pH and Hydrogen Ion Concentration Calculator
pH is a logarithmic scale for the acidity of aqueous solutions, defined as pH = -log₁₀[H⁺]. Pick any one of pH, pOH, [H⁺] or [OH⁻], and the tool instantly derives the other three plus an acid / neutral / alkaline label (at 25 °C with Kw = 10⁻¹⁴).
Which value are you entering?
Enter a valid number (pH / pOH: -2 to 16; concentration must be > 0)
pH
7.00
NeutralPure water at 25 °C is exactly pH 7.00 with [H⁺] = [OH⁻] = 10⁻⁷ M. Note that human blood is intentionally slightly alkaline (pH 7.35–7.45), not neutral.
pH 0 – 14 scale
- pOH
- 7.00
- Hydrogen ion [H⁺]
- 1.00 × 10⁻⁷ M
- Hydroxide ion [OH⁻]
- 1.00 × 10⁻⁷ M
Common substances at 25 °C
Formula
pH = -log₁₀[H⁺] │ pOH = -log₁₀[OH⁻] │ pH + pOH = 14 │ [H⁺] · [OH⁻] = Kw = 1.0 × 10⁻¹⁴ (at 25 °C)
- · All conversions assume pure water at 25 °C with Kw = 1.0 × 10⁻¹⁴; as temperature rises Kw increases, so the neutral point shifts below pH 7 (e.g. ≈ 6.63 at 50 °C).
- · Every step of 1 on the pH scale is a 10× change in [H⁺]: a pH-3 solution has 1,000× the hydrogen-ion concentration of a pH-6 one.
- · Bands used here follow the common textbook convention: strong acid pH ≤ 1, acid 1–4, weak acid 4–7, neutral 7 ± 0.05, weak base 7–10, base 10–13, strong base > 13.
- · Concentrations are entered in mol/L (M). Convert mass-percent, g/L or normality to molarity first before using this tool.
- · The tool handles the simple [H⁺] ↔ pH mapping only — polyprotic acids (H₂SO₄, H₃PO₄), buffer solutions and Ka / Kb dissociation equilibria need the Henderson-Hasselbalch equation separately.
- · References: IUPAC Gold Book "pH" entry (doi:10.1351/goldbook.P04524); CRC Handbook of Chemistry and Physics, 97 th ed. (2016–17), §5.
Frequently asked
Is pH 7 always neutral?
pH 7 is neutral only for pure water at 25 °C, where Kw = 10⁻¹⁴ and [H⁺] = [OH⁻] = 10⁻⁷ M. Change the temperature and both Kw and the neutral pH shift — neutral water is pH ≈ 7.47 at 0 °C, 6.63 at 50 °C, and 6.14 at 100 °C. So pH 6.5 in hot water can already be neutral rather than acidic. Note also that human blood at pH 7.35–7.45 is intentionally slightly alkaline, not neutral.
Why does a pH change of 1 mean a 10× change in concentration?
Because pH = -log₁₀[H⁺] is a logarithmic scale that compresses concentrations from 10⁻¹ to 10⁻¹⁴ into the friendly 1–14 range. A 1-unit drop in pH means 10× more H⁺; pH 5 vs pH 2 is a 1,000× difference. So "acid rain at pH 4 versus typical rain at pH 5.6" sounds small but is actually a 40× hydrogen-ion difference.
Can I use this for buffer solutions or weak-acid Ka problems?
This tool only does the basic [H⁺] ↔ pH mapping, assuming the acid or base is fully ionised. For buffer solutions (weak acid HA with its conjugate base A⁻) use Henderson–Hasselbalch: pH = pKa + log([A⁻]/[HA]). For weak-acid equilibria, first solve the quadratic Ka = [H⁺][A⁻]/[HA] for [H⁺], then come back here to convert to pH. Polyprotic acids (H₂SO₄, H₃PO₄, H₂CO₃) each have their own Ka per dissociation step and must be handled stepwise.
Can pH be negative or greater than 14?
Yes. The 0–14 range is just the common case, not a hard limit. Concentrated 12 M HCl has [H⁺] ≈ 12 M, giving pH = -log(12) ≈ −1.1; fuming sulfuric acid (oleum) is even lower. On the basic side, saturated NaOH can reach pH > 15. This tool accepts -2 to 16, which covers virtually every secondary-school and first-year-university question. Note that outside 0–14 the "concentration = 10⁻ᵖᴴ" shortcut becomes inaccurate because activity coefficients drift from 1; academic chemistry switches to the H₀ Hammett acidity function for those regimes.
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