How Structure Affects Acid Strength | CHEMISTRY BY BURHAN ISHAQUE

I'll introduce some generalizations that will permit us to connect molecular structure with acidity in related compounds. The main ways in which structure affects acidity in solution depend on:

1. The strength of the bond to the atom from which the proton is lost

2. The electronegativity of the atom from which the proton is lost

3. Electron delocalization in the conjugate base

Bond Strength.  The effect of bond strength is easy to see by comparing the acidities of the hydrogen halides.

In general, bond strength decreases going down a group in the periodic table. As the halogen X becomes larger, the H–X bond becomes longer and weaker and acid strength increases. This is the dominant factor in the series HCl, HBr, HI and also contributes to the relative weakness of HF.

With HF, a second factor concerns the high charge-to-size ratio of F⁻. Other things being equal, processes that give ions in which the electric charge is constrained to a small volume are less favorable than processes in which the charge is more spread out. The strong H–F bond and the high charge-to-size ratio of F⁻ combine to make HF the weakest acid of the hydrogen halides.

Because of the conjugate relationship between acidity and basicity, the strongest acid (HI) has the weakest conjugate base (I⁻), and the weakest acid (HF) has the strongest conjugate base (F⁻).

Electronegativity.  The effect of electronegativity on acidity is evident in the following series involving bonds between hydrogen and the second-row elements C, N, O, and F.

As the atom (A) to which H is bonded becomes more electronegative, the polarization ^δ+H–A^δ− becomes more pronounced and the equilibrium constant K_a for proton transfer increases.

Bond strength is more important than electronegativity when comparing elements in the same group of the periodic table as the pK_a’s for the hydrogen halides show. Fluorine is the most electronegative and iodine the least electronegative of the halogens, but HF is the weakest acid while HI is the strongest. Electronegativity is the more important factor when comparing elements in the same row of the periodic table.

In many acids, the acidic proton is bonded to oxygen. Such compounds can be considered as derivatives of water. Among organic compounds, the ones most closely related to water are alcohols. Most alcohols are somewhat weaker acids than water.

Electronegative atoms in a molecule can affect acidity even when they are not directly bonded to the ionizable proton. Compare ethanol (CH₃CH₂OH) with a related compound in which a CF₃ group replaces the CH₃ group.

We see that the substitution of C–H bonds by C–F increases the acidity of the O–H proton by 4.7 pK_a units, which corresponds to a difference of 10^4.7 in K_a. The simplest explanation for this enhanced acidity is that the electronegative fluorines attract electrons and that this attraction is transmitted through the bonds, increasing the positive character of the O–H proton.

The greater positive character, hence the increased acidity, of the O–H proton of 2,2,2-trifluoroethanol can be seen in the electrostatic potential maps displayed in Figure below.

Electrostatic potential maps of ethanol and 2,2,2-trifluoroethanol. As indicated by the more blue, less green color in the region near the OH proton in 2,2,2-trifluoroethanol, this proton bears a greater degree of positive charge and is more acidic than the OH proton in ethanol. The color scale is the same in both maps.

We can also explain the greater acidity of CF₃CH₂OH relative to CH₃CH₂OH by referring to the equations for their ionization

The conjugate base of 2,2,2-trifluoroethanol, the anion CF₃CH₂O⁻, is stabilized by its three fluorines, which attract electrons from the negatively charged oxygen, dispersing the negative charge. Because of this stabilization, the equilibrium for ionization of CF₃CH₂OH lies farther to the right than that of CH₃CH₂OH.

Structural effects that are transmitted through bonds are called inductive effects. A substituent induces a polarization in the bonds between it and some remote site.

The same kind of inductive effects that make CF₃CH₂OH a stronger acid than CH₃CH₂OH makes the trifluoro derivative of acetic acid more than 4 pK_a units stronger than acetic acid.

Inductive effects depend on the electronegativity of the substituent and the number of bonds between it and the affected site. As the number of bonds between the two units increases, the inductive effect decreases. For instance, the two butanoic acids differ in the placement of chlorines by two bonds and differ by more than 1.5 pK_a units.

Electron Delocalization in the Conjugate Base.  With a pK_a of −1.4, nitric acid is almost completely ionized in water. If we look at the Lewis formula of nitric acid in light of what we have said about inductive effects, we can see why. The N atom in nitric acid is not only electronegative in its own right, but bears a formal charge of +1, which enhances its ability to attract electrons away from the –OH group. But inductive effects are only part of the story. When nitric acid transfers its proton to water, nitrate ion is produced

Nitrate ion is stabilized by electron delocalization, which we can represent in terms of resonance between three equivalent contributing structures:

The negative charge is shared equally by all three oxygens. Stabilization of nitrate ion by electron delocalization increases the equilibrium constant for its formation

Acetic acid, with a pK_a of 4.7, derives its acidity from the electron delocalization of its conjugate base acetate ion. Acetic acid is over 11 pK_a units lower than ethanol

The two equivalent resonance contributors with the negative charge shared equally on both oxygen atoms stabilize this structure.

Electron delocalization in the conjugate base is also responsible for the increased acidity of phenols compared to alcohols.

This electron delocalization is represented by resonance among the various contributing structures:

Organic chemistry involves a good bit of reasoning by analogy and looking for trends. At the beginning of this section we listed three ways that structure can affect acidity. The last two—electronegativity of the atom from which the proton is lost, and electron delocalization in the conjugate base—are both related to the stability of the conjugate base. A useful trend emerges: factors that stabilize the conjugate base increase the acidity of the parent acid.