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Understanding the Dissociation Constant (Ka)

In the Brønsted–Lowry framework, an acid donates a proton to a solvent, producing a conjugate base and a conjugate acid. When water is the solvent, the equilibrium can be written as:

HA + H₂O <=> H₃O⁺ + A^-

The acid’s strength is quantified by the dissociation constant Ka, defined as the ratio of the concentrations of the products to the reactants (excluding water in aqueous solutions):

Ka = [H₃O⁺][A^-]/[HA]

A larger Ka indicates that the reaction favors the dissociated form, signifying a stronger acid. Conversely, a small Ka reflects a weaker acid.

Why pKa Is More Convenient

Ka values span many orders of magnitude— from as high as 10⁷ for strong acids to as low as 10^-12 for weak acids. To manage this range, chemists use the negative base‑10 logarithm of Ka, called pKa:

pKa = -log₁₀(Ka)

Thus, a strong acid with Ka = 10⁷ has pKa = -7, while a weak acid with Ka = 10^-12 has pKa = 12. The inverse relationship means a lower pKa corresponds to a stronger acid.

Converting pKa Back to Ka

When you encounter a pKa value in a reference table but need the Ka for calculations, simply reverse the logarithm:

Ka = 10^-pKa

For integer pKa values (e.g., –7) the conversion is straightforward. For fractional values (e.g., 7.5), you can use a scientific calculator or look up the result in a table. Remember that the operation involves base‑10 exponentiation, not natural logarithms.

Using this method, you can quickly obtain Ka from any reported pKa, facilitating accurate acid–base computations.