few functional groups here, and we'll learn all about functional groups some day. Won't that be exciting? The one we're interested in now is the carboxylic acid group. One of the carbons is attached to an oxygen by a double bond and a second oxygen by a single bond, with that second oxygen itself being attached to hydrogen. This arrangement of atoms makes it easy for a certain reaction to occur. That reaction is a Brønsted-Lowry acid-base reaction. Here's a mechanism...
Okay, those arrows look terrible. I'll have to work on that next time. Anyway, that's the same as the mechanism for the Brønsted-Lowry acid-base reaction I covered in my last post. This reaction can happen in the human body, assuming aspirin gets into the body in the first place. Aspirin is not found in nature. I've sometimes heard it said that aspirin was found in willow bark, but that's incorrect. Willow bark contains salicylic acid. Aspirin is acetylsalicylic acid. The stuff in willow bark would look just like aspirin, but the group attached to that oxygen on top would be replaced with a hydrogen. Aspirin became the analgesic of choice, as opposed to salicylic acid or another derivative of it, because it lacked the irritating side effects of those drugs.

So the reaction can and does occur in the human body. When this happens, the conjugate base, or acetylsalicylate, is formed. The conjugate base is ionic (with a negative charge on the oxygen that gave up hydrogen) and cannot cross cell membranes. Fortunately, this isn't a huge problem because, like other Brønsted-Lowry acid-base reactions, the reaction that turns aspirin into its conjugate base is reversible. My textbook notes that it's the acid that is present in the stomach, and the base that is present in the intestines. This should be pretty intuitive. When conditions are highly acidic, aspirin, which is only a weak acid, even when it does protonate something, will get protonated right away by the acid around it. In basic conditions, there's not much acid around to do the reverse reaction and what is there might be even weaker than aspirin as an acid anyway.

So yeah, aspirin.

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.