MCAT topics → Water and Solutions → Ka, Kb, pKa, pKb and Acid Strength

Ka, Kb, pKa, pKb and Acid Strength

MCAT trap: Inverts the relationship between pKa magnitude and acid strength. A higher pKa means a weaker acid; pKa = -log Ka, so larger pKa corresponds to smaller Ka and less dissociation.

Ka and pKa are the MCAT's primary language for quantifying acid strength. Ka is the equilibrium constant for the dissociation reaction HA ⇌ H⁺ + A⁻ — a larger Ka means more dissociation, meaning a stronger acid. pKa is just -log(Ka), which flips the scale: lower pKa = stronger acid. The same logic applies to bases with Kb and pKb. These aren't just definitions to memorize — the exam uses them as tools to rank compounds, predict solution pH, and analyze biochemical behavior in passage contexts.

The MCAT tests this concept from multiple angles. Straightforward recall questions ask you to interpret Ka or pKa values directly. Calculation questions require you to set up an ICE table for a weak acid and approximate [H⁺] to find pH. Passage-based questions might give you a table of pKa values and ask you to identify the dominant species at physiological pH, or to predict which compound acts as an acid vs. a buffer component. The Ka × Kb = Kw relationship shows up when a question gives you the Ka of an acid and asks about the strength of its conjugate base — a classic trap for students who think these values are independent.

The biggest pitfall is the pKa direction inversion: students see a larger number and assume it means stronger, but the negative log flips everything. Acetic acid (pKa ~4.75) is a stronger acid than ammonium (pKa ~9.25) — the smaller number wins. A related trap is assuming that equal concentrations of a weak and strong acid produce the same [H⁺]. They don't. A 0.1 M solution of acetic acid has a much higher pH than 0.1 M HCl because weak acids only partially dissociate. Getting these directional relationships locked in is what separates students who just know the formulas from students who can actually use them under pressure.

Common misconceptions

Common mistake
Wrong: A higher pKa means a stronger acid.
Right: A higher pKa means a weaker acid; pKa = -log Ka, so larger pKa corresponds to smaller Ka and less dissociation.
pKa = -log(Ka), so the negative sign reverses the intuitive direction. A Ka of 10⁻² gives pKa = 2 (strong-ish weak acid), while a Ka of 10⁻¹⁰ gives pKa = 10 (very weak acid). The math is unambiguous: higher pKa means smaller Ka means less dissociation means weaker acid. Anchor this with a concrete example — formic acid (pKa 3.7) is a stronger acid than acetic acid (pKa 4.75), not the other way around.
Common mistake
Wrong: A weak acid at the same concentration as a strong acid produces the same [H⁺] and therefore the same pH.
Right: A weak acid only partially dissociates, producing lower [H⁺] and a higher pH than a strong acid at the same concentration.
The key word in 'weak acid' is partial dissociation. Even if you dissolve 0.1 mol of acetic acid in water, only a small fraction actually releases H⁺ — the Ka tells you exactly how much. A strong acid like HCl at 0.1 M dissociates completely, giving [H⁺] = 0.1 M and pH = 1. Acetic acid at 0.1 M gives [H⁺] around 0.0013 M and pH around 2.9. Same concentration, very different [H⁺] — that's the whole point of Ka.
Common mistake
Wrong: Ka and Kb of a conjugate pair are independent values that must be looked up separately.
Right: For a conjugate acid-base pair, Ka × Kb = Kw = 10⁻¹⁴ at 25°C, so knowing one determines the other.
Ka and Kb for a conjugate pair are linked through water's autoionization constant: Ka × Kb = Kw = 1×10⁻¹⁴ at 25°C. This means if you know the Ka of an acid, you immediately know the Kb of its conjugate base — you don't need a separate table. For example, if acetic acid has Ka = 1.8×10⁻⁵, then acetate ion has Kb = 10⁻¹⁴ / 1.8×10⁻⁵ ≈ 5.6×10⁻¹⁰. Stronger acids have weaker conjugate bases, and this equation quantifies exactly how much weaker.
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What the exam tests

  1. Know the definitions of Ka and pKa and be able to use pKa = -log(Ka) to convert between them — and correctly interpret which direction means stronger or weaker acid.
  2. Calculate the pH of a weak acid solution by writing the Ka expression, setting up an ICE table, and using the small-x approximation when valid.
  3. Use the relationship Ka × Kb = Kw = 1×10⁻¹⁴ to determine the strength of a conjugate base from its parent acid's Ka, or vice versa.
  4. Given a set of pKa values, rank acids by strength and predict which form (protonated or deprotonated) predominates at a specified pH — especially relevant for amino acids and buffer systems.

Can you avoid these mistakes?

Compound A has pKa = 3.2 and Compound B has pKa = 7.8. Which is the stronger acid, and approximately how many times greater is its Ka than the other compound's Ka?
You have a 0.10 M solution of a weak acid HA with Ka = 1.0×10⁻⁵. Set up an ICE table and use the small-x approximation to find [H⁺] and the pH of this solution. Then check whether the approximation was valid.
The Ka of ammonium ion (NH₄⁺) is 5.6×10⁻¹⁰. Without looking anything up, calculate the Kb of ammonia (NH₃). Is ammonia a strong or weak base?
A molecule has three ionizable groups with pKa values of 2.1, 6.0, and 10.5. At physiological pH (7.4), which groups are protonated and which are deprotonated? What is the net charge state of the molecule, assuming it starts with all groups protonated?

Related topics

pH, pOH, and the Ion Product of Water Chemical Equilibrium and Keq Water — Polarity, Hydrogen Bonding, Anomalous Density Brønsted-Lowry and Lewis Acids and Bases Buffers and Henderson-Hasselbalch

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