Acids and Bases

Acid Dissociation Constant

The acid dissociation constant, $ K_a $ , is a quantitative measure of the strength of an acid in solution. It represents the equilibrium constant for the dissociation of an acid, HA, into its conjugate base, $ A^- $ , and a proton, $ H^+ $ .

$$ HA(aq) \rightleftharpoons H^+(aq) + A^-(aq) $$

The expression for $ K_a $ is:

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

where $ [H^+]] $ , $ [A^-]] $ , and $ [HA]] $ represent the equilibrium concentrations of the hydrogen ion, conjugate base, and the undissociated acid, respectively.

A larger $ K_a $ value indicates a stronger acid, meaning it dissociates more readily. A smaller $ K_a $ value indicates a weaker acid. $ pK_a $ , which is defined as $ pK_a = -\log_{10} K_a $ , is often used instead of $ K_a $ because it is easier to manage numerically. A smaller $ pK_a $ value indicates a stronger acid.

Weak Acids and Bases These notes will cover the characteristics of weak acids and bases and how to calculate their pH using the $ K_a $ and $ K_b $ values. This will include ICE tables and approximations.

Polyprotic Acids This will focus on acids that can donate more than one proton and how to calculate the different $ K_a $ values and their influence on pH.

Titration Curves This will explain how to interpret titration curves involving weak acids and relate them to $ K_a $ and $ pK_a $ .

Buffers This will discuss buffer solutions, their capacity, and how the $ K_a $ of the weak acid component is relevant to their effectiveness. It will cover the Henderson-Hasselbalch equation.

Relationship between $ K_a $ and $ K_b $ This section will explain the mathematical relationship between the acid dissociation constant ( $ K_a $ ) and the base dissociation constant ( $ K_b $ ) for a conjugate acid-base pair. This involves the ion-product constant for water, $ K_w $ .