pH Calculator
Enter a pH, pOH, [H⁺] or [OH⁻] value to convert it to all the others, or find the pH of a strong acid, strong base, weak acid or weak base with every step shown.
Scientific notation works: 1.8e-5 or 1.8×10^-5.
Assumes dilute aqueous solutions at 25 °C, where Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ and pH + pOH = 14.
How to use the pH Calculator
- Pick a mode: conversions and strong acids/bases, or weak acid/base.
- Choose what you know – pH, pOH, [H⁺], [OH⁻] or a strong acid/base concentration – and type the value.
- For weak acids or bases, pick a compound or enter Ka/Kb (or pKa/pKb) and the concentration.
- Press Calculate pH to see pH, pOH, both ion concentrations and the working.
- Read the scale to see whether the solution is acidic, neutral or basic, and Copy the result.
pH, pOH and concentration formulas
- pH = −log₁₀[H⁺] and [H⁺] = 10−pH
- pOH = −log₁₀[OH⁻] and [OH⁻] = 10−pOH
- At 25 °C, water's ion product is Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴, so pH + pOH = 14.
Example: a solution with pH 3.5 has [H⁺] = 10⁻³·⁵ = 3.16 × 10⁻⁴ mol/L, pOH = 10.5 and [OH⁻] = 3.16 × 10⁻¹¹ mol/L. Each pH unit is a tenfold change in [H⁺].
pH of weak acids and bases
A weak acid HA only partly dissociates: HA ⇌ H⁺ + A⁻, with Ka = [H⁺][A⁻] ÷ [HA]. If x is the amount that dissociates from a starting concentration C, then Ka = x² ÷ (C − x). Solving the quadratic gives
x = [H⁺] = (−Ka + √(Ka² + 4·Ka·C)) ÷ 2
For 0.1 M acetic acid (Ka = 1.8 × 10⁻⁵): x = 1.333 × 10⁻³ mol/L, so pH ≈ 2.88 and only 1.33% of the acid is ionized. The shortcut x ≈ √(Ka·C) works when less than 5% dissociates. Weak bases use Kb and [OH⁻] in the same way, then pH = 14 − pOH.
pH of common substances
| Substance | Typical pH |
|---|---|
| Battery acid | 0 – 1 |
| Stomach acid | 1.5 – 3.5 |
| Lemon juice | 2 – 2.6 |
| Vinegar | 2.4 – 3.4 |
| Coffee | 4.5 – 5 |
| Pure water (25 °C) | 7.0 |
| Blood | 7.35 – 7.45 |
| Seawater | about 8.1 |
| Baking soda solution | 8.3 – 9 |
| Household ammonia | 11 – 12 |
| Bleach | 12.5 – 13.5 |
Limits of these calculations
The formulas assume dilute solutions at 25 °C and use concentrations in place of activities. In concentrated solutions (above about 0.1 M for strong electrolytes) real pH can differ by a few tenths. Kw also changes with temperature – at 37 °C neutral water has a pH of about 6.8. Sulfuric acid's second proton is not fully released, so the "2 ions" option slightly overestimates its acidity.
Frequently asked questions
How do you calculate pH from [H+]?
Take the negative base-10 logarithm: pH = −log₁₀[H⁺]. For [H⁺] = 1 × 10⁻⁵ mol/L the pH is 5. For 2.5 × 10⁻⁴ mol/L the pH is about 3.60.
How do you get [H+] from pH?
Raise 10 to the power of minus the pH: [H⁺] = 10⁻ᵖᴴ. A pH of 7.4 means [H⁺] = 10⁻⁷·⁴ ≈ 3.98 × 10⁻⁸ mol/L.
Why do pH and pOH add up to 14?
In water at 25 °C, [H⁺] × [OH⁻] = Kw = 1.0 × 10⁻¹⁴. Taking −log₁₀ of both sides gives pH + pOH = 14. At other temperatures Kw changes, so the sum is slightly different.
What is the pH of 0.1 M acetic acid?
About 2.88. With Ka = 1.8 × 10⁻⁵, solving x² ÷ (0.1 − x) = Ka gives [H⁺] = 1.33 × 10⁻³ mol/L, so only about 1.3% of the acid molecules are ionized.
What is the difference between a strong and a weak acid?
A strong acid such as HCl dissociates almost completely in water, so [H⁺] equals its concentration. A weak acid such as acetic acid dissociates only partly, described by its acid dissociation constant Ka.
Can pH be negative or above 14?
Yes, for very concentrated strong acids or bases, such as 10 M HCl, though at those concentrations activities differ from concentrations and the simple formulas are less accurate.