Introduction to Acids and Bases — AP Chemistry
1. Acid-Base Definitions and Conjugate Pairs ★★☆☆☆ ⏱ 4 min
The AP Chemistry exam expects you to know three hierarchical definitions of acids and bases, each with a different scope. The Arrhenius definition, limited to aqueous solutions, defines acids as substances that increase $[\text{H}_3\text{O}^+]$ and bases as substances that increase $[\text{OH}^-]$ when dissolved in water.
The most commonly tested definition on the AP exam is Brønsted-Lowry: acids are proton ($\text{H}^+$) donors, and bases are proton acceptors. This leads directly to the concept of conjugate acid-base pairs: every acid donates a proton to form its conjugate base, and every base accepts a proton to form its conjugate acid. By definition, conjugate pairs differ by exactly one proton. A key rule: the stronger an acid, the weaker its conjugate base, and vice versa.
The most general definition is Lewis: Lewis acids accept an electron pair, and Lewis bases donate an electron pair. This covers reactions without proton transfer, but is less commonly tested in introductory problems.
Exam tip: When asked to identify conjugate pairs on the AP exam, always confirm the two species differ by exactly one proton. Common distractors use pairs differing by two protons, so counting H atoms will eliminate wrong answers quickly.
2. Autoionization of Water and the pH Scale ★★★☆☆ ⏱ 4 min
Water is amphoteric, meaning it can act as either a Brønsted-Lowry acid or base depending on its reaction partner. In pure water, two water molecules undergo reversible autoionization:
2\text{H}_2\text{O}(l) \rightleftharpoons \text{H}_3\text{O}^+(aq) + \text{OH}^-(aq)
The equilibrium constant for this reaction is the ion product of water, $K_w$. Pure liquid water is omitted from the equilibrium expression, giving:
K_w = [\text{H}_3\text{O}^+][\text{OH}^-]
At 25°C, $K_w = 1.0 \times 10^{-14}$. The pH scale simplifies working with small hydronium concentrations, defined as $\text{pH} = -\log[\text{H}_3\text{O}^+]$, and pOH is $\text{pOH} = -\log[\text{OH}^-]$. At 25°C, this gives the key relationship $\text{pH} + \text{pOH} = 14$.
A neutral solution has equal concentrations of hydronium and hydroxide: $[\text{H}_3\text{O}^+] = [\text{OH}^-]$, which only gives $\text{pH} = 7$ at 25°C. Acidic solutions have $\text{pH} < 7$, basic solutions have $\text{pH} > 7$ at 25°C. $K_w$ increases with temperature (autoionization is endothermic), so neutral pH decreases as temperature increases.
Exam tip: Never automatically assume neutral solutions have pH = 7. Always check if the problem gives a non-room temperature or a different $K_w$ value, and apply the definition of neutrality correctly.
3. Acid and Base Dissociation Constants ($K_a$ and $K_b$) ★★★☆☆ ⏱ 4 min
Strong acids and bases dissociate completely in dilute aqueous solution, so no equilibrium constant is needed for introductory calculations. Weak acids and bases only partially dissociate, so we use equilibrium constants to describe their strength.
For a general weak acid $\text{HA}$, dissociation in water is:
\text{HA}(aq) + \text{H}_2\text{O}(l) \rightleftharpoons \text{A}^-(aq) + \text{H}_3\text{O}^+(aq)
The acid dissociation constant $K_a$ is:
K_a = \frac{[\text{H}_3\text{O}^+][\text{A}^-]}{[\text{HA}]}
For a general weak base $\text{B}$, reaction with water is:
\text{B}(aq) + \text{H}_2\text{O}(l) \rightleftharpoons \text{BH}^+(aq) + \text{OH}^-(aq)
K_b = \frac{[\text{BH}^+][\text{OH}^-]}{[\text{B}]}
Larger $K_a$ corresponds to a stronger weak acid, and larger $K_b$ corresponds to a stronger weak base. For any conjugate acid-base pair, the key relationship is:
K_a \times K_b = K_w
This relationship lets you calculate $K_b$ of a conjugate base from $K_a$ of the parent acid, and vice versa, and it only applies to conjugate pairs.
Exam tip: The $K_a \times K_b = K_w$ relationship only applies to conjugate pairs. Never use it to relate an acid and an unrelated base, as this will always give an incorrect result.
4. AP Style Concept Check ★★★☆☆ ⏱ 2 min
Common Pitfalls
Why: Students memorize the 25°C value and forget that $K_w$ changes with temperature, changing $\text{p}K_w$.
Why: Students confuse charge change with proton change, or forget that one proton adds both one H atom and +1 charge.
Why: Students forget that pure liquids have an activity of 1 in equilibrium expressions and are omitted.
Why: Students misremember the definition of the pH scale and forget the negative sign.
Why: Students memorize the relationship but forget the requirement that it only applies to pairs that differ by one proton.
Why: Students learn water as the solvent first, so they forget it can act as either proton donor or acceptor depending on the reaction partner.