pOH
pOH is the negative base-10 logarithm of the hydroxide ion concentration [OH⁻] in a solution, serving as a measure of the basicity of an aqueous solution at a given temperature. At 25 °C, pH and pOH are complementary: they always sum to 14, making pOH a convenient way to express basic conditions. pOH is particularly useful in analytical chemistry when working with alkaline solutions, as it directly reflects the concentration of OH⁻ ions produced by bases.
Key Formula
pOH = -log₁₀[OH⁻] and pH + pOH = 14
LaTeX: \text{pOH} = -\log_{10}[\text{OH}^-] \quad \text{and} \quad \text{pH} + \text{pOH} = 14
| Symbol | Meaning | Unit |
|---|---|---|
| pOH | Negative log of hydroxide ion concentration | dimensionless |
| [OH⁻] | Molar concentration of hydroxide ions | mol/L |
| 14 | pKw at 25 °C (sum of pH and pOH) | dimensionless |
Worked Example
Problem
A cleaning solution has [OH⁻] = 0.01 mol/L. Calculate the pOH and pH of the solution at 25 °C.
Solution
Step 1: Calculate pOH: pOH = -log[OH⁻] = -log(0.01) = -log(10⁻²) = 2 Step 2: Use the relationship pH + pOH = 14 Step 3: pH = 14 - pOH = 14 - 2 = 12 Step 4: pH = 12 confirms the solution is strongly basic.
Answer
pOH = 2; pH = 12
Relationship Between [OH⁻], pOH, and pH at 25 °C
| [OH⁻] (mol/L) | pOH | pH | Classification |
|---|---|---|---|
| 1 | 0 | 14 | Strongly basic |
| 0.1 | 1 | 13 | Strongly basic |
| 0.001 | 3 | 11 | Moderately basic |
| 1 × 10⁻⁷ | 7 | 7 | Neutral |
| 1 × 10⁻¹⁰ | 10 | 4 | Acidic |
| 1 × 10⁻¹⁴ | 14 | 0 | Strongly acidic |
Interactive Tools
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Related Terms
pH Scale
The pH scale is a logarithmic measure of the hydrogen ion concentration [H⁺] in a solution, ranging from 0 to 14 at 25 °C, where values below 7 indicate acidic conditions, 7 is neutral, and above 7 is basic or alkaline. Introduced by Danish chemist Søren Peder Lauritz Sørensen in 1909, the scale compresses a trillion-fold range of H⁺ concentrations into a convenient 0–14 range. pH measurement is critical in agriculture, biology, medicine, food science, and environmental monitoring.
Water Ionization Constant (Kw)
The water ionization constant (Kw) is the equilibrium constant for the self-ionisation of water, defined as the product of the molar concentrations of hydrogen ions and hydroxide ions in pure water at a given temperature. At 25 °C, Kw = 1.0 × 10⁻¹⁴ mol²/L², which establishes the fundamental link between pH and pOH. Kw increases with temperature (water ionises more at higher temperatures), so the neutral pH shifts below 7 at elevated temperatures.
Base (Chemistry)
A base is a substance that accepts protons (H⁺ ions) from an acid or donates electron pairs, producing hydroxide ions (OH⁻) in aqueous solution, a bitter taste, slippery feel, and a pH above 7. Bases neutralise acids in reactions that form salt and water, making them essential in biological systems such as blood buffering, as well as in industrial processes like soap making. Common examples include sodium hydroxide (NaOH), ammonia (NH₃), and sodium bicarbonate (NaHCO₃).
The "p" comes from German "Potenz" (power) and "OH" represents the hydroxide ion (oxygen-hydrogen). The notation mirrors the pH convention introduced by Sørensen and was developed as an analogue measure for alkaline solutions.