In our discussion of acid-base equilibria, we saw how to calculate the equilibrium constant for any reaction involving two acids of known acid strength. By comparing differences in acidity within carefully selected groups of structurally related compounds, it is possible to gain insight into the relationship between molecular structure and chemical reactivity. In this topic we will make correlations between changes in pKa values and various atomic, molecular, and experimental variables:
- electronegativity
- bond strength
- atomic size
- resonance
- solvents
We'll start by looking at HF and its relatives, H2O, NH3, and CH4.
Charge Stability and Electronegativity
Equation 1 describes the general reaction that we want to consider.
Relative Acidities of the Hydrogen HalidesIn aqueous solution the relative acidities of the hydrogen halides is HI > HBr > HCl > HF.
It should be clear that the balance of factors which contribute to trends in acidity as you go across a row in the periodic table is different than the balance that determines the relative acidities as you go down a group. The bottom line is that you have to be flexible. Although there are no hard and fast rules to go by, the trends can generally be rationalized in terms of the balance between two factors or three factors. The last factor we will examine in this topic is resonance.
Charge Stability and Resonance
Figure 1 compares the dissociation reactions of methanol and phenol. As you can see, phenol is approximately 1,000,000 times more acidic than methanol. Clearly differences in bond strengths or electronegativity are not responsible for this difference in acidity since an O-H bond is broken in both cases and the negative charge resides on an oxygen atom in both conjugate bases.
Figure 1
Resonance to the Rescue
In our discussion of resonance theory we saw that when an orbital containing a lone pair of electrons can overlap with a p orbital of an adjacent multiple bond as shown in Figure 2, the energy of the system decreases.
Figure 2
If the Shoe Fits..
As Figure 3 indicates, the phenolate ion, C6H5O:-, meets this structural requirement.
Figure 3
Lending a Helping Hand
Notice that Figure implies that the orbital overlap leads to changes in the hybridization of the oxygen atom. This change decreases the separation between the orbitals on the oxygen and the adjacent carbon and increases their overlap.
This orbital overlap leads to charge delocalization as indicated by the resonance structures shown in Figure 4.
Figure 4
Spread 'em Out
Now consider the difference in pK
a values of methanol and acetic acid. Figure 5 provides a comparison.
Figure 5
Here We Go Again
As you can see, the O-H hydrogen in acetic acid is 11 orders of magnitude more acidic than that in methanol. Conversely, the methoxide ion, CH3O:-, is 11 orders of magnitude more basic than the acetate ion. Clearly the carbonyl group affects the acidity of the adjacent O-H group dramatically. We attribute this increase in acidity to the stabilization that the double bond affords the negative charge as it develops in the conjugate base of acetic acid. The interaction between a lone pair of electrons on the oxygen and the adjacent carbonyl group is shown in Figure 6.
Figure 6
Spread 'em Again
Remember, the formalism shown in Figure 6 is our way of indicating that the negative charge in the acetate ion can experience the nuclear attraction of both oxygen atoms. This increase in Coulombic attraction reduces the potential energy of the acetate ion relative to that of the methoxide ion where the charge is localized on one oxygen atom.
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