# Ionizable hydrogen atoms for acids

Given that hydrogen atoms that are attached to a compound must be ionizable for the compound to be an acid, what prevents a hydrogen from being ionized? Is it solubility. If the bond is nonviolent, water could not dissolve the compound, and so hydrogen atoms are not ionizable. Is this correct?

Not all hydrogen atoms in a compound are equally acidic. The acidity of a hydrogen at a certain position in a molecule depends on its chemical environment, i.e. the solvent (if the compound is dissolved), which atoms and functional groups it is attached to, and how well the negative charge of the corresponding conjugated base is stabilized. A quantitative measure for the acidity of a proton is its $pK_a$ value which is defined as follows (reference):

$$pK_a=-\mathrm{log}_{10}K_a$$

with $K_a$ being the equilibrium constant of the dissociation reaction in water $\ce{HA + H2O \rightleftharpoons H3O+ + A-}$, in which the hydrogen is abstracted from the compound by heterolytic bond cleavage. $[\ce{H2O}]$ can be assumed to be constant (approx. 55 mol/l) and is therefore included in $K_a$:

$$K_{a}=\frac{[A^{-}][H^{+}]}{[HA]}$$

Let's take formic acid ($\ce{HCOOH}$) as an example. As the name of the compound suggests, it has at least one acidic proton, and if all hydrogens would be equally acidic and dissociate from the molecule, the reaction would look like this:

$$\ce{HCOOH + 2H2O \rightleftharpoons 2H3O+ + COO^{2-}}$$

However, this does not happen in aqueous solution, only one of the protons is acidic:

$$\ce{HCOOH + H2O \rightleftharpoons H3O+ + HCOO-}$$

The reason for this is that both hydrogens are situated at different positions in the structure:

With a $pK_a$ of 3.77, the hydrogen of the carboxyl group is significantly more acidic than the hydrogen bonded to the carbon atom. The typical $pK_a$ of an aldehyde hydrogen (H bonded to the carbon of a carbonyl group) is in the range of 17 (source), so such a position cannot be deprotonated in water.

• And the reason the hydrogen of the carboxyl group is more acidic than the hydrogen directly bonded to the central carbon atom has to do with the electronegativity of carbon (or lack thereof). May 8 '14 at 16:34
• Your argument would be much stronger if you consider aqueous solution in your equations, i.e. $\ce{HA + H2O <=> H3+O + A^{-}}$. May 12 '14 at 14:53
• @Martin Thank you for your suggestion. I have edited my equations accordingly. May 12 '14 at 15:17

Ionization here (in terms of acids) refers to heterolytic bond clevage, not homolytic bond clevage (which would create a hydrogen atom, complete with its one electron). Unfortunately in the Brønsted-Lowry definition of acids, we are dealing with hydrogen protons - i.e. hydrogen atoms stripped of their lone electrons.

Of course we could take a whole hydrogen atom from a molecule, i.e. we could do this with water, and what we'd end up with is a hydrogen atom and a hydroxyl group. Unfortunately the lone hydrogen atom is rather unstable by itself - that's why at standard state conditions you find not the hydrogen radical but instead diatomic hydrogen. But then the bare hydrogen proton is rather unstable too in its own right; it's solvated in water actually (as the hydronium ion; it may even be solved by multiple waters).