pH Calculator
Calculate pH from [H⁺] concentration, or find [H⁺] and [OH⁻] from pH. Includes strong acid/base calculations and buffer pH reference.
Enter your values above to see the results.
Tips & Notes
- ✓The pH scale is logarithmic — a difference of 1 pH unit represents a 10-fold change in [H⁺]. pH 3 is 10× more acidic than pH 4, and 100× more acidic than pH 5. Always think in terms of powers of 10.
- ✓For strong acids, [H⁺] = molarity of the acid (assuming complete dissociation). For 0.01 M HCl: [H⁺] = 0.01 M = 10⁻² M → pH = 2. For 0.1 M NaOH: [OH⁻] = 0.1 M → pOH = 1 → pH = 13.
- ✓Use the Henderson-Hasselbalch equation for buffer pH: pH = pKa + log([A⁻]/[HA]). For a 50:50 mixture of conjugate base and acid, pH = pKa. This is why buffer solutions resist pH change.
- ✓At 25°C, pH + pOH = 14 (the water autoionization constant Kw = 10⁻¹⁴). At 37°C (body temperature), Kw = 2.4 × 10⁻¹⁴, so pH + pOH ≈ 13.62 — physiological pH 7.4 is slightly basic relative to neutral.
- ✓Temperature affects pH measurement significantly — calibrate pH meters at the measurement temperature using standard buffers. Many meters have automatic temperature compensation (ATC), but verify it is engaged for accurate readings.
Common Mistakes
- ✗Applying pH = -log[H⁺] directly to weak acids without accounting for partial dissociation — 0.1 M acetic acid (pKa 4.76) has pH ≈ 2.87, not 1 as it would be for a strong acid at the same concentration.
- ✗Confusing concentration with activity — pH meters measure hydrogen ion activity (aH⁺), not concentration, and these differ in concentrated solutions. pH = -log(aH⁺) = -log(γ × [H⁺]), where γ is the activity coefficient (≈1 at low ionic strength).
- ✗Forgetting that pH is temperature-dependent — pure water has pH 7.0 at 25°C but pH 6.14 at 60°C and pH 7.47 at 0°C. Always specify temperature when reporting precise pH values.
- ✗Adding excess acid or base without checking pH after each addition — buffer capacity is finite. Adding 1 mL of 1 M HCl to 10 mL of pH 7 phosphate buffer can drop pH by 0.5-2 units depending on buffer concentration.
- ✗Not rinsing the pH electrode between measurements — carryover from a previous sample introduces measurement error. Rinse with deionized water and gently blot (do not rub) with lens paper before each reading.
pH Calculator Overview
pH is the master variable of aqueous chemistry — it governs solubility, reaction rates, enzyme activity, protein structure, corrosion, and the safety of drinking water, all through the concentration of a single ion: H⁺. The logarithmic scale compresses a 10¹⁴-fold range of [H⁺] concentrations into a 0-14 scale that chemists and engineers work with daily.
pH and ion concentration relationships:
pH = −log₁₀[H⁺] | [H⁺] = 10^(−pH) | pOH = 14 − pH (at 25°C) | [H⁺][OH⁻] = 10⁻¹⁴
EX: 0.005 M HCl (strong acid, fully dissociates) → [H⁺] = 0.005 = 5 × 10⁻³ M → pH = −log(5 × 10⁻³) = 2.30. pOH = 14 − 2.30 = 11.70. [OH⁻] = 10⁻¹¹·⁷⁰ = 2.0 × 10⁻¹² M.Henderson-Hasselbalch for buffer pH:
pH = pKa + log([A⁻]/[HA]) | Buffer range: pKa ± 1 pH unit
EX: Phosphate buffer pH 7.2 using H₂PO₄⁻/HPO₄²⁻ (pKa = 7.20). log([HPO₄²⁻]/[H₂PO₄⁻]) = 7.2 − 7.2 = 0 → ratio = 1:1. Mix equal moles of monosodium phosphate and disodium phosphate for pH 7.2 buffer.pH scale — common solutions reference:
| pH | [H⁺] (mol/L) | [OH⁻] (mol/L) | Example Solution | Character |
|---|---|---|---|---|
| 0 | 1 | 10⁻¹⁴ | 1 M HCl | Strongly acidic |
| 1 | 10⁻¹ | 10⁻¹³ | Gastric acid, 0.1 M HCl | Strongly acidic |
| 3 | 10⁻³ | 10⁻¹¹ | Vinegar, lemon juice | Acidic |
| 5 | 10⁻⁵ | 10⁻⁹ | Coffee, rain | Weakly acidic |
| 7 | 10⁻⁷ | 10⁻⁷ | Pure water (25°C) | Neutral |
| 7.4 | 4×10⁻⁸ | 2.5×10⁻⁷ | Human blood | Slightly basic |
| 9 | 10⁻⁹ | 10⁻⁵ | Baking soda, seawater | Basic |
| 11 | 10⁻¹¹ | 10⁻³ | Household ammonia | Strongly basic |
| 14 | 10⁻¹⁴ | 1 | 1 M NaOH | Strongly basic |
| Buffer System | pKa | Useful pH Range | Common Application |
|---|---|---|---|
| Citric acid / citrate | 3.13, 4.76, 6.40 | 3.0 – 6.2 | Food chemistry, cell biology |
| Acetic acid / acetate | 4.76 | 3.8 – 5.8 | Electrophoresis, biochemistry |
| MES | 6.15 | 5.5 – 6.7 | Protein biochemistry |
| PIPES | 6.76 | 6.1 – 7.5 | Cell culture |
| Phosphate (H₂PO₄⁻/HPO₄²⁻) | 7.20 | 5.8 – 8.0 | Biochemistry, PBS |
| HEPES | 7.55 | 6.8 – 8.2 | Cell culture, clinical |
| Tris (Tris-HCl) | 8.06 | 7.0 – 9.0 | Molecular biology, gels |
| Borate | 9.24 | 8.0 – 10.0 | Electrophoresis |
Frequently Asked Questions
pH = -log₁₀[H⁺], where [H⁺] is in mol/L. The negative logarithm converts the very small concentrations typical of [H⁺] into a convenient 0-14 scale. Examples: [H⁺] = 1 × 10⁻³ M → pH = 3; [H⁺] = 5 × 10⁻⁸ M → pH = -log(5 × 10⁻⁸) = 7.30; [H⁺] = 0.05 M → pH = 1.30. Reverse: [H⁺] = 10^(-pH). At pH 7.40 (blood): [H⁺] = 10^(-7.40) = 3.98 × 10⁻⁸ M = 39.8 nM.
At 25°C, pH + pOH = pKw = 14, where Kw = [H⁺][OH⁻] = 1 × 10⁻¹⁴ is the water ionization constant. From pH, find pOH = 14 - pH, then [OH⁻] = 10^(-pOH). Example: pH 9 → pOH = 5 → [OH⁻] = 10⁻⁵ M = 10 µM. At pH 7 (neutral): [H⁺] = [OH⁻] = 10⁻⁷ M. At pH < 7: acidic, [H⁺] > [OH⁻]. At pH > 7: basic, [OH⁻] > [H⁺]. In blood at 37°C, pKw ≈ 13.62, so the neutral point is pH 6.81 — blood pH 7.4 is slightly above the neutral point for its temperature.
A buffer resists pH change upon addition of small amounts of acid or base, by containing a weak acid and its conjugate base in comparable concentrations. Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA]). Example: acetate buffer with pKa 4.76, [CH₃COO⁻] = 0.1 M, [CH₃COOH] = 0.05 M → pH = 4.76 + log(0.1/0.05) = 4.76 + log(2) = 4.76 + 0.301 = 5.06. Buffer capacity is maximum when [A⁻]/[HA] = 1, i.e., pH = pKa. Effective buffering range is pKa ± 1 pH unit.
pH scale reference points: gastric acid pH 1-2 (very acidic); lemon juice pH 2-2.5; vinegar pH 2.5-3.5; coffee pH 5-5.5; normal rain pH 5.6; pure water pH 7.0 at 25°C; blood pH 7.35-7.45; seawater pH 7.8-8.3; baking soda solution pH 8.5; household ammonia pH 11-12; bleach pH 12-13; drain cleaner (NaOH) pH 13-14. In physiology: blood pH 7.35-7.45 is strictly maintained; pH <7.35 is acidosis; pH >7.45 is alkalosis. A shift of 0.1 pH unit in blood is clinically significant.
For a weak acid HA with concentration C and acid dissociation constant Ka: [H⁺] = √(Ka × C) when Ka << C (approximation valid when degree of dissociation < 5%). pH = -log[H⁺] = ½ × (pKa - log C). Example: 0.1 M acetic acid, Ka = 1.75 × 10⁻⁵ (pKa = 4.756). [H⁺] = √(1.75 × 10⁻⁵ × 0.1) = √(1.75 × 10⁻⁶) = 1.322 × 10⁻³ M. pH = -log(1.322 × 10⁻³) = 2.88. Verify approximation: degree of dissociation = 1.322 × 10⁻³ / 0.1 = 1.32% < 5% ✓.
Enzymes are proteins whose active site shape and charge distribution depend on the ionization state of amino acid residues — which changes with pH. Most enzymes have a narrow pH optimum: pepsin (stomach) pH 2; salivary amylase pH 6.8; most intracellular enzymes pH 7-7.5; alkaline phosphatase pH 9-10. Outside the optimum by ±1-2 pH units, activity drops sharply. DNA stability also depends on pH — extreme pH denatures the double helix by disrupting hydrogen bonds. In fermentation, metabolic acid production lowers pH and eventually inhibits the microorganism unless buffered. Buffer selection in biotechnology is critical: phosphate (pH 5.8-8.0), Tris (pH 7.0-9.0), HEPES (pH 6.8-8.2), citrate (pH 3.0-6.2).