Acid Base Reactions
By James Ashenhurst
Walkthrough of Acid-Base Reactions (3) – Acidity Trends
Last updated: March 21st, 2019
Let’s review what’s been talked about so far in this series on acid-base reactions:
An acid base reaction involves the donation of a proton (H+) from an acid to a base.
- The species which loses H+ is the acid
- The species which gains H+ is the base
- The conjugate base is what becomes of the acid after it loses H+
- The conjugate acid is what becomes of the base after it gains H+
All else being equal, charged species are more unstable than neutral species. Since an acid is becoming more negative upon loss of a proton, the stability of the new lone pair on the conjugate base is a key factor in determining how favorable the reaction will be.
In other words:
- any factor which stabilizes the conjugate base will increase the acidity.
- any factor which destabilizes the conjugate base will decrease the acidity.
For our purposes, a roughly equivalent word for “stability” is “basicity”. The more “unstable” the pair of electrons on a species is, the more basic it is. Stabilizing a lone pair lowers the basicity; destabilizing the lone pair increases basicity.
Last time we went through the effect of electronegativity on acidity and basicity. But there are at least six more factors that can act to stabilize negative charge in organic chemistry. So let’s go through each of these in turn, and explicitly tie together the relationship between the stability of negative charge and acidity/basicity.
1. The less charge the better.
All else being equal, the lower the charge density, the more stable a species is. So stability increases as we go from O(2-) to HO(-) to H2O. This is demonstrated in the acidity/basicity trends of water and related species, shown here. H3O(+) is the most acidic (most stable conjugate base); HO(-) is the least acidic (least stable conjugate base).
High charge densities tend to be less stable than low charge densities. The more “spread out” a charge can be, the more stable it will be. So as we go down the periodic table from F(-) to I(-), notice that the magnitude of the negative charge doesn’t change, but the volume that it occupies does. Iodide ion, being considerably larger than fluoride ion (206 pm vs. 119 pm) is more stable, which means that H-I is a stronger acid than H-F.
3. Electron withdrawing groups
The last post described how stability of a negative charge increases with increasing electronegativity of the atom. Adding electron withdrawing groups to an atom can have the a similar effect to that of increasing electronegativity. For example, in the molecule below, the stabilization of the negative charge increases with every hydrogen that is substituted for chorine.
Resonance is another means by which negative charge can be dispersed in a molecule. If the conjugate base has a charge which can interact with adjacent double bonds or p orbitals, its stability will increase. This leads to increased acidity of the conjugate acid.
Increasing the “s-character” of an orbital has a similar effect as that of increasing electronegativity and adding electron withdrawing groups; it brings the negative charge closer to the positively charged nucleus, which , according to Coulomb’s Law, is a favorable interaction. So as we go from sp3 (alkane) to sp2 (alkene) to sp (alkyne) hybridization, the stability of the negative charge increases. So alkynes are remarkably acidic compared to alkanes.
A special case is that of aromaticity – usually dealt with in the first few weeks of second-semester organic chemistry – which is a special type of stability exhibited by some conjugated molecules. Without going into details, I’ll just say that if the conjugate base is aromatic, this will lead to a tremendous increase in acidity relative to similar (but non-aromatic) molecules.
I know this has been yet another long post but hopefully it at least ties together the issue of charge stability and acidity/basicity. Whenever the issue of stability comes up, so should the issue of acidity and basicity. These items are woven together.
So far I’ve just talked about trends. But what do you do when you need to compare the acidity of two species that aren’t related by a trend, like, propanol and H2S? How do you proceed? In this case we are going to resort to the results of experimental measurements, and in the next post I’ll start to discuss how we do that.
Next Post: Walkthrough of Acid-Base Reactions (4) – pKa