Hydroxide Ion Concentration Calculator

Calculate the hydroxide ion [OH-] concentration from the pOH of a solution to quickly determine the alkalinity of basic mixtures.

Initial Concentration of Weak Base (M)

Base Dissociation Constant (Kb)

Equilibrium [OH⁻]

1.3416e-3

Solution pOH2.87

Solution pH11.13

Degree of Ionization1.34

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The Alkaline Equilibrium

While acids donate protons, weak bases (like ammonia or organic amines) act as proton acceptors. Because they are "weak," they do not dissociate completely. Instead, they must steal a hydrogen proton from a water molecule in a reversible equilibrium reaction, forcing the water to leave behind a hydroxide ion ([OH⁻]).

To calculate exactly how many hydroxide ions are produced in a basic solution, chemists use the Base Dissociation Constant (Kb).

Solving for [OH⁻]

Just like with weak acids, calculating the exact equilibrium state requires an ICE (Initial, Change, Equilibrium) table. However, if the base is weak enough that it ionizes less than 5%, we can bypass the complex quadratic math using the standard approximation formula.

The Equation

\begin{aligned} [OH⁻] = \sqrt{K_b \times C} \end{aligned}

Where:

[OH⁻]=

Equilibrium Hydroxide Ion Concentration (Molar)

K_b

Base Dissociation Constant

Initial Concentration of Weak Base

Example Calculation

You prepare a 0.5 M solution of Ammonia ( $K_b = 1.8 \times 10^{-5}$ ). To find the [OH⁻]:

Multiply $K_b$ by Concentration: $(1.8 \times 10^{-5}) \times 0.5 = 9.0 \times 10^{-6}$
Take the square root: $\sqrt{9.0 \times 10^{-6}}$
Result: $3.0 \times 10^{-3}$ M

The solution has an [OH⁻] of $0.003$ M. From here, you can easily calculate the pOH, and subsequently, the pH of the basic mixture.

Frequently Asked Questions

A weak base is a chemical compound that does not fully ionize in water. Unlike strong bases (like NaOH) which release 100% of their hydroxide ions instantly, weak bases establish an equilibrium where only a small fraction of molecules react with water.

First, take the negative logarithm of the [OH⁻] to find the pOH (pOH = -log[OH⁻]). Then, subtract that value from 14 at standard room temperature (pH = 14 - pOH).

No. The approximation formula only applies to weak bases that reach equilibrium. For a strong base like Potassium Hydroxide (KOH), the [OH⁻] is simply exactly equal to the initial concentration of the base.

If you only have the Ka of the conjugate acid, you can easily calculate the Kb. At 25°C, the product of Ka and Kb must equal exactly 1.0 x 10^-14 (Kw). Simply divide Kw by the Ka to find your Kb.

As a solution becomes extremely dilute, a larger percentage of the base ionizes. Once ionization exceeds 5%, the mathematical assumption that 'Initial Concentration ≈ Equilibrium Concentration' becomes invalid, requiring the full quadratic formula.

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