Defining Acids and Bases
Few concepts in chemistry have been more carefully developed than the definitions of acids and bases. Three major frameworks exist, each more general than the last:
Arrhenius definition (1884): An acid produces H⁺ ions in water; a base produces OH⁻ ions. Simple and useful, but limited to aqueous solutions.
Brønsted-Lowry definition (1923): An acid is a proton donor; a base is a proton acceptor. This works in any solvent and explains reactions between acids and bases in the gas phase.
Lewis definition (1923): An acid is an electron-pair acceptor; a base is an electron-pair donor. The most general definition, covering reactions like the formation of complex ions.
For most introductory chemistry, the Brønsted-Lowry framework is the most useful starting point.
Strong vs. Weak Acids and Bases
Strong acids dissociate completely in water — every molecule gives up its proton: - HCl → H⁺ + Cl⁻ (100% dissociation) - HNO₃, H₂SO₄, HClO₄, HBr, HI are the other common strong acids
Weak acids only partially dissociate, reaching an equilibrium: - CH₃COOH ⇌ H⁺ + CH₃COO⁻ (only ~1% dissociation for acetic acid at typical concentrations)
Similarly, strong bases (NaOH, KOH, Ca(OH)₂) dissociate completely, while weak bases (NH₃, amines) partially accept protons from water.
The degree of dissociation is quantified by the acid dissociation constant Ka — a larger Ka means a stronger acid.
The pH Scale
pH measures the hydrogen ion concentration in aqueous solution:
pH = −log[H⁺]
The scale typically runs from 0 to 14, though values outside this range are possible for very concentrated solutions: - pH 0–6: acidic (more H⁺ than OH⁻) - pH 7: neutral (equal H⁺ and OH⁻ at 25°C) - pH 8–14: basic/alkaline (more OH⁻ than H⁺)
Each unit change in pH represents a 10-fold change in hydrogen ion concentration. Lemon juice (pH ≈ 2) is 100,000 times more acidic than black coffee (pH ≈ 7).
Familiar pH Values
| Substance | Approximate pH |
|---|---|
| Stomach acid (HCl) | 1–2 |
| Lemon juice | 2 |
| Vinegar (acetic acid) | 3 |
| Coffee | 5 |
| Milk | 6.5 |
| Pure water | 7 |
| Blood | 7.4 |
| Baking soda | 8.3 |
| Ammonia | 11 |
| Bleach | 12–13 |
| Drain cleaner (NaOH) | 14 |
Neutralization Reactions
When an acid and a base react, they undergo neutralization: the proton from the acid is transferred to the base, producing water and a salt. The general pattern:
Acid + Base → Salt + Water
The classic example is hydrochloric acid reacting with sodium hydroxide:
HCl + NaOH → NaCl + H₂O
In net ionic form, all strong acid-strong base neutralizations reduce to the same reaction:
H⁺ + OH⁻ → H₂O
The "salt" formed varies with the acid and base used. NaCl is table salt; NH₄Cl is ammonium chloride (used in cough drops); Ca₃(PO₄)₂ is calcium phosphate (a component of bones).
Neutralization Is Not Always Neutral
When a strong acid reacts with a weak base (e.g., HCl + NH₃ → NH₄Cl), the resulting salt solution is acidic because the ammonium ion (NH₄⁺) is itself a weak acid. When a weak acid reacts with a strong base (e.g., CH₃COOH + NaOH → CH₃COONa + H₂O), the resulting solution is basic. Only strong acid + strong base neutralization gives a truly neutral (pH 7) solution.
Conjugate Acid-Base Pairs
In the Brønsted-Lowry framework, every acid-base reaction involves two conjugate pairs. When acetic acid donates a proton to water:
CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺
- CH₃COOH (acid) and CH₃COO⁻ (conjugate base) are a pair
- H₂O (base) and H₃O⁺ (conjugate acid) are a pair
The stronger the acid, the weaker its conjugate base. HCl is a strong acid; Cl⁻ is an extremely weak base. Acetic acid is a weak acid; acetate (CH₃COO⁻) is a moderately strong base.
Buffers: Resisting pH Change
A buffer is a solution that resists changes in pH when small amounts of acid or base are added. Buffers contain a weak acid and its conjugate base (or a weak base and its conjugate acid) in comparable amounts.
Human blood is buffered at pH 7.35–7.45 primarily by the bicarbonate buffer system:
CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻
When acid is added (H⁺ increases), bicarbonate (HCO₃⁻) absorbs the protons. When base is added (H⁺ decreases), carbonic acid (H₂CO₃) releases protons. This tight pH control is critical — blood pH outside 7.2–7.6 is life-threatening.
Acid-Base Titration
Titration is a technique for determining the concentration of an unknown acid or base solution by slowly adding a known-concentration solution until neutralization is complete. The point of exact neutralization is the equivalence point.
A pH indicator changes color near the equivalence point, or a pH meter can track the change directly. The volume of titrant used, combined with its concentration, gives the moles of acid/base that reacted — and from stoichiometry, the concentration of the unknown.
Example: A 25.00 mL sample of HCl solution requires 32.45 mL of 0.1000 M NaOH to reach the equivalence point.
Moles NaOH = 0.03245 L × 0.1000 mol/L = 0.003245 mol Since HCl + NaOH react 1:1, moles HCl = 0.003245 mol [HCl] = 0.003245 mol / 0.02500 L = 0.1298 M
Real-World Importance of Acid-Base Chemistry
Industrial: Sulfuric acid is the most produced industrial chemical in the world, used in fertilizer production, metal refining, and oil refining.
Environmental: Acid rain (pH as low as 4.2) results from SO₂ and NOₓ emissions reacting with water vapor. It damages ecosystems, corrodes buildings, and leaches metals from soil.
Medicine: Many drugs are weak acids or bases; their absorption in the body depends critically on the pH of the stomach and intestines.
Food science: Leavening agents in baking (baking powder = baking soda + acid) rely on acid-base reactions to produce CO₂ bubbles that make bread and cakes rise.