Acid-Base Reactions
Grade 11 · Science · Study Guide
Explore the chemistry of acids and bases — from pH and proton transfer to neutralization reactions, titration, and buffer systems.
Acids and bases are fundamental in chemistry and biology. The Arrhenius definition (1887) defines an acid as a substance that produces H⁺ ions in water and a base as one that produces OH⁻ ions. This was expanded by Brønsted and Lowry, who defined acids as proton (H⁺) donors and bases as proton acceptors. This broader view explains why ammonia (NH₃) is a base even though it contains no OH⁻ — it accepts protons. The Lewis definition goes further still: Lewis acids accept electron pairs; Lewis bases donate them.
Strong acids (HCl, HNO₃, H₂SO₄, HClO₄, HBr, HI) dissociate completely in water. Weak acids (acetic acid, carbonic acid) dissociate only partially, establishing an equilibrium. Similarly, strong bases (NaOH, KOH) dissociate completely; weak bases (NH₃) only partially. The strength of an acid is quantified by its Ka (acid dissociation constant) — larger Ka means stronger acid.
pH = −log[H⁺]. At 25°C, pure water has pH = 7 (neutral). Acids have pH < 7; bases have pH > 7. The pH scale is logarithmic — a change of 1 pH unit represents a 10-fold change in [H⁺]. pOH = −log[OH⁻], and pH + pOH = 14 at 25°C. These relationships let you calculate H⁺ and OH⁻ concentrations from pH and vice versa.
Neutralization is the reaction between an acid and a base to form salt and water: acid + base → salt + water. For example, HCl + NaOH → NaCl + H₂O. Titration uses a carefully measured volume of solution of known concentration (the titrant) added to a solution of unknown concentration until the equivalence point is reached (moles of acid = moles of base). An indicator changes color at or near the equivalence point.
Buffers resist changes in pH when small amounts of acid or base are added. They consist of a weak acid and its conjugate base (or weak base and conjugate acid) in similar concentrations. Blood is buffered at pH ~7.4 by the carbonic acid/bicarbonate system (H₂CO₃/HCO₃⁻), essential for survival.
| Strong Acids | Strong Bases |
|---|---|
| HCl (hydrochloric acid) | NaOH (sodium hydroxide) |
| HNO₃ (nitric acid) | KOH (potassium hydroxide) |
| H₂SO₄ (sulfuric acid) | Ba(OH)₂ (barium hydroxide) |
| HBr (hydrobromic acid) | Ca(OH)₂ (calcium hydroxide) |
| HClO₄ (perchloric acid) | LiOH (lithium hydroxide) |
| pH | [H⁺] (mol/L) | Classification |
|---|---|---|
| 0 | 1.0 | Very strong acid |
| 3 | 1.0 × 10⁻³ | Acidic (e.g., vinegar) |
| 7 | 1.0 × 10⁻⁷ | Neutral (pure water) |
| 10 | 1.0 × 10⁻¹⁰ | Basic (e.g., baking soda) |
| 14 | 1.0 × 10⁻¹⁴ | Very strong base |
| Theory | Acid | Base |
|---|---|---|
| Arrhenius | Produces H⁺ in water | Produces OH⁻ in water |
| Brønsted-Lowry | Proton donor | Proton acceptor |
| Lewis | Electron pair acceptor | Electron pair donor |
Key terms
- Brønsted-Lowry Acid
- A proton (H⁺) donor in a chemical reaction.
- Brønsted-Lowry Base
- A proton (H⁺) acceptor in a chemical reaction.
- pH
- pH = −log[H⁺]; measures hydrogen ion concentration. Acidic < 7, neutral = 7, basic > 7.
- Strong Acid
- An acid that dissociates completely in water (e.g., HCl, H₂SO₄, HNO₃).
- Weak Acid
- An acid that partially dissociates in water, establishing an equilibrium (e.g., acetic acid).
- Neutralization
- Reaction of acid + base → salt + water; at the equivalence point moles of H⁺ = moles of OH⁻.
- Titration
- Procedure using a known concentration solution (titrant) to determine the unknown concentration of another solution.
- Buffer
- Solution that resists pH change; contains a weak acid and its conjugate base in similar concentrations.
Practice questions
What is the pH of a solution with [H⁺] = 1.0 × 10⁻⁴ mol/L?
- 2
- 4
- 6
- 10
Show answer
pH = −log[H⁺] = −log(1.0 × 10⁻⁴) = 4.
According to Brønsted-Lowry theory, an acid is defined as a:
- Proton acceptor
- Proton donor
- Hydroxide producer
- Electron donor
Show answer
Brønsted-Lowry acids donate H⁺ (protons) to bases. This definition is broader than Arrhenius and applies in non-aqueous contexts too.
At 25°C, pH + pOH = ______.
Show answer
The ion product of water Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C, which means pH + pOH = 14.
What are the products of a neutralization reaction between HCl and NaOH?
- HNaClO₂ and H₂
- NaCl and H₂O
- NaH and ClOH
- Cl₂ and Na(OH)₂
Show answer
Neutralization: HCl + NaOH → NaCl + H₂O. Acid + base → salt + water.
Which of the following is a strong acid?
- Acetic acid (CH₃COOH)
- Carbonic acid (H₂CO₃)
- Nitric acid (HNO₃)
- Phosphoric acid (H₃PO₄)
Show answer
HNO₃ is one of six common strong acids that dissociate completely in water. Acetic, carbonic, and phosphoric acids are all weak acids.
Fun facts
- Stomach acid (HCl, pH ~1–2) is strong enough to dissolve metals, yet the stomach lining regenerates itself every few days to withstand this harsh environment.
- The ocean has absorbed roughly 30% of human CO₂ emissions since industrialization, causing ocean acidification — pH has dropped from ~8.2 to ~8.1. That 0.1 unit drop represents a ~26% increase in [H⁺], threatening coral skeletons made of calcium carbonate.
- Your blood must maintain pH 7.35–7.45 — a range of just 0.1 units. A drop to pH 7.0 is life-threatening; the carbonate buffer system and kidneys work constantly to keep it in range.
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