The Chemical Bond
Chemical Thermodynamics and the Equilibrium State
Acids and Bases
|Let us revisit a few favorite examples to discuss the underlying chemical thermodynamics governing the relative strength of acids. Why is ethanol a weaker acid than acetic acid? Why is acetic acid a weaker acid than trifluoracetic acid? When each of these acids disassociates, it forms a negatively charged conjugate base. Characteristics of the conjugate base which allow the negatively charged species to exist at a lower energy will reduce the enthalpy of the right side of the equation shifting the equilibrium towards disassociation. The internal energy differences among ethoxide, acetate, and trifluoroacetate underly the differences in the value of the respective acid constants. Acetic acid is a stronger acid than ethanol because resonant acetate holds its negative charge at lower energy than ethoxide. Trifluoroacetic acid is a stronger acid than acetic acid because the electronegative fluorines of trifluoroacetate stabilize the negative charge of the species by induction.|
Similar reasoning, applies to the oxygen acids, the strength of which can be roughly estimated by the value of m in the variant empirical formula XOm(OH)n, as compared to the hydrogen number, n. The more oxygens, the more diffuse the negative charge on the anion. Sulfuric acid (H2SO4 rearranged to SO2(OH)2)is stronger than carbonic acid (H2CO3 rearranged to CO(OH)2) for example.
As a final example, HCl, HBr, and HI are strong, while HF is weak because the small anion radius of F- crowds the negative charge into a small space. A small metal sphere has a lower capacitance than a large metal sphere. The same amount of charge exists at a higher voltage on the small sphere. Likewise, we might say that fluoride ion has lower capacitance than chloride ion. Fluoride is higher energy than chloride ion, so hydrogen fluoride is the weaker acid.