Acid and Base Properties

Once you can classify a compound as an acid or a base, the next question the MCAT asks is how strong — and what that strength does to a solution's pH. This subtopic covers the autoionization of water, pH and pOH, the standard lists of strong acids and bases, the equilibrium constants Ka and Kb for weak acids and bases, how a conjugate acid and base's strengths are quantitatively linked, how molecular structure affects acid strength, and what happens to pH when an acid and a base neutralize each other.

Key Takeaways

• Autoionization of water: H₂O + H₂O ⇌ H₃O⁺ + OH⁻; Kw = [H₃O⁺][OH⁻] = 1.0×10⁻¹⁴ at 25°C, giving [H₃O⁺] = [OH⁻] = 1.0×10⁻⁷ M in pure water. Kw is fixed at a given temperature but changes with temperature.

• pH/pOH: pH = −log[H⁺], pOH = −log[OH⁻], pH + pOH = 14 at 298 K. pH < 7 = acidic, pH > 7 = basic, pH = 7 = neutral. Estimate a p-value from n×10⁻ᵐ as roughly m − 0.n.

• Strong acids/bases (HCl, HBr, HI, H₂SO₄, HNO₃, HClO₄, HClO₃; LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)₂, Sr(OH)₂, Ba(OH)₂) dissociate completely; everything else is weak.

• Ka/Kb: Ka = [H₃O⁺][A⁻]/[HA] for a weak acid, Kb = [B⁺][OH⁻]/[BOH] for a weak base — smaller values mean weaker dissociation. Ka < 1.0 = weak acid, Kb < 1.0 = weak base.

• Ka × Kb = Kw for any conjugate acid-base pair — a strong member forces the other to be weak, since their product is fixed.

• Induction: electronegative atoms near the acidic proton withdraw electron density, weaken the bond, and increase acid strength.

• Salt hydrolysis: strong acid + strong base → neutral; strong acid + weak base → acidic (conjugate acid hydrolysis); weak acid + strong base → basic (conjugate base hydrolysis); weak acid + weak base → depends on relative strengths.

Autoionization of Water and Kw

Water molecules react with each other in a process called the autoionization (or self-ionization) of water: one water molecule donates a proton to another, producing a hydronium ion (H₃O⁺) and a hydroxide ion (OH⁻). The reaction is reversible, and it's central to understanding the acidity or basicity of any aqueous solution.

H₂O(l) + H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)

The equilibrium constant for this reaction is the water dissociation constant, Kw:

Kw = [H₃O⁺][OH⁻]

Water itself doesn't appear in the expression — as with any equilibrium constant, pure liquids and solids are left out because their concentrations don't meaningfully change.

At 25°C, Kw = 1.0×10⁻¹⁴. In pure water, H₃O⁺ and OH⁻ are produced in a 1:1 ratio, so [H₃O⁺] = [OH⁻] = 1.0×10⁻⁷ M. Kw is an equilibrium constant, so at a fixed temperature it's fixed too — the product of [H₃O⁺] and [OH⁻] in any aqueous solution at 298 K must equal 1.0×10⁻¹⁴, whether that solution is neutral, acidic, or basic. Kw does change if the temperature changes, since it's temperature-dependent like any equilibrium constant.

pH, pOH, and the p-Scale

Because [H₃O⁺] and [OH⁻] are typically very small numbers, chemists track them using a p-scale: the negative logarithm of a concentration.

pH = −log[H⁺]

pOH = −log[OH⁻]

Adding an acid to water increases [H₃O⁺]. For example, hydrochloric acid dissociates completely: HCl(aq) → H⁺(aq) + Cl⁻(aq). The rising [H⁺] lowers pH (pH < 7). Adding a base increases [OH⁻] — sodium hydroxide dissociates as NaOH(aq) → Na⁺(aq) + OH⁻(aq) — which lowers [H₃O⁺] in turn and raises pH (pH > 7).

Because Kw links [H₃O⁺] and [OH⁻] at a fixed temperature, pH and pOH are linked too. At 298 K:

pH + pOH = 14

So as pH rises, pOH falls by the same amount, and vice versa. At 298 K: pH < 7 (pOH > 7) means an excess of H₃O⁺ — acidic; pH > 7 (pOH < 7) means an excess of OH⁻ — basic; pH = pOH = 7 means the two are equal — neutral.

Worked example: What is the pH of a 0.010 M HCl solution? HCl is a strong acid (see below), so it dissociates completely: [H⁺] = 0.010 M = 1.0×10⁻² M. pH = −log(1.0×10⁻²) = 2.00.

Estimating p-Values Without a Calculator

The MCAT doesn't allow calculators, so it helps to have a shortcut for estimating a p-value from a concentration written in proper scientific notation, n×10⁻ᵐ (where 1 ≤ n < 10). Using log(xy) = log(x) + log(y):

p-value ≈ m − 0.n

Here 0.n means sliding n's decimal point one place to the left — dividing n by 10. This works as an approximation because log(n) for n between 1 and 10 is always a decimal between 0 and 1 (log 1 = 0, log 10 = 1): the closer n is to 1, the closer log(n) is to 0, and the closer n is to 10, the closer log(n) is to 1. Sliding the decimal is a fast stand-in for that logarithm — it's not exact, but it gets you close enough to pick the right answer choice.

For example, [H⁺] = 4×10⁻⁵ M gives an estimated pH ≈ 5 − 0.4 = 4.6 (the precise value is 4.40 — close enough to distinguish it from the neighboring answer choices on a multiple-choice test).

Strong Acids and Bases

Strong acids and bases dissociate completely into their ions in aqueous solution, which is what makes them so effective at shifting a solution's pH. There are only a handful of each, and the MCAT expects you to recognize them on sight:

Strong Acids:

• Hydrochloric acid (HCl)

• Hydrobromic acid (HBr)

• Hydroiodic acid (HI)

• Sulfuric acid (H₂SO₄)

• Nitric acid (HNO₃)

• Perchloric acid (HClO₄)

• Chloric acid (HClO₃)

Strong Bases:

• Lithium hydroxide (LiOH)

• Sodium hydroxide (NaOH)

• Potassium hydroxide (KOH)

• Rubidium hydroxide (RbOH)

• Cesium hydroxide (CsOH)

• Calcium hydroxide (Ca(OH)₂)

• Strontium hydroxide (Sr(OH)₂)

• Barium hydroxide (Ba(OH)₂)

Any acid or base not on one of these lists should be treated as weak.

Weak Acids and Bases: Ka and Kb

Weak acids and bases only partially dissociate in water, settling into an equilibrium where both the ionized and unionized forms are present together.

For a weak acid HA:

HA(aq) + H₂O(l) ⇌ H₃O⁺(aq) + A⁻(aq)

Ka = [H₃O⁺][A⁻] / [HA]

The smaller the Ka, the weaker the acid — the less it dissociates. As with Kw, water is left out of the expression because it's a liquid. Acetic acid (CH₃COOH), the acid in vinegar, is a classic weak acid example.

For a weak base BOH:

BOH(aq) ⇌ B⁺(aq) + OH⁻(aq)

Kb = [B⁺][OH⁻] / [BOH]

The smaller the Kb, the weaker the base. Ammonia (NH₃) is a classic weak base example. As a general rule, a species is a weak acid if Ka < 1.0 and a weak base if Kb < 1.0.

Worked example: Acetic acid has Ka = 1.8×10⁻⁵. What is the pH of a 0.100 M acetic acid solution? Setting up the equilibrium (assuming the dissociated amount, x, is small relative to 0.100 M):

Ka ≈ x² / 0.100 → x² = (1.8×10⁻⁵)(0.100) = 1.8×10⁻⁶ → x = [H₃O⁺] ≈ 1.3×10⁻³ M

pH = −log(1.3×10⁻³) ≈ 2.87

Conjugate Acid-Base Pairs: Ka × Kb = Kw

Picture a weak acid HA dissociating in water while its conjugate base A⁻ simultaneously acts as a base toward water. Add those two equilibria together and the water molecules on each side cancel out — what's left is exactly the autoionization of water. Because the reactions combine this way, their equilibrium constants combine the same way: the product of the individual Ka and Kb equals Kw.

Ka(acid) × Kb(conjugate base) = Kw

Kb(base) × Ka(conjugate acid) = Kw

This relationship makes the strength of a conjugate acid-base pair a package deal: if one member is strong (a large Ka or Kb), the other must be weak (a small Kb or Ka), since their product is fixed at Kw. For acetic acid and its conjugate base, acetate (CH₃COO⁻): Kb(acetate) = Kw / Ka = (1.0×10⁻¹⁴) / (1.8×10⁻⁵) ≈ 5.6×10⁻¹⁰ — a very small Kb, consistent with acetate being a very weak base.

Induction and Acid Strength

Molecular structure affects how strong an acid is, and one major factor is induction: electronegative atoms positioned near an acidic proton pull electron density out of the bond holding that proton. This weakens the bond and makes the proton easier to lose, so acids with electronegative atoms closer to the acidic hydrogen are stronger than acids without them.

A supplied example illustrating this principle (verified independently, not from the source transcript, which named this comparison without completing it): bicarbonate (HCO₃⁻) is a noticeably stronger acid than water itself (pKa ≈ 10.3 for HCO₃⁻ acting as an acid, versus pKa ≈ 15.7 for water). The carbonate-based framework surrounding HCO₃⁻'s remaining O–H bond is more electronegative and more effective at withdrawing electron density than water's own structure, which weakens that O–H bond and makes the proton easier to donate than one on a plain water molecule.

Salt Formation and Neutralization Reactions

A neutralization reaction occurs when an acid and a base react to form a salt and, often, water:

Acid + Base → Salt + Water

Because acids and bases can each be strong or weak, there are four possible combinations, and each produces a different outcome for the resulting solution's pH.

Strong acid + strong base (e.g., HCl + NaOH → NaCl + H₂O): both reactants dissociate completely, and the net ionic equation is just H⁺ + OH⁻ → H₂O. Na⁺ and Cl⁻ are spectator ions, taking no further part in the reaction. Because the acid and base neutralize each other completely, the resulting solution is neutral (pH = 7).

Strong acid + weak base (e.g., HCl + NH₃ → NH₄Cl): the strong acid dissociates completely, but ammonia is a weak base that reacts with H⁺ to form NH₄⁺ rather than fully ionizing beforehand. Cl⁻ is a spectator ion. The resulting solution is slightly acidic, because the ammonium ion (NH₄⁺) — the conjugate acid of a weak base — can hydrolyze in water to regenerate H₃O⁺.

Weak acid + strong base (e.g., HClO + NaOH → NaClO + H₂O): sodium hydroxide dissociates completely, while hypochlorous acid (HClO) only partially dissociates and reacts with the resulting OH⁻ ions. Na⁺ is a spectator ion. The resulting solution is slightly basic, because the hypochlorite ion (ClO⁻) — the conjugate base of a weak acid — can hydrolyze in water to regenerate OH⁻.

Weak acid + weak base (e.g., HClO + NH₃ → NH₄ClO): both reactants only partially dissociate. The resulting solution's pH depends on the relative strengths of the two species involved and how far each of their conjugate ions hydrolyzes — there's no single direction that applies to every weak-acid/weak-base pairing.

Common MCAT Mistakes

• Assuming a stronger acid always has a weaker conjugate base by some fixed amount. The relationship is multiplicative, not additive — Ka × Kb = Kw is a fixed product, so a large Ka forces a very small Kb, but the exact split depends on the specific Ka value.

• Forgetting Kw itself changes with temperature. Kw = 1.0×10⁻¹⁴ only holds at 25°C; at a different temperature, [H₃O⁺] = [OH⁻] = 1.0×10⁻⁷ M and pH = 7 for "neutral" no longer apply.

• Trying to memorize every acid's strength instead of the two short lists. Any acid or base not on the strong-acid/strong-base lists is weak by default — there's no need to memorize weak acids individually.

• Assuming every acid-base neutralization gives a neutral (pH 7) solution. Only strong acid + strong base lands at pH 7. Strong-weak combinations land off-neutral because the resulting conjugate ion hydrolyzes.

MCAT-Style Concept Check

Question: A 0.050 M solution of a weak acid HA has a Ka of 4.0×10⁻⁸. Which of the following best describes the corresponding Kb of its conjugate base, A⁻, at 25°C?

• A) Kb ≈ 4.0×10⁻⁸ (equal to Ka)

• B) Kb ≈ 2.5×10⁻⁷ (larger than Ka, since Ka is small)

• C) Kb ≈ 1.6×10⁻²¹ (much smaller than Ka)

• D) Kb cannot be determined without knowing [HA] at equilibrium

Answer: B

Explanation: Ka × Kb = Kw, so Kb = Kw / Ka = (1.0×10⁻¹⁴) / (4.0×10⁻⁸) = 2.5×10⁻⁷. Because Ka is small (a weak acid), Kb comes out comparatively larger — consistent with A⁻ being a moderately effective weak base. This calculation depends only on Ka and Kw, not on the solution's concentration, which is why (D) is a trap.

FAQ

Why doesn't water appear in the Kw, Ka, or Kb expressions?

Equilibrium constant expressions exclude pure liquids and solids, since their concentrations are treated as effectively constant and don't change the equilibrium position. Water acts as the solvent in these reactions, so it's left out just like any other pure liquid.

How many strong acids and strong bases do I need to memorize?

Seven strong acids (HCl, HBr, HI, H₂SO₄, HNO₃, HClO₄, HClO₃) and eight strong bases (LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)₂, Sr(OH)₂, Ba(OH)₂). Anything not on these two lists is weak.

Why does a strong acid + weak base reaction produce an acidic solution instead of a neutral one?

The strong acid fully dissociates, but the weak base only partially reacts, converting into its conjugate acid (e.g., NH₃ → NH₄⁺). That conjugate acid can then hydrolyze in water to regenerate H₃O⁺, pushing the resulting solution's pH below 7.

How does induction make an acid stronger?

Electronegative atoms near the acidic proton pull electron density away from the bond holding that proton, weakening the bond and making the proton easier to release. The closer and more numerous the electronegative atoms, the stronger the inductive effect and the stronger the acid.