pOH Calculator
pOH is a measure of the hydroxide ion concentration [OH⁻] in a solution, representing its alkalinity or basicity. It's the logarithmic counterpart to pH, which measures hydrogen ion concentration [H⁺]. The "p" stands for "power" or "potenz" (German for power), and OH represents hydroxide ions.
pOH helps quantify basicity, predict chemical reactions, calculate equilibrium constants, and determine suitable conditions for biological and industrial processes. It's essential in acid-base chemistry, buffer solutions, and titration calculations.
Key pOH concepts:
- Logarithmic scale: pOH = -log₁₀[OH⁻] (like pH = -log₁₀[H⁺])
- Inverse relationship: Higher [OH⁻] → lower pOH → more basic
- Complementary to pH: pH + pOH = pKw (14 at 25°C)
- Temperature dependent: Kw changes with temperature
- Ion product constant: [H⁺]×[OH⁻] = Kw = 1.0×10⁻¹⁴ at 25°C
This calculator finds any acid-base parameter when you know one value:
- Find pOH: Enter [OH⁻] → Get pOH = -log₁₀[OH⁻]
- Find pH: Enter [H⁺] → Get pH = -log₁₀[H⁺]
- Find [OH⁻]: Enter pOH → Get [OH⁻] = 10⁻ᵖᴼᴴ
- Find [H⁺]: Enter pH → Get [H⁺] = 10⁻ᵖᴴ
The calculator provides:
- Complete acid-base profile from one input
- Visual pH/pOH scale with indicator position
- Solution classification (acidic/neutral/basic)
- Ion concentration ratios and comparisons
- Common solution presets for reference
- Multiple unit support (M, mM, µM)
Reference pOH and pH values of common solutions at room temperature (25°C):
| Solution | pH | pOH | [H⁺] (M) | [OH⁻] (M) | Acid-Base Type |
|---|---|---|---|---|---|
| Battery Acid | 0.5 | 13.5 | 3.16×10⁻¹ | 3.16×10⁻¹⁴ | Very Strong Acid |
| Lemon Juice | 2.0 | 12.0 | 1.00×10⁻² | 1.00×10⁻¹² | Strong Acid |
| Vinegar | 3.0 | 11.0 | 1.00×10⁻³ | 1.00×10⁻¹¹ | Weak Acid |
| Coffee | 5.0 | 9.0 | 1.00×10⁻⁵ | 1.00×10⁻⁹ | Weak Acid |
| Pure Water (25°C) | 7.0 | 7.0 | 1.00×10⁻⁷ | 1.00×10⁻⁷ | Neutral |
| Sea Water | 8.0 | 6.0 | 1.00×10⁻⁸ | 1.00×10⁻⁶ | Weak Base |
| Baking Soda Solution | 9.0 | 5.0 | 1.00×10⁻⁹ | 1.00×10⁻⁵ | Weak Base |
| Milk of Magnesia | 10.5 | 3.5 | 3.16×10⁻¹¹ | 3.16×10⁻⁴ | Strong Base |
| Household Ammonia | 11.5 | 2.5 | 3.16×10⁻¹² | 3.16×10⁻³ | Strong Base |
| 1M NaOH | 14.0 | 0.0 | 1.00×10⁻¹⁴ | 1.00 | Very Strong Base |
pOH 0-2: Very strong base (1M NaOH, concentrated KOH)
pOH 3-5: Strong base (ammonia, baking soda solution)
pOH 6-8: Weak base to neutral (sea water, blood, milk)
pOH 9-11: Weak acid (coffee, rainwater, urine)
pOH 12-14: Strong acid (vinegar, lemon juice, stomach acid)
Below are answers to frequently asked questions about pOH calculations:
Use the formula: pOH = -log₁₀[OH⁻] where [OH⁻] is in moles per liter (M).
- Measure [OH⁻]: Determine hydroxide concentration (e.g., 0.001 M)
- Take logarithm: log₁₀(0.001) = -3
- Apply negative sign: -(-3) = 3
- Result: pOH = 3.0
For [OH⁻] = 2.5×10⁻⁵ M: log₁₀(2.5×10⁻⁵) = log₁₀(2.5) + log₁₀(10⁻⁵) = 0.3979 - 5 = -4.6021 → pOH = 4.60
Calculator tip: Use the "log" button on your calculator. For [OH⁻] = 0.002 M: Press 0.002, then "log", then "×(-1)" or "±" to change sign.
Common concentration unit conversions for acid-base calculations:
1 M = 1000 mM = 1,000,000 µM
1 mM = 0.001 M = 1000 µM
1 µM = 0.000001 M = 0.001 mM
1 mol/L = 1 M = 1000 mmol/L
To convert: Multiply or divide by powers of 1000
Example conversions:
0.005 M = 5 mM = 5000 µM
250 µM = 0.00025 M = 0.25 mM
Our calculator handles all conversions automatically based on your selected units.
pOH monitoring is critical in water treatment for disinfection, corrosion control, and precipitation prevention.
| Process | Optimal pOH | Purpose | Consequences if Wrong |
|---|---|---|---|
| Chlorination | 6.0-7.0 | Maximize disinfectant efficiency | Poor disinfection or chlorine waste |
| Coagulation | 5.5-6.5 | Optimal floc formation | Poor turbidity removal |
| Corrosion Control | 7.0-8.5 | Prevent pipe corrosion | Metal leaching, pipe damage |
| Softening | 10.0-10.5 | Precipitate hardness ions | Scale formation, inefficient softening |
| Swimming Pools | 6.5-7.5 | Comfort and disinfection | Eye irritation, algae growth |
Key calculation: Treatment plants continuously monitor and adjust pOH (via pH) using acids/bases to maintain optimal conditions. Lime (Ca(OH)₂) raises pOH (lowers pH), while CO₂ injection lowers pOH (raises pH).
Living organisms maintain strict pOH/pH ranges through buffer systems:
- Human blood: pH 7.35-7.45 (pOH 6.55-6.65) - Critical for life
- Stomach acid: pH 1.5-3.5 (pOH 10.5-12.5) - Protein digestion
- Pancreatic juice: pH 7.5-8.8 (pOH 5.2-6.5) - Neutralizes stomach acid
- Cellular cytoplasm: pH ~7.2 (pOH ~6.8) - Enzyme function
- Lysosomes: pH 4.5-5.0 (pOH 9.0-9.5) - Waste digestion
- Plant sap: pH 5.0-6.5 (pOH 7.5-9.0) - Nutrient transport
Buffer systems: Bicarbonate (blood), phosphate (cells), and protein buffers maintain stable pOH. Deviations cause acidosis/alkalosis (blood) or enzyme denaturation. Medical tests often measure pOH/pH to diagnose disorders.
The relationship comes from the water autoionization constant Kw = [H⁺][OH⁻] = 1.0×10⁻¹⁴ at 25°C.
Kw = [H⁺][OH⁻] = 1.0×10⁻¹⁴
Take -log₁₀ of both sides:
-log(Kw) = -log([H⁺][OH⁻])
pKw = -log[H⁺] + (-log[OH⁻])
14 = pH + pOH
Since Kw varies with temperature, pKw = pH + pOH also varies: 13.99 at 20°C, 13.26 at 50°C, 12.26 at 100°C.
Practical implication: You only need to measure either pH or pOH to know both. In neutral water at 25°C: [H⁺] = [OH⁻] = 10⁻⁷ M, so pH = pOH = 7. Adding acid increases [H⁺] (lowers pH) and decreases [OH⁻] (raises pOH), keeping sum = 14.
Weak acids/bases partially dissociate, requiring equilibrium calculations using Ka or Kb.
| Type | Formula | Example Calculation | Approximation |
|---|---|---|---|
| Weak Acid | [H⁺] = √(Ka×C) pOH = 14 + log[H⁺] | 0.1M CH₃COOH (Ka=1.8×10⁻⁵) [H⁺]=√(1.8×10⁻⁶)=1.34×10⁻³ pH=2.87, pOH=11.13 | Valid when C≫Ka |
| Weak Base | [OH⁻] = √(Kb×C) pOH = -log[OH⁻] | 0.1M NH₃ (Kb=1.8×10⁻⁵) [OH⁻]=√(1.8×10⁻⁶)=1.34×10⁻³ pOH=2.87, pH=11.13 | Valid when C≫Kb |
| Buffer Solution | Henderson-Hasselbalch pOH = pKb + log([acid]/[base]) | NH₃/NH₄⁺ buffer pKb=4.75, [acid]=[base] pOH=4.75, pH=9.25 | Exact for buffers |
Key concept: For weak electrolytes, use equilibrium constants (Ka, Kb) not just concentration. Strong acids/bases (HCl, NaOH) fully dissociate: [H⁺] or [OH⁻] = concentration. Weak ones (acetic acid, ammonia) require equilibrium calculations.