Knowing how to determine the relative strength of organic acids and bases is the key to understanding, (not just memorising) the organic reactions that make up the rest of this organic course.
You already know that a strong acid is a molecule that will readily donate a proton, and a strong base will gain one. You should already know some strong inorganic acids and bases, but you may not know many organic ones. That’s okay – this isn’t a place to memorise a great big list of compounds, instead you just need to know three rules, (and the reasons behind them) and remember one important thing: it all comes down the the stability of the conjugate acid or base.
The key to a strong acid/base is that the conjugate base/acid must be stable in solution. A strong acid has a weak conjugate base. A strong base has a weak conjugate acid. Confused? Prove it to yourself using some inorganic acids or bases. This means that when you’re looking at a list of acids or bases, sometimes it is easiest to look at their conjugate bases or acids and determine their stability and relative strength.
There are three things that will influence the conjugate base/acid’s stability:
Let’s start by looking at how the acidity is affected, then deal with bases later…
Electronegativity is the measure of how strongly an atom attracts electrons. It increases across a period and decreases down a group – i.e. fluorine is the most electronegative element, and francium is the least. Highly electronegative atoms are happy to hold a negative charge and exist as their negative ions.
There are two parts to how electronegativity dictates acidity:
a) the atom bearing the charge in the conjugate ion:
b) inductive effect
The atom bearing the charge in the conjugate ion is the more significant of the two effects. If the charge-bearing atom is highly electronegative, (fluorine, chlorine etc) the anion (conjugate base) will be more stable than if you had an electropositive atom in its place.
The inductive effect is a little more subtle, and involves having electronegative atoms that are not directly attached to the acidic proton. When you lose that proton, an electronegative atom further down the molecule, (provided it isn’t too far, the inductive effect falls away very quickly with distance) will help stabilise the negative charge.
The acidity of a compound changes with the hybridisation of the carbon. Ethyne (sp) is more acidic than ethene (sp2) or ethane (sp3), for example. Orbitals with higher s character, (a higher proportion of s orbital to p orbital; sp orbitals have 50% s character, sp2 33% and sp3 25%) are more stable, as the electrons in s orbitals are close to the nucleus. It follows that in an anion, electrons would rather be in a high s character orbital, as they will be closer to the nucleus and therefore more stable.
This is actually an extension of electronegativity, as an sp carbon is more electronegative than an sp3 carbon.
3. Resonance stabilisation:
An anion that has a resonance form is going to be far more stable than one where the charge is localised to one atom. Spreading out the charge through resonance will increase acidity. We will talk more about resonance stabilisation in a few weeks.
To summarise, having electronegative atoms, carbons with high s character and resonance structures for the anion will all make a compound more acidic. Of course, they will also make something a weaker base. A strong base will have electropositive atoms, (to hold the positive charge), low s character and will not be able to form resonance structures.
For practice, place the following in order of decreasing acidity:
CH3CH2CH2OH, CH3CH2CO2H, CH3CHClCO2H
Some links that you might find useful: