Acid Strength
Acid Strength
In the previous exercise, I circled the protons that needed to be removed. But, how did I know which ones to remove? I needed to identify which protons were the most acidic and were most likely to be pulled off by a base.
In organic chemistry, you will be faced with the question on which proton in a molecule is the most acidic. You might even be asked to compare several compounds and decide which is the strongest acid. I hated these questions on tests when I took organic chemistry. I would usually just guess and move on. Somehow, I thought I probably had to memorize some magical list somewhere that ranked acid strengths. But, I never saw that list and probably would not have memorized it even if I did see it. The good news is that you do not need to memorize any list. The way to compare acid strengths is much simpler than that and is quite easy to understand.
Whenever we want to know how strong an acid is, we look at its conjugate base. Another way to say this is, we need to look at what the acid becomes after its proton is removed.
Then, we look at the conjugate base, A-, and determine how happy it is. A- wants to be made if it is really happy. If it is happy and wants to be made, then H-A is a good acid that would gladly give away its H+ proton so it can make A-. A- does not want to be made if it is not happy. Then, H-A would be a weak or poor acid because it would tightly hold onto its H+ proton so it would not have to make A-.
So, how do we tell if A- is happy and stable or not? We need to remember that nature hates localized charges. Localized charges are those that are stuck in one spot. In general, the more we can spread out the negative charge, the better. Also, the more the negative charge is on atoms that like negative charges, the better.
There are four things we look for.
1. What type of atom is the negative charge on?
2. Can it be stabilized with resonance?
3. Does the inductive effect help or hurt it?
4. What type of orbital is the lone pair of electrons in?
Rule 1. Look at the atom that has the negative charge.
First, we need to know where several elements are in the periodic table. We learned these in Chapter 1, but let’s make sure we remember them.
2. Fill in the missing elements in the periodic table below.
Consider this problem. Which of the following compounds is the most acidic?
Whenever we solve a problem like this, we ALWAYS look at the conjugate bases. Let’s draw those.
The real question is which of these negative species is happier? We have N- versus O-. Remembering LiBe BCNOFNe, I see that N and O are in the same row of the periodic table. Between these two, which likes a negative charge on it better? Well, fluorine is the most electronegative element. So, the closer an element is to fluorine, the more electronegative it is and the more it wants the negative charge. With O being to the right of N, we know that O- is happier or more stable than N-.
Important point
If we are comparing two atoms in the same horizontal row of the periodic table, the element farthest to the right likes the negative charge better. Most of the time in organic chemistry, you will be comparing C, N, and O, so know where those are.
3. In each pair below, circle the compound that is the MOST acidic.
R-NH2 vs. R-CH3
H2O vs. NH3
R-CH3 vs. R-OH
Consider this problem. Which of the following compounds is the most acidic?
Once again, we ALWAYS look at the conjugate bases.
Here is where things get a little tricky. If I compare the electronegativities of I and Cl, I see that Cl is closer to fluorine, so it is more electronegative. I might then think, incorrectly, that Cl- is happier than I-. Why is that wrong? Another trend comes into play that ends up being more important than electronegativity. This trend is size. We previously learned that as we go down the periodic table, atoms get larger—because there are more electron shells for the atom. When we compare the sizes of I- and Cl-, we see that I- is quite a bit larger.
Being larger means that the negative charge can spread out and smear over a much larger area. Remember, nature hates localized charges. It is much better the more the negative charge spreads out. Cl- is unhappier than I- because its negative charge is confined to a smaller space.
Important Point
If we are comparing two atoms in the same vertical column of the periodic table, the lower one is larger and likes the negative charge better. This may seem counterintuitive at first because the larger one is usually less electronegative, but size trumps electronegativity in this case.
4. In each pair below, circle the compound that is the MOST acidic.
HBr vs. HI
CH3-SH vs. CH3-OH
HF vs. HCl
Rule 2. Can it be stabilized with resonance?
Rule 1 showed us that spreading out negative charge over a larger atom is very beneficial. Spreading out the charge trumps electronegativity arguments we can make. It would be even better if we could spread out a negative charge over two or more atoms! This is sometimes possible if we can make resonance forms.
Consider ethanol and acetic acid. Acetic acid is about 100 billion times more acidic than ethanol. Why is this?
Whenever we compare acidity, we always, always, ALWAYS look at the conjugate bases and compare them.
So, which of the negative charges is happier or more stable? First, we compare what atoms the negative charge is on. We see that they are both negative oxygen atoms. This is a tie and cannot explain the big difference in acidity of the two compounds. So, then what is it?
The answer is resonance. We can draw a good, valid resonance form for acetate. We cannot do this for ethoxide. The acetate negative charge can be spread over two oxygen atoms while the ethoxide negative charge must remain solely on one oxygen atom. Therefore, it is a much more localized charge. And nature hates localized charges.
Resonance also explains why sulfuric acid is such a strong acid.
Three resonance forms can be drawn for the conjugate base of sulfuric acid. The negative charge can be spread over three electronegative oxygen atoms. This makes this a very happy conjugate base and sulfuric acid a strong acid.
Important Point
If resonance forms can be drawn to spread out a charge over several atoms, it makes it much happier.
5. In each pair below, circle the compound that is the MOST acidic.
Rule 3. Does the inductive effect help or hurt it?
The inductive effect, or induction, is changing the electron density through the sigma bond framework.
Consider this problem. Which of the following compounds is the most acidic?
Whenever we compare acidity, we ALWAYS look at the conjugate bases and compare them.
So, we begin to compare them. They both have negative oxygen atoms, so that is a tie. Neither one has a resonance form to share the negative charge, so they are still tied. So, what can make a difference? One has chlorine atoms attached to the alpha carbon and one has hydrogen atoms attached. We know that chlorine is more electronegative than hydrogen. We can think of the chlorine atoms as vacuum cleaners for negative electrons. They begin sucking the negative electrons through the sigma bond framework and help spread out the negative charge. This does not happen near as much with the hydrogen atoms since hydrogen atoms are not as electronegative. Notice, the electrons are being sucked through the sigma bonds.
It is better the more a charge is spread over a molecule. The chlorine atoms take part of the negative charge upon themselves spreading the charge more.
It is better the closer the electronegative atoms are to the negative charge.
is much happier than
because the Cl atoms are now one bond farther away and cannot suck the electrons from the negative oxygen as well.
Electronegative atoms like F, Cl, Br help stabilize a negative charge through induction. Groups that donate electron density would hurt a negative charge. Silicon-based groups or carbon-based groups (alkyl groups) tend to donate electron density and destabilize negative charges.
Consider this problem. Which of the following compounds is the most acidic?
Whenever we compare acidity, we ALWAYS look at the conjugate bases and compare them.
The negative charge is on the same type of atom. This is a tie. There are no resonance forms for either one. How does the inductive effect influence it?
The CH3 groups are not electronegative. In fact, alkyl groups like CH3 are a little electron donating. They push electron density into the molecule. You could also say that instead of sucking electron density out of the molecule like a halogen does, an alkyl group “blows” electron density into the molecule.
So, which of the protonated compounds is most likely to give up its H+ to make the conjugate base? Definitely not the one that makes something very unhappy. It will hold onto its H+ as much as it can in order to avoid making something so unhappy. It is the weaker acid.
Important Point
Electronegative groups help stabilize negative charges and electron donating groups (alkyls) hurt negative charges.
6. In each pair below, circle the compound that is the MOST acidic.
Rule 4. What type of orbital is the lone pair of electrons in?
Finally, if the negative charge is on carbon atoms, you need to look at what type of hybrid orbital the lone pair of electrons is in. Carbon atoms in organic chemistry are sp, sp2, or sp3 hybridized. Here they are with a blue (+) drawn where the nucleus of the atom is located. These hybrid orbitals are not all the same shape and size.
Round atomic s orbitals hold their electrons close to its nucleus. Dumbbell shaped atomic p orbitals hold their electrons above and below the nucleus. The sp hybrid orbitals are made from mixing one s and one p atomic orbital. They have 50% “s” character. The sp2 hybrid orbitals are made from mixing one s and two p atomic orbitals. They have 33.3% “s” character. The sp3 hybrid orbitals are made by mixing one s and three p atomic orbitals. They have 25% “s” character. The more “s” character a hybrid orbital has the more it holds its electrons close to the nucleus. So, sp hybrid orbitals hold their electrons closer to its positive nucleus. Sp2 hybrid orbitals hold their electrons a little farther away from the positive nucleus and sp3 hybrid orbitals hold their electrons the farthest away from the positive nucleus. Lone pairs of negative electrons prefer to be in sp hybrid orbitals because they are then held closer to the positive nucleus.
Consider this problem. Which of the following compounds is the most acidic?
Whenever we compare acidity, we ALWAYS look at the conjugate bases and compare them.
The negative charge is on a carbon atom for each one. This is a tie.
There are no resonance forms for any of them.
The inductive effect might have a small effect.
The largest effect is the hybrid orbital the lone pair of electrons is in.
So,
is the most acidic of these compounds because it makes the happiest conjugate base. In fact, it is 10000000000000000000 times more acidic than a proton on an alkene!
Important Point
Protons on sp carbons are more acidic than on sp2 or sp3 carbon atoms. This usually means protons on triple bonds, C≡C-H, are more acidic than on double bonds, C=C-H, and single bonds C-C-H.
7. Label the circled protons from most to least acidic (1=most, 3=least).
8. Methyl mercaptan (CH3SH) is a stronger acid than methanol (CH3OH). Explain.
Answers
2.
3.
4.
5.
6.
7.
8. Methyl mercaptan (CH3SH) is a stronger acid than methanol (CH3OH). Explain.
We look at the conjugate bases and see that CH3S- is more stable and happier with its negative charge than CH3O- since S is larger than O and can spread out the negative charge more. Therefore, methyl mercaptan is more likely to give up its proton making it the stronger acid.