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Acids, Bases, and pH Explained

From the tartness of lemon juice to the slipperiness of soapy water, acids and bases are part of everyday experience. Understanding what makes a substance acidic or basic — and how the pH scale quantifies it — is fundamental to chemistry and biology.

Defining Acids and Bases

Chemists have developed several complementary definitions of acids and bases, each more general than the last.

The Arrhenius definition (1884) is the simplest: an acid is a substance that produces hydrogen ions (H⁺) when dissolved in water; a base is a substance that produces hydroxide ions (OH⁻) in water. Hydrochloric acid (HCl) dissociates into H⁺ and Cl⁻; sodium hydroxide (NaOH) dissociates into Na⁺ and OH⁻.

The Brønsted–Lowry definition (1923) is broader: an acid is a proton (H⁺) donor; a base is a proton acceptor. This allows reactions to be described in non-aqueous solvents and covers substances the Arrhenius model misses — for example, ammonia (NH₃) has no OH⁻ to donate but clearly acts as a base by accepting a proton to form ammonium (NH₄⁺).

The even broader Lewis definition defines acids as electron-pair acceptors and bases as electron-pair donors. Lewis acids and bases are important in organic chemistry and in reactions that involve no proton transfer at all.

Strong vs. Weak Acids and Bases

A strong acid dissociates completely in water. Hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and nitric acid (HNO₃) are strong acids. In a 0.1 mol/L solution of HCl, essentially every HCl molecule has donated its proton.

A weak acid dissociates only partially. Acetic acid (CH₃COOH), the acid in vinegar, is a classic example. In a 0.1 mol/L solution only about 1.3% of molecules have donated their proton. The equilibrium constant for this dissociation is the acid dissociation constant, K_a; a smaller K_a means a weaker acid.

The same distinction applies to bases. Sodium hydroxide and potassium hydroxide are strong bases (full dissociation). Ammonia and the amine groups in biological molecules are weak bases.

The pH Scale

The pH scale is a way to express the concentration of hydrogen ions in a solution on a convenient logarithmic scale. It is defined as:

pH = −log₁₀[H⁺]

where [H⁺] is the molar concentration of hydrogen ions. Because the scale is logarithmic, each whole-number step represents a tenfold change in acidity. A solution at pH 3 is 10 times more acidic than one at pH 4, and 100 times more acidic than one at pH 5.

• pH 0–6: acidic (more H⁺ than OH⁻)
• pH 7: neutral (equal H⁺ and OH⁻, as in pure water at 25°C)
• pH 8–14: basic/alkaline (more OH⁻ than H⁺)

Calculating pH and pOH

For a strong acid, pH calculation is straightforward because complete dissociation means [H⁺] equals the initial acid concentration. For a 0.01 mol/L HCl solution: [H⁺] = 0.01 = 10⁻², so pH = −log(10⁻²) = 2.

For a weak acid you must set up an equilibrium expression using K_a and solve for [H⁺], which requires either an exact quadratic or an approximation when K_a is very small relative to the initial concentration.

The complementary measure pOH = −log₁₀[OH⁻]. At 25°C, pH + pOH = 14 (because the ion product of water, K_w, equals 10⁻¹⁴). So if pH = 11, then pOH = 3, and [OH⁻] = 10⁻³ mol/L.

Neutralisation and Salt Formation

When an acid and a base react together, the products are water and a salt. The general equation is: acid + base → salt + water. For hydrochloric acid and sodium hydroxide: HCl + NaOH → NaCl + H₂O. This is a neutralisation reaction. In a titration, a known concentration of base is added to an acid (or vice versa) until the solution reaches the equivalence point at pH 7.

Note that not all salts are neutral. Sodium acetate (from weak acid + strong base) produces a basic solution because the acetate ion hydrolyses water to regenerate some OH⁻. Ammonium chloride (from weak base + strong acid) produces a slightly acidic solution.

Buffers

A buffer is a solution that resists changes in pH when small amounts of acid or base are added. Most buffers consist of a weak acid and its conjugate base (the salt form). The blood buffer system, for example, is a mixture of carbonic acid (H₂CO₃) and bicarbonate ions (HCO₃⁻). Normal blood pH is 7.35–7.45; a drop to 7.2 (acidosis) or a rise to 7.6 (alkalosis) is life-threatening.

Buffers work because the weak acid can donate a proton to neutralise added base, while the conjugate base can accept a proton to neutralise added acid. The Henderson–Hasselbalch equation predicts the pH of a buffer solution: pH = pK_a + log([A⁻]/[HA]), where [A⁻] is the conjugate base concentration and [HA] is the weak acid concentration.

Summary

Acids donate protons (or produce H⁺ ions); bases accept protons (or produce OH⁻ ions). The pH scale runs from 0 (most acidic) to 14 (most basic), with 7 as neutral, and each unit represents a tenfold change in [H⁺]. Strong acids and bases dissociate completely; weak ones establish equilibria characterised by K_a or K_b. Neutralisation produces salt and water. Buffers, essential to living systems, exploit weak acid–conjugate base pairs to hold pH steady against perturbation.