AP Chemistry: Acids and Bases

Acids and bases are a core unit in AP Chemistry and a frequent source of both easy points and costly mistakes. The exam expects you to move comfortably between models (especially Brønsted-Lowry theory), calculations (pH, $K_a$, $K_b$), and representations (particle diagrams, titration curves). Mastery is less about memorizing isolated formulas and more about knowing which chemical story you are in: strong vs. weak, before vs. after equivalence, buffered vs. unbuffered, and what species actually exist in solution.

The Brønsted-Lowry framework: the unit’s grammar

AP Chemistry leans heavily on Brønsted-Lowry definitions:

• Acid: proton donor
• Base: proton acceptor

This definition is powerful because it explains conjugate acid-base pairs. When an acid donates $H^{+}$, it becomes its conjugate base; when a base accepts $H^{+}$, it becomes its conjugate acid.

Example:

• $HF$ donates $H^{+}$ to become $F^{-}$ (conjugate base)
• $N{H}_3$ accepts $H^{+}$ to become $N{H}_4^{+}$ (conjugate acid)

A practical exam skill is recognizing that “stronger acid” means “weaker conjugate base.” That relationship shows up in reasoning questions, equilibrium direction, and qualitative ranking.

Strong versus weak acids and bases: what “strong” really means

“Strong” does not mean concentrated. It means complete (or essentially complete) dissociation in water.

• Strong acid example: $HCl\to H^{+}+C{l}^{-}$
• Weak acid example: $C{H}_{3}COOH\rightleftharpoons H^{+}+C{H}_{3}CO{O}^{-}$

The same logic applies to bases:

• Strong base: $NaOH\to N{a}^{+}+O{H}^{-}$
• Weak base: $N{H}_3+H_{2}O\rightleftharpoons N{H}_4^{+}+O{H}^{-}$

The key calculation difference

• For strong acids/bases, pH or pOH is driven directly by initial concentration (after stoichiometry, if mixed).
• For weak acids/bases, you must use an equilibrium approach with $K_a$ or $K_b$.

Connecting $K_a$ and $K_b$

For a conjugate acid-base pair: $$K_a{K}_b=K_w$$ with $K_w=1.0\times 10^{-14}$ at 25°C.

This relationship is often the fastest path when given only one constant.

pH, pOH, and what the math is actually saying

At 25°C:

• $pH=-\log [H^{+}]$
• $pOH=-\log [O{H}^{-}]$
• $pH+pOH=14$

AP problems frequently mix equilibrium chemistry with logarithms. Two habits prevent common errors:

1. Track the chemistry first (what reactions occur? what is left after neutralization?).
2. Logarithms come last (convert to pH after you have $[H^{+}]$ or $[O{H}^{-}]$).

Weak acid pH via ICE tables (the standard method)

For a weak acid $HA$: $$HA\rightleftharpoons H^{+}+A^{-}$$ $$K_a=\frac{[H^{+}][A^{-}]}{[HA]}$$

If the initial concentration is $C$ and the change is $x$, then at equilibrium:

• $[H^{+}]=x$
• $[A^{-}]=x$
• $[HA]=C-x$

So: $$K_a=\frac{x^2}{C-x}$$

AP often allows the approximation $C-x\approx C$ when $x$ is small. A quick check is that the percent ionization should typically be less than about 5% for that approximation to be safe.

Buffers: resisting pH change is about moles, not magic

A buffer is a mixture of:

• a weak acid and its conjugate base (like $C{H}_{3}COOH/C{H}_{3}CO{O}^{-}$), or
• a weak base and its conjugate acid (like $N{H}_3/N{H}_4^{+}$)

Buffers resist pH change because added $H^{+}$ or $O{H}^{-}$ is consumed by the conjugate partner, shifting equilibrium rather than allowing free $H^{+}$ to accumulate.

Henderson-Hasselbalch: what it’s for

For an acid buffer: $$pH=p{K}_a+\log \left(\frac{[A^{-}]}{[HA]}\right)$$

This equation is most useful when you already have both components present. It is especially effective in titration buffer regions. But the deeper idea the exam tests is ratio control: pH depends on the ratio of conjugate base to weak acid, not their absolute values.

Buffer capacity and its practical meaning

Buffer capacity is the amount of acid or base the buffer can neutralize before pH changes dramatically. It increases when:

• the total concentration of buffer components is higher
• the buffer is closer to $[A^{-}]\approx [HA]$ (maximum capacity near $pH\approx p{K}_a$)

A common misconception is that any weak acid solution is a “buffer.” It is not. Without a meaningful amount of conjugate base present, there is little resistance to added acid or base.

Titrations: reading the curve and knowing the chemistry region-by-region

Titration problems can look different, but the workflow is consistent: stoichiometry first, then equilibrium (if needed).

Strong acid–strong base titration

• Before equivalence: pH determined by excess strong acid or base
• At equivalence: $pH\approx 7$ (neutral) at 25°C
• After equivalence: pH determined by excess strong base

The curve has a steep vertical jump centered near pH 7.

Weak acid–strong base titration (a high-frequency AP scenario)

This curve has distinct regions:

1. Initial weak acid: pH from $K_a$ equilibrium (ICE table)
2. Buffer region: mixture of $HA$ and $A^{-}$; Henderson-Hasselbalch applies well
3. Half-equivalence point: $[HA]=[A^{-}]$ so

$$pH=p{K}_a$$
This is one of the cleanest conceptual checkpoints on the exam.

1. Equivalence point: solution contains mainly $A^{-}$ (the conjugate base), so pH is greater than 7 due to hydrolysis:

$$A^{-}+H_{2}O\rightleftharpoons HA+O{H}^{-}$$ Use $K_b=K_w/K_a$ for the equilibrium.

1. After equivalence: pH from excess strong base

Weak base–strong acid titrations mirror this logic, with pH less than 7 at equivalence due to the conjugate acid.

Choosing an indicator

AP may ask which indicator is appropriate. The indicator’s transition range should fall within the steep part of the curve near the equivalence point. That depends on whether equivalence is acidic, basic, or neutral, not on personal preference.

Particle diagrams: where students lose “free response” points

Particle diagrams are not art; they are chemical accounting. Common errors come from misrepresenting dissociation, forgetting spectator ions, or showing species that cannot coexist in large amounts.

Frequent diagram mistakes to avoid

• Drawing a strong acid as mostly undissociated molecules. Strong acids should appear essentially fully ionized.
• Drawing a weak acid as fully ionized. Weak acids should be mostly $HA$, with smaller amounts of $H^{+}$ and $A^{-}$.
• Inventing free __MATH_INLINE_56__ without water. In aqueous solutions, protons are associated with water (often represented as $H_3{O}^{+}$).
• Ignoring charge balance. Solutions must be electrically neutral overall.
• Misunderstanding buffers. A buffer diagram must show significant amounts of both members of the conjugate pair.

If you can narrate a particle diagram in words (what is present and why), you can usually draw it correctly.

High-yield habits for the AP exam

• Start every acid-base problem by listing the major species present after mixing and neutralization.
• Use moles for stoichiometry in titrations and buffer “addition” problems; convert to concentrations only when needed.
• Know the conceptual anchors: strong vs. weak, half-equivalence ($pH=p{K}_a$), and why equivalence points are not always at 7.
• Treat $K_a$, $K_b$, and $K_w$ as a connected system, not separate topics.
• For diagrams and explanations, prioritize extent of ionization, dominant species, and electrical neutrality.

Acids and bases reward clarity. When you can identify the governing chemistry in each situation, the calculations become shorter, the graphs make sense, and the free-response explanations write themselves.