pH Calculator,
find pH, pOH, [H⁺] and [OH⁻].

Field guide

pH, pOH, and the chemistry of acidic and alkaline solutions.

What is pH?

pH is a logarithmic measure of the concentration of hydrogen ions (H⁺) in an aqueous solution. The term stands for potential of hydrogen, and it was introduced by the Danish chemist Søren Peder Lauritz Sørensen in 1909. The formula is:

where [H⁺] is the molar concentration of hydrogen ions in mol/L (moles per litre). Because pH is a negative logarithm, a lower pH means a higher hydrogen ion concentration — and a more acidic solution. A solution with [H⁺] = 0.001 mol/L (10⁻³) has a pH of 3; a solution with [H⁺] = 0.0000001 mol/L (10⁻⁷) has a pH of 7.

The logarithmic scale: what each unit means

Because pH is logarithmic, each whole-number change represents a tenfold change in hydrogen ion concentration. A solution at pH 3 has ten times more H⁺ than a solution at pH 4, and one hundred times more than pH 5. This is why small changes in pH can represent large chemical differences. Blood pH, for instance, is tightly regulated between 7.35 and 7.45 — a range that looks tiny on the scale but corresponds to meaningful changes in ion concentration that affect enzyme activity throughout the body.

What is pOH?

pOH is the equivalent measure for hydroxide ions (OH⁻):

The critical relationship connecting pH and pOH comes from the ionic product of water (K_w). At 25 °C, pure water partially dissociates into H⁺ and OH⁻ ions with an equilibrium constant:

Taking the negative logarithm of both sides gives the fundamental identity:

This means that knowing any one of the four quantities (pH, pOH, [H⁺], [OH⁻]) immediately determines all three others. This calculator performs all four conversions simultaneously.

Acids, bases, and neutral solutions

The pH scale conventionally runs from 0 to 14, though values outside this range are chemically possible for extremely concentrated strong acids and bases:

• pH below 7 (acidic): The solution has more H⁺ than OH⁻. Common examples include stomach acid (pH 1.5–3.5), lemon juice (pH 2–3), vinegar (pH 2.4–3.4), and coffee (pH 4.5–5).
• pH = 7 (neutral): Equal concentrations of H⁺ and OH⁻ (each 10⁻⁷ mol/L at 25 °C). Pure distilled water at exactly 25 °C. Note that tap water is rarely exactly pH 7 due to dissolved CO₂ and mineral ions.
• pH above 7 (alkaline / basic): The solution has more OH⁻ than H⁺. Examples include blood (pH 7.35–7.45), seawater (pH 7.8–8.3), baking soda solution (pH 8.3), and bleach (pH 11–12).

Why does temperature affect pH?

The relationship pH + pOH = 14 is only exactly true at 25 °C. The value 14 comes from −log₁₀(K_w), and K_w changes with temperature. At 37 °C (body temperature), K_w ≈ 2.4 × 10⁻¹⁴, so the neutral point shifts to pH 6.81, not 7.00. This does not mean blood is acidic at body temperature — it means the definition of neutral changes slightly. For most practical calculations and this calculator, the 25 °C value is the standard.

Strong acids and weak acids: a key distinction

The pH calculation above assumes complete dissociation — that all of the acid or base has reacted with water to release H⁺ or OH⁻. This is only exactly true for strong acids (HCl, H₂SO₄, HNO₃) and strong bases (NaOH, KOH). Weak acids like acetic acid (vinegar) or carbonic acid only partially dissociate, so the pH of a 0.1 mol/L acetic acid solution is not pH 1.0 — it is about pH 2.87. Calculating pH for weak acids and bases requires using the Henderson–Hasselbalch equation and knowing the acid dissociation constant K_a.

This calculator is designed for strong acids and bases, or situations where you know the actual H⁺ or OH⁻ concentration directly (not just the concentration of the parent acid or base).

Practical applications

pH measurement and calculation appears across science, medicine, and industry:

• Medicine and physiology: Blood pH (7.35–7.45) is monitored in intensive care. Metabolic acidosis (pH below 7.35) and alkalosis (above 7.45) require rapid clinical intervention.
• Agriculture: Soil pH determines which nutrients are available to plants. Most crops prefer slightly acidic soil (pH 6–6.8). Acidic soils are limed to raise pH; alkaline soils can be treated with sulfur to lower it.
• Water treatment: Drinking water is regulated to pH 6.5–8.5. pH below 6.5 corrodes pipes; above 8.5, chlorination becomes less effective.
• Food science: pH is central to food safety (Clostridium botulinum does not grow below pH 4.6, the basis of home canning safety), fermentation (sourdough, yogurt, wine), and flavour chemistry (acids brighten flavours; alkalinity gives pretzels and ramen noodles their characteristic taste and colour).
• Laboratory chemistry: Buffer preparation, protein chemistry (enzymes are pH-sensitive), and titration calculations all depend on accurate pH work.

Disclaimer

This calculator uses the relationship pH + pOH = 14 at 25 °C and assumes complete ionic dissociation. Results are valid for strong electrolytes at dilute to moderate concentrations. For weak acids, buffers, or solutions at temperatures other than 25 °C, additional thermodynamic data is required.