pH Calculator
Convert between pH, pOH and ion concentration, and find the pH of strong acids, weak acids, bases and buffers. Every answer shows the working and where it sits on the pH scale.
How to Work Out pH
pH measures how acidic or alkaline a solution is by tracking the concentration of hydrogen ions in it. Because those concentrations span many orders of magnitude, the scale is logarithmic: every whole pH unit represents a tenfold change. Use the calculator for an instant answer, then read the formula behind it.
pH from hydrogen ion concentration
The definition of pH. Take the hydrogen ion concentration in moles per litre and apply the negative base-10 logarithm. A concentration of 0.001 mol/L gives a pH of 3, and a tenfold dilution raises the pH by exactly one unit.
pH = −log₁₀[H⁺]Concentration from pH, and the link to pOH
Reverse the logarithm to get the concentration back. Water itself sets the relationship between the two ions: at 25 °C the ion product Kw is 1.0 × 10⁻¹⁴, which is why pH and pOH always add up to 14 at that temperature.
[H⁺] = 10⁻ᵖᴴ · pH + pOH = 14Strong acids and strong bases
Strong acids such as HCl dissociate essentially completely, so the hydrogen ion concentration equals the acid concentration. For a strong base, find the hydroxide concentration first, take its pOH, then subtract from 14. Multiply by the number of ionisable protons for diprotic acids like H₂SO₄.
Acid: pH = −log₁₀(C) · Base: pH = 14 − (−log₁₀(C))Weak acids, weak bases and buffers
Weak acids only partly dissociate, so you need the dissociation constant Ka and an equilibrium calculation. For a buffer — a weak acid sitting alongside its conjugate base — the Henderson–Hasselbalch equation gives the pH directly, and the pH equals the pKa whenever the two concentrations match.
Ka = x²/(C − x) · pH = pKa + log₁₀([A⁻]/[HA])pH Calculator
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pH Formulas
The equations behind every calculation on this page, with a worked example for each. All values assume 25 °C.
| What you want | Formula | Worked example | When to use it |
|---|---|---|---|
| pH from hydrogen ion concentration | pH = −log₁₀[H⁺] | [H⁺] = 1 × 10⁻³ → pH 3.00 | Any solution where [H⁺] is known |
| Hydrogen ion concentration from pH | [H⁺] = 10⁻ᵖᴴ | pH 4 → 1 × 10⁻⁴ mol/L | Reversing a meter reading |
| pOH from hydroxide concentration | pOH = −log₁₀[OH⁻] | [OH⁻] = 1 × 10⁻² → pOH 2.00 | Alkaline solutions |
| pH from pOH | pH = 14 − pOH | pOH 2 → pH 12.00 | Converting between the two scales |
| Ion product of water | Kw = [H⁺][OH⁻] = 1 × 10⁻¹⁴ | [H⁺] = 1 × 10⁻⁵ → [OH⁻] = 1 × 10⁻⁹ | Finding the partner ion |
| pH of a strong acid | pH = −log₁₀(n × C) | 0.01 mol/L HCl → pH 2.00 | HCl, HNO₃, H₂SO₄ (n = 2) |
| pH of a strong base | pH = 14 + log₁₀(n × C) | 0.01 mol/L NaOH → pH 12.00 | NaOH, KOH, Ca(OH)₂ (n = 2) |
| pH of a weak acid | Ka = x² / (C − x), pH = −log₁₀(x) | 0.1 mol/L acetic acid → pH 2.87 | Partially dissociating acids |
| pH of a weak base | Kb = x² / (C − x), pH = 14 + log₁₀(x) | 0.1 mol/L ammonia → pH 11.13 | Ammonia, amines |
| Buffer pH (Henderson–Hasselbalch) | pH = pKa + log₁₀([A⁻]/[HA]) | pKa 4.76, equal amounts → pH 4.76 | Acid and conjugate base mixtures |
| pKa from Ka | pKa = −log₁₀(Ka) | Ka = 1.8 × 10⁻⁵ → pKa 4.74 | Reading data tables |
| Percent dissociation | (x ÷ C) × 100 | x = 1.34 × 10⁻³, C = 0.1 → 1.34% | Judging acid strength |
pH of Common Substances
Typical values for everyday liquids and laboratory solutions. Real samples vary with concentration, temperature and source, so treat these as approximate.
| Substance | Typical pH | Approx. [H⁺] (mol/L) | Classification |
|---|---|---|---|
| Battery acid | 0.5 | 3 × 10⁻¹ | Strongly acidic |
| Gastric acid | 1.5 – 3.5 | 3 × 10⁻² | Strongly acidic |
| Lemon juice | 2.0 – 2.6 | 1 × 10⁻² | Acidic |
| Vinegar | 2.4 – 3.4 | 4 × 10⁻³ | Acidic |
| Orange juice | 3.3 – 4.2 | 5 × 10⁻⁴ | Acidic |
| Black coffee | 4.8 – 5.2 | 1 × 10⁻⁵ | Weakly acidic |
| Rainwater | 5.0 – 5.6 | 1 × 10⁻⁵ | Weakly acidic |
| Milk | 6.5 – 6.8 | 3 × 10⁻⁷ | Weakly acidic |
| Pure water at 25 °C | 7.0 | 1 × 10⁻⁷ | Neutral |
| Human blood | 7.35 – 7.45 | 4 × 10⁻⁸ | Weakly alkaline |
| Seawater | 7.5 – 8.4 | 1 × 10⁻⁸ | Weakly alkaline |
| Baking soda solution | 8.3 | 5 × 10⁻⁹ | Alkaline |
| Milk of magnesia | 10.5 | 3 × 10⁻¹¹ | Alkaline |
| Household ammonia | 11.0 – 11.5 | 1 × 10⁻¹¹ | Strongly alkaline |
| Bleach | 12.5 – 13.0 | 3 × 10⁻¹³ | Strongly alkaline |
| Drain cleaner (NaOH) | 13.5 – 14.0 | 3 × 10⁻¹⁴ | Strongly alkaline |
pH FAQ
Answers to the questions people most often get stuck on when calculating pH.
pH is the negative base-10 logarithm of the hydrogen ion concentration in moles per litre, written pH = −log₁₀[H⁺]. To reverse it, [H⁺] = 10⁻ᵖᴴ. A solution with a hydrogen ion concentration of 0.01 mol/L therefore has a pH of 2, and diluting it tenfold moves the pH to 3.
At 25 °C, pH + pOH = 14. This follows from the ion product of water, Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴. Because Kw changes with temperature, the sum is only exactly 14 at 25 °C — at higher temperatures the neutral point sits below 7.
A strong acid dissociates almost completely, so nearly every molecule releases a hydrogen ion. A weak acid settles at an equilibrium where most molecules stay intact, producing far fewer ions. At 0.1 mol/L, hydrochloric acid sits near pH 1 while acetic acid is closer to pH 2.87 — roughly a 75-fold difference in hydrogen ion concentration.
Yes. The 0 to 14 range is a convention covering the concentrations met in most laboratory work, not a hard limit. Concentrated hydrochloric acid can show a negative pH and concentrated sodium hydroxide can exceed 14. At those concentrations, however, ion activity rather than simple concentration governs the true value, so calculated figures become unreliable.
Use the Henderson–Hasselbalch equation: pH = pKa + log₁₀([A⁻] ÷ [HA]), where [A⁻] is the conjugate base concentration and [HA] is the weak acid concentration. When the two are equal the logarithm is zero, so the pH equals the pKa. The equation holds best when the ratio stays between about 1:10 and 10:1.
Yes. The ion product of water rises with temperature, so neutral water at 50 °C has a pH of about 6.63 rather than 7.00. The water is still neutral, because hydrogen and hydroxide concentrations remain equal — only the numerical value of neutrality shifts. This is why pH meters include temperature compensation.
Because the scale is logarithmic, one unit is a tenfold change in hydrogen ion concentration. A solution at pH 4 is ten times more acidic than one at pH 5 and a hundred times more acidic than one at pH 6. This is why small shifts in the pH of blood or ocean water represent large chemical changes.