When we discuss acidity in organic, it is important to remember that this is a relative value. For instance, we might be comparing the acid strength of something that is not normally associated as being acidic. An example of this might be comparing the acidity of two separate alcohols. So, what would make two alcohols, both being the same functional group, vary in acidity?

First, make sure you can identify which compound is acting as an acid, here a proton donor, and which is the conjugate base. Once the conjugate base is identified, you can clearly see that there is a center of electron density (the negative charge). So, how can this charge be stabilized? Well, if you could remove the charge somehow, that would certainly stabilize (which means lower in energy) the molecule. This is where inductive effects and resonance effects come into play.

Inductive Effects

Atoms have different affinities for electrons or different electronegativities. This is particularly important when we discuss dipole moments in molecules. For instance, H-Br has a dipole moment that points towards the bromine since it is more electronegative than hydrogen. Likewise, water has a dipole that points towards the oxygen. The reason for this dipole is the difference in the electronegativity between the two adjacent atoms.

In light of this phenomena that we have already seen with dipoles, the theory of inductive effects can be discussed. Inductive effects are the pull/push of electrons through the sigma bond framework of a molecule. Remember that each bond to an atom contains two electrons and those electrons are shared. In the same way that a dipole exists between unequal sharing, inductive effects pull or push electron density. This allows us to explain several different phenomena. Let's take for example a set of carboxylic acids.

In the case of these two acids, the top being acetic acid and the bottom 2-fluoroacetic acid, the pKa values are rather different. Fluoroacetic acid has a pKa of about 2.6 whereas acetic acid has a pKa of 4.76. This is a rather large difference. So, the question is why the difference? Inductive effects! Take a look below.

We can see from the picture that the flourine is pulling electron density away from the center of negative charge on the conjugate base. Anytime you can move charge away from one spot and spread it around multiple spots, you cause a stabilizing effect. Therefore, the electron withdrawing potential of the fluorine pulls electron density away from the oxygen with the negative charge and hence, stabilizes the conjugate base. Finally, if one stabilizes the conjugate, the acid becomes stronger. This is the reason for the difference in pKa for the two compounds.

Inductive effects are both distance dependent and additive. This means that in order for inductive effects to be considered, they must be relatively close to the position they are affecting. The reason for the distance requirement is that this is a through bond (sigma bonds!) effect. Therefore, the further the electronegative atom is from the acidic proton, the less it can pull electron density.

As you can see from the picture above, there is only a slight difference in the pKa values of butanoic acid and 4-chlorobutanoic acid. This is primarily due to the fact the chlorine atom is too far away to have any real and significant effect on the acidity of the carboxylic acid. Just as a side note, butanoic (also called butyric) acid smells awful and is a component of human vomit and rancid butter. Hungry?

Resonance Effects

Resonance effects are the other component to determining the acidity of compounds. These effects are stronger than inductive effects and can act over longer distances. The reason for this is that resonance involves the movement of pi electrons! Remember, inductive effects don't actually move any pi bonds. Since a picture is worth a thousand words, take a look below.

As you can see from the picture, both of these alcohols are attached to 6-membered rings but only one of them has pi electrons. The phenol has pi electrons that can be moved around the ring (known as resonance structures). This greatly increases the stability of the anion that would result from deprotonation.

In the above picture, notice that the negative charge on the oxygen of the phenol (called a phenoxide) can be moved all the way around the ring since it is adjacent to pi electrons. Moving the electrons (and hence the negative charge) to other places rather than just the oxygen. Think of a hot potato. If you have a hot potato and you hold it with your hands without some type of glove, chances are high that your hand will get burned. However, if you pass the hot potato to another person quickly, then you are less likely to get burned. Now, if you pass the potato across multiple people then you are even less likely to get burned. This same reasoning works with charges. The more you can delocalize the charge (move it around, resonance) the more stabile the molecule.

At this point, you may be wondering how electronegativity affects the molecule. After all, shouldn't the oxygen be pulling electron density away from the ring and not pushing its electrons into a ring? Isn't oxygen more electronegative than carbon? The answer is that anytime resonance and inductive effects compete against each other, resonance wins. Therefore, the resonance effect on oxygen outweighs the electronegativity. So let's look at an example.

Which is more acidic?

To solve this problem, we need to think about which molecule will stabilize the conjugate base more. Remember, the more stabile the conjugate base, the more acidic. This requires us to know what the conjugate base is and if we see from above, it is the phenoxide. Now that we know what the conjugate base is, we need to look at which molecule will stabilize it more. Both molecules have a ring with pi electrons (aromatic ring), so that cannot be the difference. The substituents are, however, different. The compound on the left has a methoxy group and the compound on the right has a nitro group. The best thing to do at this point is draw everything out.

What we can now easily see it the the methoxy group, with its lone pairs, actually pushes electron density onto the ring. This destabilizes the conjugate base since electrons are being pushed together. On the other hand, the nitro group acts as an electron sink and both inductively and through resonance pulls the electrons from the phenoxide oxygen. See if you can draw the resonance structures that prove this fact! Therefore, since the conjugate base is more stabilized, the compound on the right is more acidic. Again, in this example, the resonance effect of the lone pairs on oxygen outweigh its electronegativity.

A final point to make regarding this subject. Make sure you are comparing apples to apples and not apples to oranges. Take for instance a carboxylic acid and a phenol. In general, carboxylic acids have pKa values around 5 while phenols have pKa values around 10. However, there are many more resonance structures for a phenol vs a carboxylic acid. This would lead some to believe that phenols should be more acidic. This, however, is not the case and carboxylic acids are indeed more acidic as we can see by their pKa values. The comparison between the two is one that you simply cannot make as they are different compounds all together.

Happy studying!