pH Calculator

Calculate pH from hydrogen ion concentration [H⁺], or find [H⁺] from pH. Also shows pOH, [OH⁻], and acidity classification.

Acid-Base

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How to Use

1
Select calculation direction
Choose whether to calculate pH from hydrogen ion concentration [H⁺] or find [H⁺] from a known pH value.
2
Enter your known value
Input the concentration in mol/L (M) or the pH value as a decimal number. The tool applies the definition pH = -log₁₀[H⁺] to compute the result.
3
Read the result and classification
The calculator returns the result along with the acid-base classification: strongly acidic (pH < 3), weakly acidic, neutral (pH 7), weakly basic, or strongly basic (pH > 11).

About

pH is one of the most fundamental parameters in chemistry, biology, medicine, and environmental science. The scale, introduced by Danish chemist Søren Pédersen in 1909, quantifies the acidity or basicity of an aqueous solution on a logarithmic scale derived from the molar concentration of hydrogen ions. Because the scale is logarithmic, each unit change in pH represents a tenfold change in [H⁺] — a solution at pH 4 is ten times more acidic than one at pH 5.

Understanding pH is critical across many disciplines. In biochemistry, enzyme activity, protein folding, and cellular metabolism all depend on maintaining pH within narrow ranges — human blood is tightly regulated between pH 7.35 and 7.45 by carbonate, phosphate, and protein buffer systems. In agriculture, soil pH determines nutrient availability and microbial community composition. In industrial chemistry, pH control affects reaction rates, product quality, and corrosion behavior.

This calculator uses the exact relationship pH = -log₁₀[H⁺] and its inverse [H⁺] = 10⁻ᵖᴴ to provide instant conversions. Whether you are calibrating laboratory equipment, balancing an aquarium, formulating a cosmetic product, or preparing for an exam, accurate pH calculation is the foundation of sound chemical reasoning.

FAQ

What is pH and why is the scale from 0 to 14?

pH is defined as the negative base-10 logarithm of the hydrogen ion concentration: pH = -log₁₀[H⁺]. At 25°C, pure water autoionizes to give [H⁺] = 1×10⁻⁷ mol/L, corresponding to pH 7. The practical scale runs from 0 to 14 because concentrations outside 1 mol/L to 10⁻¹⁴ mol/L are rarely encountered in aqueous solution, though the mathematical scale is theoretically unbounded in both directions.

How does temperature affect pH?

pH is temperature-dependent because the autoionization constant of water (Kw) changes with temperature. At 25°C, Kw = 1×10⁻¹⁴, giving neutral pH = 7. At 37°C (body temperature), Kw ≈ 2.4×10⁻¹⁴, shifting neutral pH to approximately 6.81. This means a blood pH of 7.4 at 37°C is still basic, not neutral. For precise measurements, always report the temperature at which pH was measured.

What is the difference between pH and pOH?

pH measures hydrogen ion concentration while pOH measures hydroxide ion concentration, defined as pOH = -log₁₀[OH⁻]. At 25°C, pH + pOH = 14, the value of pKw. This relationship means knowing either quantity at a given temperature allows calculation of the other. Strong bases are often easier to describe in terms of pOH: 0.1 mol/L NaOH gives pOH = 1, which means pH = 13.

Why do strong and weak acids give different pH values at the same concentration?

Strong acids (HCl, H₂SO₄, HNO₃) dissociate completely in water, so [H⁺] equals the initial acid concentration. Weak acids (acetic acid, carbonic acid) only partially dissociate; their equilibrium is described by an acid dissociation constant Ka. A 0.1 mol/L solution of HCl gives pH ≈ 1.0, while the same concentration of acetic acid (Ka = 1.8×10⁻⁵) gives pH ≈ 2.9 because only about 1.3% of the acid molecules ionize.

What is the Henderson-Hasselbalch equation?

The Henderson-Hasselbalch equation approximates the pH of a buffer solution: pH = pKa + log([A⁻]/[HA]), where pKa = -log(Ka), [A⁻] is the concentration of the conjugate base, and [HA] is the concentration of the weak acid. It is derived from the Ka expression and is most accurate when both acid and conjugate base concentrations are between 10⁻³ and 1 mol/L. Buffers are most effective within one pH unit of the pKa.

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