Strong Acid and Base pH Calculator
Inputs
| Solute type | Strong acid |
|---|---|
| Concentration | 0.01 M |
| Ionizable H⁺ or OH⁻ per unit | 1 |
Strong Acid and Base pH Calculator
Find the pH and pOH of a strong acid or strong base from its molar concentration. Handles polyprotic acids and bases such as H₂SO₄ and Ca(OH)₂ through the H⁺/OH⁻ count, assuming complete dissociation at 25 °C.
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This is a strongly acidic solution — a low pH with a high hydrogen-ion concentration.
Details
Strong acids, strong bases, and complete dissociation
A strong acid or strong base ionizes completely when it dissolves in water. Every HCl molecule becomes one H⁺ and one Cl⁻; every NaOH unit becomes one Na⁺ and one OH⁻. Because nothing stays un-ionized, the concentration of the reactive ion is fixed directly by how much solute was added — no equilibrium constant is needed. That is what makes the pH of a strong acid or base a one-step logarithm rather than the quadratic that weak acids require.
The formula
For a strong acid, the hydrogen-ion concentration equals the number of ionizable protons per formula unit, n, times the molar concentration C:
[H+]=nCpH=−log10(nC)For a strong base, the same logic applies to hydroxide ions, and the pH follows from the ion product of water:
[OH−]=nCpOH=−log10(nC)pH=14−pOH| Symbol | Quantity | Unit |
|---|---|---|
| C | Concentration of the acid or base | mol/L (M) |
| n | Ionizable H⁺ or OH⁻ per formula unit | unitless |
| pH | Acidity measure | unitless |
| pOH | Basicity measure | unitless |
The relationship pH + pOH = 14 holds at 25 °C, where the ion product of water is Kw = [H⁺][OH⁻] = 1 × 10⁻¹⁴.
The count n for polyprotic acids and bases
The factor n accounts for solutes that release more than one reactive ion. Sulfuric acid, H₂SO₄, is diprotic and contributes two H⁺ per molecule, so a 0.005 M solution behaves as 0.01 M in hydrogen ions. Calcium hydroxide, Ca(OH)₂, and barium hydroxide, Ba(OH)₂, each release two OH⁻ per formula unit. Monoprotic acids such as HCl and HNO₃, and single-hydroxide bases such as NaOH and KOH, use n = 1.
| Solute | Type | n | 0.01 M gives |
|---|---|---|---|
| HCl | strong acid | 1 | pH 2.00 |
| H₂SO₄ | strong acid | 2 | pH 1.70 |
| NaOH | strong base | 1 | pH 12.00 |
| Ca(OH)₂ | strong base | 2 | pH 12.30 |
Sulfuric acid is only approximately strong in its second ionization, so treating it as fully diprotic is a good but not exact model at higher concentrations.
Worked example
What is the pH of 0.005 M sulfuric acid?
[H+]=nC=2×0.005=0.01 M pH=−log10(0.01)=2.00Then pOH = 14 − 2.00 = 12.00, and the hydroxide concentration is [OH⁻] = 10⁻¹² M. With a pH well below 7, the solution is strongly acidic.
Strong versus weak
Only a handful of acids are strong: HCl, HBr, HI, HNO₃, HClO₃, HClO₄, and H₂SO₄. The common strong bases are the group 1 hydroxides (LiOH, NaOH, KOH, RbOH, CsOH) and the heavier group 2 hydroxides (Ca(OH)₂, Sr(OH)₂, Ba(OH)₂). Everything else — acetic acid, citric acid, ammonia, and the like — is weak and ionizes only partially. For a weak acid the actual [H⁺] is much smaller than n·C and must be found from the acid-dissociation constant Ka. Using this calculator on a weak acid would badly overstate its acidity.
The dilute limit
The plain formula assumes the solute supplies essentially all of the H⁺ or OH⁻, ignoring the tiny amount water contributes on its own. That is safe for typical laboratory concentrations. Below about 10⁻⁶ M, though, water's self-ionization is no longer negligible, and a strong acid can never raise the pH above 7 no matter how dilute it becomes. A rigorous treatment solves [H⁺] = C + Kw/[H⁺] together with the charge balance; for very dilute solutions the true pH sits closer to neutral than −log₁₀(n·C) suggests.
Frequently Asked Questions (FAQ)
How do you calculate the pH of a strong acid?
A strong acid dissociates completely in water, so the hydrogen-ion concentration equals the number of ionizable protons times the acid concentration: [H⁺] = n·C. The pH is then the negative base-10 logarithm of that value, pH = −log₁₀(n·C). For example, 0.01 M HCl (n = 1) gives pH = −log₁₀(0.01) = 2. The seven common strong acids are HCl, HBr, HI, HNO₃, HClO₃, HClO₄, and H₂SO₄.
How do you find the pH of a strong base?
For a strong base you first find the pOH from the hydroxide concentration, pOH = −log₁₀(n·C), then convert to pH with pH = 14 − pOH at 25 °C. For example, 0.01 M NaOH (n = 1) has pOH = 2 and pH = 12. A 0.01 M solution of Ca(OH)₂ releases two OH⁻ per formula unit (n = 2), giving [OH⁻] = 0.02 M, pOH ≈ 1.70, and pH ≈ 12.30.
What do I enter for a diprotic acid like H₂SO₄?
Set the number of ionizable H⁺ (or OH⁻) per formula unit to match the solute. Sulfuric acid, H₂SO₄, releases two protons, so use 2: a 0.005 M solution then behaves as 0.01 M in H⁺ and gives pH 2. Likewise Ba(OH)₂ and Ca(OH)₂ release two hydroxide ions, so a base with those uses 2. Monoprotic acids (HCl, HNO₃) and single-hydroxide bases (NaOH, KOH) use 1.
Does this work for weak acids like acetic acid?
No. This calculator assumes complete dissociation, which is only accurate for strong acids and bases. Weak acids and bases (acetic acid, ammonia, and most organic acids) ionize only partially, so their actual [H⁺] is far lower than n·C and depends on the acid-dissociation constant Ka. For a weak acid, calculate [H⁺] from Ka and the concentration first, then take its negative logarithm.
Why is the pH not exactly 6 for 10⁻⁶ M HCl?
The formula pH = −log₁₀(n·C) ignores the H⁺ that water itself contributes. That contribution is negligible for ordinary concentrations, but near 10⁻⁶ M and below it becomes significant, and a strong acid can never push the pH above 7. A rigorous result comes from solving [H⁺] = C + Kw/[H⁺] together with charge balance; for very dilute solutions the true pH sits closer to 7 than the simple formula predicts.
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pH Calculator
Calculate pH, pOH, [H⁺], and [OH⁻] of a solution. Enter hydrogen-ion concentration, hydroxide concentration, pH, or pOH and get the other three at 25 °C.