pH Calculator

pH stands for "potential of hydrogen." It is defined as the negative base-10 logarithm of the hydrogen ion activity (approximately the concentration in dilute solutions): pH = −log₁₀[H⁺]. A lower pH means higher acidity, a higher pH means higher alkalinity. Each unit change represents a tenfold change in hydrogen ion concentration.

In aqueous solutions at 25 °C, pH + pOH = 14. This relationship comes from the autoionization constant of water: K_w = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴. Taking negative logarithms of both sides gives pH + pOH = pK_w = 14.00. This calculator uses this relationship to convert between pH, pOH, [H⁺], and [OH⁻] in all three input modes.

The relationship pH + pOH = 14 is only valid at 25 °C because K_w is temperature-dependent. At 0 °C, pK_w = 14.95 (neutral pH = 7.47). At 100 °C, pK_w = 12.26 (neutral pH = 6.14). This calculator assumes 25 °C. For accurate pH at other temperatures, you need to use the appropriate pK_w value and adjust the pH + pOH sum accordingly.

Strong acids (HCl, HNO₃, H₂SO₄, HClO₄) dissociate completely in water, so [H₊] equals the acid concentration. Weak acids (acetic acid, citric acid, carbonic acid) dissociate only partially, so [H₊] is much lower than the acid concentration. For example, 0.1 M HCl gives pH = 1.00, while 0.1 M acetic acid (K_a = 1.8 × 10⁻⁵) gives pH = 2.87. This calculator accepts the actual [H₊] you enter, so you must use the correct free ion concentration for your specific acid or base.

The calculator assumes water-based solutions at 25 °C where K_w = 1.0 × 10⁻¹⁴. For non-aqueous solvents or mixed-solvent systems, the ion product constant is entirely different (e.g., for ethanol, the equivalent constant is about 10⁻¹⁹µ), and the conventional pH scale does not apply. For such systems, use non-aqueous pH measurement methods and appropriate reference standards.

Many pH-relevant concentrations are expressed in scientific notation, such as 1.0 × 10⁻⁷ M for pure water. Splitting the input into a coefficient field and an optional exponent field lets you enter values exactly as they appear in textbooks and lab data. You can also enter the full decimal (e.g., 0.0000001) directly in the coefficient field and leave the exponent blank.

Use the Henderson–Hasselbalch equation: pH = pK_a + log([A⁻]/[HA]). For example, a buffer containing 0.1 M acetic acid (pK_a = 4.76) and 0.2 M sodium acetate gives pH = 4.76 + log(0.2/0.1) = 5.06. This calculator works with the final [H₊] concentration. If you know the buffer’s exact [H₊], enter it directly. For preparing buffer stock solutions, use the Stock Solution Calculator.

Calibrate your pH meter daily for routine use and before every critical measurement series. Use at least two buffer solutions bracketing your expected sample pH (typically pH 4.00, 7.00, and 10.00). Replace buffers monthly and always use fresh aliquots to avoid contamination. If you notice slow response or drifting readings, clean the electrode and recalibrate. For high-precision work (< ±0.02 pH), calibrate immediately before each measurement.

Pure water undergoes autoionization: H₂O ⇌ H⁺ + OH⁻. At 25 °C, the equilibrium constant K_w = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴. Since pure water has equal concentrations of H⁺ and OH⁻, [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ M. Taking the negative logarithm: pH = −log(1.0 × 10⁻⁷) = 7.00. Pure water exposed to air absorbs CO₂ and forms carbonic acid, lowering its pH to about 5.5–6.0.

The calculator is most accurate for dilute solutions (below 0.1 M) where activity coefficients are close to 1. For concentrated acids above 1 M, the actual (measured) pH differs from −log[H⁺] because activity coefficients deviate significantly from 1. For example, 10 M HCl has an activity-based pH of about −1.0, but the direct calculation gives −log(10) = −1.0. For concentrated solutions, experimental measurement with a properly calibrated pH meter is recommended.

Normal arterial blood pH is 7.35–7.45, tightly regulated by the bicarbonate buffering system and respiratory control of CO₂ levels. Venous blood pH is slightly lower at 7.31–7.41. Blood pH below 7.35 is acidosis, and above 7.45 is alkalosis. Life-threatening conditions occur when blood pH falls below 6.8 or rises above 7.8. The bicarbonate system (H₂CO₃/HCO₃⁻, pK_a = 6.35) maintains this narrow range through the equilibrium: CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻.

CO₂ dissolves in water to form carbonic acid (H₂CO₃), which dissociates to release H⁺: CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻. This is why open containers of pure water or dilute buffers have pH 5.5–6.0 rather than 7.0. The same process causes ocean acidification: atmospheric CO₂ dissolves in seawater, lowering pH from 8.2 to 8.1 since the Industrial Revolution. In laboratory measurements, minimize sample exposure to air for accurate pH readings of unbuffered solutions.

pH test strips and indicator solutions provide semi-quantitative pH estimates.
Universal indicator changes color across the full pH range (red at pH 0–3, orange/yellow at pH 4–6, green at pH 7, blue at pH 8–10, purple at pH 11–14).
For specific ranges, single indicators are more precise: phenolphthalein (colorless at pH < 8.3, pink at pH > 10.0), bromothymol blue (yellow at pH < 6.0, blue at pH > 7.6).
Litmus paper gives only a binary acid/base classification.

Test strips typically have ±0.5 pH unit accuracy.

The calculator flags results for quick interpretation: a warning indicator appears for acidic solutions (pH < 7), and a neutral or basic result shows an ok indicator. This is a visual helper designed for rapid scanning of results, especially useful when processing multiple samples in a laboratory or educational setting.

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