pOH Calculator: Enter [OH⁻] Concentration, Get pOH & pH Instantly
This pOH calculator converts a hydroxide ion concentration [OH⁻] (in mol/L) into a pOH value and its equivalent pH using the two fundamental equations of aqueous acid–base chemistry: pOH = −log₁₀[OH⁻] and pH + pOH = 14 (at 25 °C). pOH is the negative base-10 logarithm of the hydroxide concentration — the mirror image of pH. A low pOH (below 7) signals a basic solution; a high pOH (above 7) signals an acidic one. Use this tool in chemistry coursework, lab titrations, water quality testing, and industrial process control.
pOH = −log₁₀([OH⁻]). At 25 °C, pH + pOH = 14. Example: [OH⁻] = 0.001 mol/L → pOH = 3, pH = 11 (moderately basic). Lower pOH means more basic; higher pOH means more acidic.
When to use this calculator
- Determining whether a cleaning solution (e.g., bleach, ammonia) is safe to handle based on its hydroxide concentration and resulting pOH/pH.
- Verifying the alkalinity of swimming pool or drinking water, where EPA guidelines recommend pH 7.2–7.8 and any pOH deviation signals over-treatment with base.
- Titration endpoint verification in general chemistry labs: calculating expected pOH after adding a known volume of NaOH to a weak acid solution.
- Industrial wastewater treatment: confirming that effluent [OH⁻] is low enough to meet EPA discharge limits (pH 6–9) before releasing water.
Worked Example: 0.001 mol/L NaOH
- Given: [OH⁻] = 0.001 mol/L (NaOH fully dissociates, so [OH⁻] = molar concentration)
- Step 1 — pOH = −log₁₀(0.001) = −log₁₀(10⁻³) = 3.00
- Step 2 — pH = 14 − pOH = 14 − 3 = 11.00
- Result: pOH = 3, pH = 11 → moderately basic (comparable to household ammonia)
How it works
3 min readHow pOH Is Calculated
The two core equations are:
pOH = −log₁₀([OH⁻]) # [OH⁻] in mol/L (molarity)
pH = 14 − pOH # valid at 25 °C (298 K)The constant 14 is the pKw of water at 25 °C — the negative log of water's ion-product constant Kw = 1.0 × 10⁻¹⁴ mol²/L². Both relationships derive directly from:
Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ (at 25 °C)
pH + pOH = pKw = 14> ⚠️ Temperature note: pKw changes with temperature. At 37 °C (body temperature) pKw ≈ 13.62, so pH + pOH ≈ 13.62, not 14. This calculator uses the standard 25 °C value.
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Quick-Reference Table: [OH⁻] → pOH → pH
| [OH⁻] (mol/L) | pOH | Equivalent pH | Classification |
|---|---|---|---|
| 1.0 | 0.00 | 14.00 | Extremely basic |
| 0.1 | 1.00 | 13.00 | Strongly basic |
| 0.01 | 2.00 | 12.00 | Strongly basic |
| 0.001 | 3.00 | 11.00 | Moderately basic |
| 1 × 10⁻⁴ | 4.00 | 10.00 | Moderately basic |
| 1 × 10⁻⁵ | 5.00 | 9.00 | Mildly basic |
| 1 × 10⁻⁶ | 6.00 | 8.00 | Slightly basic |
| 1 × 10⁻⁷ | 7.00 | 7.00 | Neutral |
| 1 × 10⁻⁸ | 8.00 | 6.00 | Slightly acidic |
| 1 × 10⁻⁹ | 9.00 | 5.00 | Mildly acidic |
| 1 × 10⁻¹⁰ | 10.00 | 4.00 | Moderately acidic |
Each 10× decrease in [OH⁻] raises pOH by 1 and lowers pH by 1.
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Step-by-Step Examples
Example 1 — 0.001 mol/L NaOH
NaOH is a strong base that fully dissociates, so [OH⁻] = 0.001 mol/L exactly.
pOH = −log₁₀(0.001) = −log₁₀(10⁻³) = 3.00
pH = 14 − 3.00 = 11.00 → Moderately basicA pH of 11 is consistent with household ammonia (pH ≈ 11–12) and is corrosive to skin with prolonged contact.
Example 2 — Bleach solution at [OH⁻] = 0.01 mol/L
Sodium hypochlorite solutions are alkaline. With [OH⁻] = 0.01 mol/L:
pOH = −log₁₀(0.01) = 2.00
pH = 14 − 2.00 = 12.00 → Strongly basicEPA classifies solutions above pH 12.5 as corrosive hazardous waste (40 CFR §261.22).
Example 3 — Aquarium water at [OH⁻] = 3.16 × 10⁻⁶ mol/L
Many freshwater fish thrive at pH 8.0–8.5. At [OH⁻] = 3.16 × 10⁻⁶ mol/L:
pOH = −log₁₀(3.16 × 10⁻⁶) = 5.50
pH = 14 − 5.50 = 8.50 → Mildly basic, optimal for many species---
Common Errors to Avoid
1. Forgetting the negative sign — pOH is minus the log. log₁₀(0.001) = −3, so pOH = +3, not −3. A negative pOH only occurs when [OH⁻] > 1 mol/L (very concentrated bases).
2. Using pH + pOH = 14 at non-standard temperatures — At 0 °C, pKw ≈ 14.94; at 60 °C, pKw ≈ 13.02. Applying the 14 constant outside 20–25 °C range introduces meaningful error in precise lab work.
3. Inputting [OH⁻] in units other than mol/L — The formula requires molarity (mol/L). If you measure in mmol/L (mM), divide by 1000 first. For example, 5 mM OH⁻ = 0.005 mol/L → pOH = 2.30.
4. Confusing [OH⁻] with [H⁺] — Some instruments report [H⁺] directly. You cannot plug [H⁺] into the pOH formula. Either compute pH = −log[H⁺] first, then pOH = 14 − pH, or calculate [OH⁻] = Kw / [H⁺].
5. Assuming strong-base dissociation for weak bases — For ammonia (NH₃), [OH⁻] ≠ initial concentration. You must use the base dissociation constant Kb = 1.8 × 10⁻⁵ and solve the equilibrium expression before plugging into the pOH formula.
Frequently asked questions
What is pOH and how does it differ from pH?
pH measures hydrogen ion concentration [H⁺] (pH = −log[H⁺]), while pOH measures hydroxide ion concentration [OH⁻] (pOH = −log[OH⁻]). Both scales are linked by pH + pOH = 14 at 25 °C. A low pOH (e.g., 2) means high [OH⁻] and a strongly basic solution; a high pOH (e.g., 12) means low [OH⁻] and an acidic solution — the opposite of how pH behaves.
Why does pH + pOH equal exactly 14?
At 25 °C, the ion-product constant of water is Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ mol²/L². Taking the negative log of both sides gives −log(Kw) = −log[H⁺] + (−log[OH⁻]), or pKw = pH + pOH = 14. This value is specific to 25 °C; at 37 °C pKw drops to about 13.62, and at 0 °C it rises to about 14.94.
How do I calculate pOH if I only know the pH?
Rearrange the relationship: pOH = 14 − pH. For example, a solution with pH 9 has pOH = 14 − 9 = 5.00. From there, [OH⁻] = 10^(−5) = 0.00001 mol/L. This is the most common route in chemistry problems where pH is given but [OH⁻] is needed.
How do I find [OH⁻] if I only know the pH?
Use two steps: (1) pOH = 14 − pH, then (2) [OH⁻] = 10^(−pOH). For a solution with pH 9: pOH = 5, so [OH⁻] = 10⁻⁵ = 0.00001 mol/L. Alternatively, use Kw directly: [OH⁻] = Kw / [H⁺] = (1 × 10⁻¹⁴) / (10⁻⁹) = 10⁻⁵ mol/L.
Can pOH be negative or greater than 14?
Yes. pOH is negative when [OH⁻] > 1 mol/L — for example, 10 mol/L NaOH gives pOH = −1 and pH = 15. pOH exceeds 14 only in very concentrated strong acids where [OH⁻] is vanishingly small (e.g., 10 mol/L HCl → pOH ≈ 15, pH ≈ −1). These extremes are real but rarely encountered outside industrial chemistry.
Does this formula work for weak bases like ammonia?
Not directly. Weak bases don't fully dissociate, so [OH⁻] ≠ initial concentration. For ammonia (NH₃), you must first solve the equilibrium: Kb = [NH₄⁺][OH⁻] / [NH₃] = 1.8 × 10⁻⁵. For a 0.1 mol/L NH₃ solution, [OH⁻] ≈ √(1.8 × 10⁻⁵ × 0.1) ≈ 1.34 × 10⁻³ mol/L → pOH ≈ 2.87, pH ≈ 11.13. Then and only then do you apply the pOH formula.
What pOH range is considered safe for drinking water?
The EPA Secondary Drinking Water Regulations recommend pH 6.5–8.5, which corresponds to pOH ≈ 5.5–7.5. The WHO Guidelines for Drinking-water Quality (4th ed.) note that pH outside 6.5–9.5 may indicate chemical contamination or corrosion issues. A pOH near 6–7 (pH 7–8) is ideal for most municipal supplies.
How do I convert pOH back to [OH⁻] concentration?
Reverse the logarithm: [OH⁻] = 10^(−pOH). For pOH = 4.5: [OH⁻] = 10^(−4.5) = 3.16 × 10⁻⁵ mol/L. For pOH = 2: [OH⁻] = 10^(−2) = 0.01 mol/L. This inverse operation is called the antilog and is available on any scientific calculator as the '10^x' or 'INV log' button.
How accurate is pOH = −log[OH⁻] at high ionic strength?
At high ionic strength (concentrated salt solutions), activity effects cause deviation. The rigorous formula is pOH = −log(a_OH⁻) = −log(γ[OH⁻]), where γ is the activity coefficient (Debye–Hückel equation). For dilute solutions (ionic strength < 0.01 mol/L), γ ≈ 1 and the simplified formula is accurate to ±0.01 pOH units. For seawater (ionic strength ≈ 0.7 mol/L), γ for OH⁻ can drop to ~0.6, causing ~0.22 pOH units of error if ignored.