Organic Chemistry Acids and Bases #1

Put the following organic compounds in the order of decreasing acidity:

A.  1, 2, 4, 3
B.  3, 1, 2, 4
C.  2, 3, 4, 1
D.  3, 1, 4, 2

Answer: D


SHORT-AND-SWEET:
Two of the compounds are carboxylic acids (1 is propionic acid, 3 is trichloroacetic acid), and two are alcohols (n-hexanol and phenol).  Clearly, one of the carboxylic ACIDS (hint, hint) will be the most acidic of the four compounds (which means that we can eliminate answer C).

When an acid gives away its proton, the result is a conjugate base, which is an anion.  Nature likes things that are stable, which is NOT charged molecules.  However, if negative charge can be somehow "stabilized" nature likes that.  What will stabilize an anion and its negative charge?

1.  The atom carrying the charge.  The more electronegative (e.g. halogen) and larger (more space for negative charge to spread) the element, more stable the charge.  Also more s character to a hybridized orbital means electrons hanging out closer to the nucleus, which confers stability.
2.  The neighborhood.  If there are electronegative elements in the vicinity or several resonant forms, the charge will be spread over the entire molecule, which makes it more stable.

Trichloroacetic acid has not one, but three chlorine atoms, which will help carry the negative charge, and is therefore more acidic than the propionic acid.  What about the alcohols?  After deprotonation, phenoxide will be able to spread the negative charge over the entire ring, which will stabilize it, making phenol more acidic then n-hexanol (answer D).



THE WHOLE STORY:
Most organic chemistry reactions involve acid-base chemistry, which is why a good understanding of this topic will be your secret weapon for doing well on MCAT organic chemistry questions.

As if the topic were not confusing enough, someone had to come up with THREE different definitions of acids and bases: Bronsted-Lowry, Lewis, and Arrhenius.  Before we get to the actual definition, we would like to point out that for the purposes of MCAT you can accomplish a lot with solid understanding of just Bronsted-Lowry acids and bases.  Lewis comes in handy in certain situations, but Arrhenius you can completely discard because it is useless (sorry, Arrhenius).

Like we said, the main definition to take away is Bronsted-Lowry's, which defines an acid as molecule that donates a proton (H+) to a base, and a base as a molecule which accepts the proton.  After the acid donates its proton, it becomes the conjugate base, and similarly, after a base accepts a proton, it becomes the conjugate acid.

Lewis acids and bases are defined based on electron transfer.  Lewis acids accept electrons, and due to their "love" of electrons are also called electrophiles.  Lewis bases donate electrons, and are called nucleophile.  Many organic reactions are essentially interactions between electrophiles and nucleophiles, and an understanding of Lewis acids and bases comes in handy there.

Arrhenius defines acids and bases based on their dissociation in water and whether they produce hydronium ion, H3O+ (acids), or hydroxide ions, OH- (bases).  Now, forget the last sentence.  Moving on to stuff that you will actually use.

REMEMBER:  Most things in chemistry are relative.  

What does that mean?  It is similar to how things actually work in the real world.  Some of you (and us) were athletes in high school.  Remember being on the varsity tennis (or any other sport) team -- you rocked it!  But let's say that we sent you off to the pro tour, where you'd have to face Serena Williams or Roger Federer.  Compared to theirs, your tennis skills (and ours, too) would seem....well....sadly, not as amazing!

Similarly, in chemistry if you have reagent X, this reagent will act one way in the presence of reagent Y, and might do something completely different in the presence of the third reagent Z.

Same thing with acidity.  How acidic a particular compound will be depends on which molecule it is interacting with.  Therefore, in the right environment more or less every compound can act both as an acid and a base.

However, if you compare all of these compounds to water you can calculate their acidity.  This is expressed as acid dissociation constant Ka or as a negative log of Ka, which is pKa.

REMEMBER:  The strongest acids will have the highest dissociation constants Ka and the lowest (even negative) pKa.

From the expressions above, you can see that when discussing relative acidity we talk about Bronsted-Lowry acids.  In this sense the acidity of a molecule tells you how easily a molecule will get rid of its proton.

The result of the proton dissociation is a negatively charged conjugate base.  One of the big concepts in chemistry is that charge is annoying, so nature will tolerate it only if it is stabilized in some way.  This will apply to the conjugate base as well.

REMEMBER:  The stability of the acid's conjugate base will determine how strong the acid is.  The more stable the conjugate base, the stronger the acid.

What determines stability of conjugate base (or any anion for that matter)?

1.  The atom carrying the negative charge.
- Electronegativity:  there are some atoms which just loooooove electrons.  They are considered very electronegative, and the most electronegative of them you will find as you go to the right and up on the periodic table (fluorine is the most electronegative).  The more electronegative the atom carrying negative charge is, the more stability it confers to the whole molecule.  For example, a negatively charged oxygen will be more stable than if it were on a nitrogen or carbon.

- Size:  the larger the atom (which happens as you go down and to the right in the periodic table), the more stable the negative charge, because the charge is delocalized over a larger space.

- Hybridization state:  hybrid orbitals that have more s character are closer to the positively-charged nucleus, which stabilizes electrons.  This means electrons in orbitals with more s character make anion more stable.  sp orbitals (50% s character) are more stable than sp2 (33% s character) which are more stable than sp3 (25% s character).

2.  The neighborhood.
- Electronegative neighbors:  if electronegative atoms (such as halogens) live close to the atom carrying negative charge they will pull some of the electron density away from that atom (inductive effect), spreading the charge over a larger area, which would stabilize the conjugate base.

- Resonance effects:  resonance is the delocalization of electrons in a molecule such that the bonding cannot be expressed by one single Lewis formula.  This is one of the main effects that make carboxylic acids so acidic compared to the other organic compounds.


........Now, back to our question.  Because the question asks you to put the above molecules in order of decreasing acidity, we suggest you start by identifying the most acidic molecule.

Two of the molecules listed are in fact carboxylic ACIDS (propionic and trichloroacetic acid), which are the most acidic of organic compounds (though still weak acids compared to some of the inorganic acids).  Among the four compounds the most acidic one will probably be one of these two, and our correct answer would start with either 1 or 3 (which eliminates answer C).

How do you figure out which one is the most acidic?  Take a better look at the two acids, and ask yourself which one of them will have a more stable conjugate base.  For both of them the negative charge will rest on oxygen and both will have a stabilizing resonance effect.  However, notice the three chlorine atoms in the trichloroacetic acid.  Chlorine, a halogen, is very electronegative, and will "share" the negative charge with oxygen, stabilizing the molecule.  In fact, the pKa of this acid is 0.6, which means it is more than 10,000 times more acidic than the propionic acid, whose pKa is close to 5 (which eliminates answer A).

We are left now with the two alcohols, n-hexanol and phenol.  You probably know that alcohols and their hydroxyl proton (-OH) are not very acidic.  But how do these two compare to each other in this aspect?

What happens when n-hexanol donates its proton?  The molecule is left with a negative charge on the oxygen, which is alright.  Is there any help from its neighboring atoms?  Not really -- carbon atoms are not particularly eager to help carry the negative charge.

What about phenol?  Upon losing the proton, the negative charge is still on the oxygen, BUT the rest of the ring will help out.  The charge will delocalize over the ring through resonance stabilization, which will make the phenoxide anion more stable, therefore making the phenol (pKa 10) a better acid than n-hexanol (pKa 16).

The correct answer is D:  trichloroacetic acid (3) > propionic acid (1) > phenol (4) > n-hexanol (2).


BIG PICTURE:

1.  Given the appropriate environment, every compound can act as an acid or as a base.  Ka and pKa tell you how acidic a compound is compared to water.

2.  A stable conjugate base means strong acid.  (And vice versa - stable conjugate acid means a strong base.)

3.  Which atom carries a negative charge and who its neighbors are determine anion stability.  Electronegative, large atom with optimal hybridization state (more s character = better), that has neighbors who are equally electron-loving equals stable anion.


~The MCAT POD Team~