pH and Hydrogen Ion Concentration Calculator

pH and Acid-Base Chemistry Guide

The pH Scale

pH = −log₁₀[H⁺], where [H⁺] is the hydrogen ion concentration in mol/L. The scale runs 0–14 at 25°C: pH < 7 is acidic, pH = 7 is neutral (pure water at 25°C), pH > 7 is alkaline. Each unit change represents a 10-fold change in [H⁺]: pH 3 has 10× more H⁺ than pH 4, and 100× more than pH 5. Lemon juice ≈ pH 2. Black coffee ≈ pH 5. Pure water pH 7. Baking soda pH 9. Bleach pH 12. Human blood is maintained at pH 7.35–7.45 — outside this range is life-threatening.

Converting pH and [H⁺]

From pH to [H⁺]: [H⁺] = 10^(−pH). pH 3 → [H⁺] = 10⁻³ = 0.001 mol/L. pH 7 → [H⁺] = 10⁻⁷ mol/L. From [H⁺] to pH: pH = −log₁₀[H⁺]. [H⁺] = 0.01 → pH = −log(0.01) = −(−2) = 2. The ionic product of water: Kw = [H⁺][OH⁻] = 1×10⁻¹⁴ at 25°C. This means pH + pOH = 14. If pH = 4, then pOH = 10, [OH⁻] = 10⁻¹⁰ mol/L. Strong acids fully dissociate: 0.01 mol/L HCl → [H⁺] = 0.01 mol/L → pH = 2.

Buffer Solutions

A buffer resists pH change when small amounts of acid or alkali are added. Composed of a weak acid and its conjugate base (or weak base and conjugate acid). Example: ethanoic acid/sodium ethanoate buffer. Biological importance: blood is buffered at pH 7.4 by the carbonic acid/bicarbonate system. CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻. The Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA]) — used to calculate buffer pH and design buffers for specific pH values in lab work.

Indicators and pH Measurement

Indicators are weak acids that change colour when they donate or accept a proton — the colour depends on the ratio of acid to conjugate base form. Litmus: red in acid (pH < 7), blue in alkali. Universal indicator: continuous colour spectrum from red (pH 1) through green (pH 7) to purple (pH 14). Phenolphthalein: colourless below pH 8.3, pink above — used in alkaline endpoint titrations. Methyl orange: red below pH 3.1, yellow above pH 4.4 — used in acid endpoint titrations. For precision measure