acids bases ph

Acids and bases are fundamental chemical species defined by three major theories. The Brønsted-Lowry theory (most used in A-Level) defines an acid as a proton (H⁺) donor and a base as a proton acceptor. The Lewis theory extends this: acids are electron pair acceptors, bases are electron pair donors.

Strong acids (HCl, HNO₃, H₂SO₄) completely dissociate in water: HCl → H⁺ + Cl⁻. Weak acids (CH₃COOH, HCN) partially dissociate, establishing equilibrium: CH₃COOH ⇌ H⁺ + CH₃COO⁻. Similarly, strong bases (NaOH, KOH) fully dissociate, while weak bases (NH₃) partially accept protons: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻.

pH is defined as: pH = -log₁₀[H⁺], where [H⁺] is hydrogen ion concentration in mol dm⁻³. At 25°C, pure water has [H⁺] = 1.0 × 10⁻⁷ mol dm⁻³, giving pH 7 (neutral). The ionic product of water is Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ mol² dm⁻⁶ at 25°C.

pKa = -log₁₀Ka measures acid strength (lower pKa = stronger acid). The acid dissociation constant Ka = [H⁺][A⁻]/[HA]. For weak acids: [H⁺] = √(Ka × [HA]) (assuming minimal dissociation).

Memory Aid - ACIDS: Always Create Ion Dissociation Systems. Strong = complete, Weak = equilibrium.

Conjugate pairs: When an acid loses H⁺, it forms its conjugate base (CH₃COOH/CH₃COO⁻). A strong acid has a weak conjugate base. Buffer solutions resist pH change, containing a weak acid and its conjugate base (or weak base and conjugate acid).

Understanding acids and bases transforms everyday chemistry into predictable science. Your stomach acid (HCl, pH ~1.5) exemplifies a strong acid—it completely dissociates to digest proteins. Antacids like magnesium hydroxide neutralize excess acid through Brønsted-Lowry proton transfer: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O.

Blood pH (7.35-7.45) is maintained by a carbonic acid-bicarbonate buffer: H₂CO₃ ⇌ H⁺ + HCO₃⁻. When you exercise, CO₂ increases, forming more H₂CO₃ and lowering pH slightly; you breathe faster to expel CO₂, restoring pH—a biological buffer in action!

Think of pH like a volume dial: each unit is a 10-fold change in [H⁺]. Lemon juice (pH 2) has 100× more H⁺ than tomato juice (pH 4). This logarithmic scale means small pH changes represent huge concentration differences.

Weak acids dominate biological systems because they buffer naturally. Acetic acid in vinegar (CH₃COOH, Ka = 1.8 × 10⁻⁵) only partially ionizes, creating an equilibrium that resists drastic pH changes when diluted—unlike HCl, which stays fully ionized.

Industrial applications: pH control is critical in fermentation (beer brewing uses pH 4-6 for optimal yeast activity), textile dyeing (different dyes require specific pH ranges), and water treatment (pH 6.5-8.5 for safe drinking water).

Analogy: A strong acid is like a bag of marbles spilling completely; a weak acid is like a zip-lock bag barely open—marbles (H⁺ ions) slowly leak out, with most staying inside.

Kw's temperature dependence means pure water at 50°C has pH < 7 but remains neutral because [H⁺] = [OH⁻]. This surprises students who think pH 7 always means neutral!