Convert between pH, [H⁺], and [OH⁻]. Calculate hydrogen and hydroxide ion concentrations from pH or vice versa.
The hydrogen ion concentration [H⁺] is the master variable that defines the acidity or basicity of every aqueous solution. The pH scale, defined as pH = −log₁₀[H⁺], compresses the enormous range of possible hydrogen ion concentrations (from about 1 M in strong acid to 10⁻¹⁴ M in strong base) into a manageable 0–14 scale that chemists and biologists use daily.
Understanding the relationship between pH, [H⁺], and [OH⁻] is critical in virtually every branch of science. In biochemistry, enzyme activity depends on precise pH; a shift of just 0.1 pH units can alter reaction rates by 25% or more. In environmental science, even small changes in water pH affect aquatic ecosystem health. In medicine, blood pH must be maintained between 7.35 and 7.45, with deviations causing acidosis or alkalosis.
This calculator converts freely between pH, [H⁺], and [OH⁻], using the water autoionization constant Kw = [H⁺][OH⁻] = 10⁻¹⁴ at 25°C. Enter any one value and instantly get all the others, along with a visual pH scale showing where your solution falls and a reference table of common substances.
Converting between pH and molar concentrations requires logarithmic calculations that are easy to get wrong, especially with scientific notation. This calculator does it instantly and shows the context of where your solution sits on the acid-base spectrum. This hydrogen ion concentration calculator helps you compare outcomes quickly and reduce avoidable mistakes when making day-to-day care decisions. Use the estimate as a planning baseline and confirm final decisions with a qualified professional when risk is high.
pH = −log₁₀[H⁺]. pOH = −log₁₀[OH⁻]. pH + pOH = 14 (at 25°C). [H⁺] = 10^(−pH). [OH⁻] = Kw/[H⁺] = 10⁻¹⁴/[H⁺].
Result: [H⁺] = 0.01 M, [OH⁻] = 1×10⁻¹² M
[H⁺] = 10^(−2) = 0.01 M. pOH = 14 − 2 = 12. [OH⁻] = 10^(−12) = 1×10⁻¹² M. This is 100,000 times more acidic than pure water.
The pH concept was introduced by Danish chemist Søren Sørensen in 1909 while working at the Carlsberg Laboratory in Copenhagen. He needed a convenient way to express hydrogen ion concentrations during enzyme studies. The original definition used the hydrogen electrode; modern pH meters use glass electrodes calibrated against standard buffers.
Living systems maintain exquisitely tight pH control. Blood plasma has a pH of 7.35–7.45; deviation by just 0.2 units can be life-threatening. The stomach operates at pH 1.5–3.5 to denature proteins and kill bacteria. Lysosomes maintain pH ~4.5 for acidic hydrolase activity. Each organelle has its own optimal pH, maintained by proton pumps and buffers.
For strong acids that fully dissociate, [H⁺] equals the acid concentration: 0.01 M HCl has [H⁺] = 0.01 M and pH = 2.0. For weak acids, the calculation requires the Ka equilibrium: [H⁺] = √(Ka × C) for dilute solutions. This calculator works with the final [H⁺] regardless of whether it came from a strong or weak acid.
pH stands for "power of hydrogen" and equals −log₁₀[H⁺]. It measures the hydrogen ion activity in a solution on a scale typically from 0 to 14.
Yes. Concentrated strong acids can have negative pH (e.g., 10 M HCl has pH ≈ −1), and concentrated strong bases can exceed 14.
Because hydrogen ion concentrations span 14 orders of magnitude (from 1 to 10⁻¹⁴ M). A logarithmic scale makes these values practical to work with.
Kw is the ion product of water: Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C. It increases with temperature (about 10⁻¹³ at 60°C).
Yes. Water's Kw increases with temperature, so neutral pH drops below 7 at higher temperatures. At 37°C (body temperature), neutral pH is about 6.8.
pH measures acidity ([H⁺]), while pOH measures basicity ([OH⁻]). They always sum to pKw, which is 14 at 25°C.