Calculate pH of strong acids, weak acids, buffer solutions, and dilutions with step-by-step solutions. Includes pH converter, ICE tables, and Henderson-Hasselbalch equation. Perfect for GCSE and A-Level Chemistry.
pH measures hydrogen ion concentration
pOH measures hydroxide ion concentration
At 25°C, pH and pOH always sum to 14
At 25°C
Acid dissociation constant
For buffer solutions
Conservation of moles during dilution
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pH stands for "power of hydrogen" and measures how acidic or alkaline a solution is. It is defined as pH = -log₁₀[H⁺], where [H⁺] is the concentration of hydrogen ions in mol/dm³.
Because pH uses a logarithmic scale, each 1-unit change in pH represents a 10-fold change in hydrogen ion concentration. A solution with pH 3 has 10 times more H⁺ ions than one with pH 4, and 100 times more than pH 5.
At 25°C, the ionic product of water Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴. This means pH + pOH always equals 14. Pure water has pH = 7 (neutral), because [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ mol/dm³.
The pH scale runs from 0 (very acidic) through 7 (neutral) to 14 (very alkaline). Each substance in everyday life has a characteristic pH:
A buffer solution resists changes in pH when small amounts of acid or base are added. It contains a weak acid and its conjugate base (e.g. ethanoic acid + sodium ethanoate). The pH is calculated using the Henderson-Hasselbalch equation:
pH = pKa + log₁₀([A⁻]/[HA])
When you add water to an acid or base, the number of moles stays the same — only the concentration changes. The dilution equation C₁V₁ = C₂V₂ lets you calculate the new concentration or volume.
HCl is a strong acid, so [H⁺] = 0.5 mol/dm³.
pH = -log₁₀(0.5) = 0.30
[H⁺] = 10⁻ᵖᴴ = 10⁻⁴ = 0.0001 mol/dm³ (or 1 × 10⁻⁴)
ICE: Kb = x²/(0.05 − x). Using quadratic: x = [OH⁻] = 9.49 × 10⁻⁴.
pOH = -log₁₀(9.49 × 10⁻⁴) = 3.02.
pH = 14 − 3.02 = 10.98
pKa = -log₁₀(1.8 × 10⁻⁵) = 4.74.
pH = 4.74 + log₁₀(0.3/0.2) = 4.74 + 0.176 = 4.92
pH = -log₁₀[H⁺]. It measures how acidic or alkaline a solution is on a scale typically from 0 to 14, where 7 is neutral.
For a strong acid: [H⁺] = concentration × proton count, then pH = -log₁₀[H⁺]. Strong acids dissociate completely.
Strong acids dissociate 100% (HCl, HNO₃, H₂SO₄). Weak acids partially dissociate and have a Ka value (CH₃COOH, HF).
ICE stands for Initial, Change, Equilibrium. Set up concentrations for each species, substitute into Ka expression, and solve for x.
pH = pKa + log₁₀([A⁻]/[HA]). Used to calculate buffer pH from weak acid concentration, conjugate base concentration, and Ka.
Yes! Any strong acid with concentration > 1 mol/dm³ gives pH < 0. For example, 10 M HCl has pH = -1.
Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C. This is the ionic product of water and explains why pH + pOH = 14.
The dilution equation. Moles are conserved: concentration × volume before = concentration × volume after.
% dissociation = ([H⁺] at equilibrium / initial concentration) × 100. It shows what fraction of a weak acid has dissociated.
H₂SO₄ is diprotic — it releases 2 H⁺ ions per molecule. So [H⁺] = 2 × concentration, giving a lower pH than HCl at the same concentration.
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