pOH: the other side of the scale
Learn what pOH measures, how it relates to pH through pH + pOH = 14, and how to convert between hydroxide concentration and pH.
Quick answer
pOH = −log[OH⁻]: the negative logarithm of the hydroxide-ion concentration. At 25 °C, pH + pOH = 14.00, so either scale gives the other.
Example: For 0.0010 M NaOH (strong base), [OH⁻] = 0.0010 M, pOH = 3.00, and pH = 14.00 − 3.00 = 11.00.
What you will learn
- Define pOH as −log[OH⁻].
- Use pH + pOH = 14.00 (at 25 °C) to convert between the scales.
- Calculate pOH from hydroxide concentration and pH from pOH.
- Explain the Kw origin of the pH–pOH relationship.
The mirror image of pH
Just as pH compresses the huge range of [H⁺] into a tidy scale, pOH does the same for hydroxide:
Low pOH means basic (lots of OH⁻); high pOH means acidic (little OH⁻). For 0.0010 M NaOH, a strong base that dissociates completely: [OH⁻] = 0.0010 M, so pOH = −log(1.0 × 10⁻³) = 3.00.
Why they sum to 14
Water's autoionization, Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C, ties the two scales together. Take the negative log of both sides:
So the NaOH solution above has pH = 14.00 − 3.00 = 11.00. One scale always implies the other — reporting both is redundant but sometimes convenient.
When pOH is the easier tool
pOH shines in base-centered problems. Titrating an acid with NaOH? Track pOH as the base is added — it falls smoothly while pH jumps. Working with Kb equilibria? [OH⁻] is the natural variable, so pOH keeps the numbers small.
Conversion drill: [OH⁻] = 2.5 × 10⁻⁵ M → pOH = −log(2.5 × 10⁻⁵) = 4.60 → pH = 14.00 − 4.60 = 9.40. And backward: pOH 5.20 → [OH⁻] = 10^(−5.20) = 6.3 × 10⁻⁶ M.
Temperature changes the sum
The 14.00 sum is a 25 °C value. Because Kw grows with temperature (water ionizes more when hot), pKw shrinks: at 100 °C, Kw ≈ 5.1 × 10⁻¹³, so pH + pOH ≈ 12.29 and neutral pH is ~6.14. In problems at nonstandard temperatures, always recompute from the given Kw rather than assuming 14.
Variables and units
| Symbol | Name | Unit | Meaning |
|---|---|---|---|
| Hydroxide-ion concentration | M | Molar concentration of hydroxide ions |
Common mistakes
Every aqueous solution has both a pH and a pOH — they describe the same solution from opposite ends. A pH 3 solution has pOH 11; a pOH 2 solution has pH 12. The pair always describes the full picture.
High pOH means low [OH⁻] — the opposite. pOH 12 is barely basic (it is pH 2's mirror). Low pOH = basic, high pOH = acidic, exactly mirroring pH.
pOH is computed, not measured: find [OH⁻] (or pH) and convert. pH meters measure H⁺ activity; pOH comes from the Kw relationship.
Real-world applications
Base titrations: pOH is the natural variable when titrating with NaOH.
Industrial cleaning: alkaline process control often tracks pOH directly.
Biochemistry: enzyme activity versus hydroxide concentration.
Water treatment: lime softening calculations use hydroxide concentrations.
Frequently asked questions
Is pH + pOH always 14?
pH + pOH = 14.00 holds at 25 °C because Kw = 1.0 × 10⁻¹⁴. At other temperatures Kw differs, so the sum differs (about 12.3 at 100 °C). Always check the temperature before using 14.
Does pOH work like pH but for bases?
Same idea, mirrored: pOH = −log[OH⁻], lower pOH means more basic, and each unit is a tenfold change in [OH⁻]. pOH 3 means [OH⁻] = 10⁻³ M — a fairly strong base.
Why do pH and pOH add to 14?
pH = −log[H⁺] and pOH = −log[OH⁻], with [H⁺][OH⁻] = Kw. Taking −log of both sides of the Kw expression gives pH + pOH = pKw = 14.00. The sum rule is just the Kw equilibrium in logarithmic clothing.
Why do we use pH far more often than pOH?
Mostly convention and history — pH was defined first and became universal. But pOH is genuinely convenient for base-heavy problems: 0.010 M NaOH gives pOH = 2.00 immediately, then pH = 12.00.