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Acids and Bases

Chapter 10: Acids and Bases

5 min read Updated Mar 26, 2026
Read the entire chapter on one page Every section in order, with the sidebar tracking where you are as you scroll.
🎯 Diagnostic: Test Your Starting Level 24 questions (2 per section). No prior reading required, see what you already know.

Aim to answer every question before checking. Missed questions point you to the sections you need most.

1. (10.1) A Brønsted-Lowry acid is:
D. Lewis acids accept electron pairs; Arrhenius acids are the subset that produce H⁺ in water. Brønsted is between: a proton donor in any solvent.
2. (10.1) A Lewis base is:
B. Every Brønsted base is a Lewis base, but not every Lewis base has a proton to accept. NH₃, :Cl⁻, and water are Lewis bases.
3. (10.2) The conjugate base of NH₄⁺ is:
A. Remove one H⁺ to get the conjugate base. Conjugate pairs differ by exactly one H⁺.
4. (10.2) A strong acid has a:
C. Cl⁻ (from HCl) is essentially inert. Strong acid ↔ weak conjugate base; weak acid ↔ stronger (but still weak) conjugate base.
5. (10.3) Which is a strong acid in water?
D. The common strong acids: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄, HClO₃. HF is a weak acid despite its reactivity.
6. (10.3) Typical strong bases in water are:
B. Heavier alkaline-earth hydroxides (Ba(OH)₂) are also strong. Ammonia is a common weak base (pKa of NH₄⁺ ≈ 9.25).
7. (10.4) A weak acid is characterized by:
A. Acetic acid Ka ≈ 1.8 × 10⁻⁵; most biomolecular weak acids have Ka between 10⁻³ and 10⁻¹⁰.
8. (10.4) For a weak base, Kb is:
C. Ka × Kb = Kw = 10⁻¹⁴ at 25 °C links the two constants for a conjugate pair.
9. (10.5) pH is defined as:
D. Each unit drop in pH means a tenfold rise in [H⁺]. At 25 °C, pH + pOH = 14.
10. (10.5) At 25 °C, pure water has:
B. Pure water autoionizes slightly: 2 H₂O ⇌ H₃O⁺ + OH⁻. At 25 °C, Kw = 10⁻¹⁴ and [H⁺] = [OH⁻] = 10⁻⁷ M.
11. (10.6) The pH of 0.01 M HCl (a strong acid, fully dissociated) is:
A. [H⁺] = 0.01 M = 10⁻² M, so pH = 2.
12. (10.6) The pH of 0.001 M NaOH is:
C. [OH⁻] = 10⁻³ M → pOH = 3 → pH = 14 - 3 = 11.
13. (10.7) For a weak acid HA at concentration C:
D. The approximation assumes dissociation is small compared with C. Works when Ka × C >> Kw and dissociation is under ~5%.
14. (10.7) Acetic acid has pKa ≈ 4.76, so Ka is approximately:
B. Ka = 10^(-pKa) = 10^(-4.76) ≈ 1.75 × 10⁻⁵.
15. (10.8) Polyprotic acids such as H₃PO₄:
A. Removing each H⁺ becomes harder because the resulting anion is increasingly negatively charged.
16. (10.8) H₂SO₄ is unusual among polyprotic acids because:
C. This is why MCAT problems treating H₂SO₄ as diprotic strong acid may slightly overestimate pH. pKa₂ is about 1.99 for HSO₄⁻.
17. (10.9) The equivalence point of a titration is:
D. Endpoint (indicator color change) is the approximate observation; equivalence point is the true stoichiometric condition.
18. (10.9) At the half-equivalence point of a weak-acid/strong-base titration:
B. At half-equivalence, [HA] = [A⁻], so pH = pKa (from Henderson-Hasselbalch).
19. (10.10) An indicator signals the endpoint by:
A. Indicators are weak acids whose HIn and In⁻ forms have different colors. Transition is visible over ~pKa ± 1.
20. (10.10) Phenolphthalein (pKa ≈ 9.3) is typically useful for titrations with:
C. Good rule: pick an indicator whose pKa sits within ±1 pH unit of the expected equivalence-point pH.
21. (10.11) A buffer is:
D. Buffers neutralize added acid or base via the common-ion effect. Blood uses a bicarbonate buffer system.
22. (10.11) The Henderson-Hasselbalch equation for a buffer is:
B. When [A⁻] = [HA], pH = pKa. Equivalent to A if you flip the ratio and swap the sign.
23. (10.12) Buffer capacity is greatest when:
A. Effective buffer range is pKa ± 1. Both components should be concentrated enough to absorb added acid or base.
24. (10.12) Adding a small amount of strong acid to a buffer:
C. Until buffer is exhausted, added H⁺ is absorbed by A⁻. Once capacity is exceeded, pH plunges rapidly.

You squeeze a lemon into your water and wince at the sour taste. You wash your hands with soap and notice how slippery it feels. That sour bite and that slippery film are your body detecting acids and bases - two categories of molecules that the MCAT tests more heavily than almost any other topic in general chemistry.

Acid-base chemistry ties together equilibrium, thermodynamics, and even organic chemistry. An MCAT passage might describe a blood buffer system and ask you to calculate the pH after adding a strong acid. Another might show a titration curve and ask you to identify the pKa of an unknown acid. Yet another might ask why a particular amino acid side chain is protonated at physiological pH. Every one of those questions uses the tools you will build in this chapter.

The good news: acid-base chemistry follows a small set of rules. Once you understand how protons move between molecules - and how to track that movement mathematically - these passages become some of the most predictable on the entire exam.

Shock Absorbers and Overflowing Cups

Think of this chapter in two halves. The first half is about definitions and strength: what makes something an acid or a base, and how completely it gives up or accepts protons. The second half is about measurement and control: pH calculations, titration curves, and buffers.

Titration is like slowly filling a cup to the brim - nothing dramatic happens until that one last drop sends water cascading over the edge (the equivalence point). Buffers are shock absorbers for pH - they absorb the bumps of added acid or base so the solution barely changes. If you understand these two metaphors deeply, you already have the intuition for the hardest problems in this chapter.

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