Calculate the pH for a 0.75 M solution of potassium hypochlorite
Potassium hypochlorite, KClO, dissociates to give the hypochlorite ion, ClO–, which acts as a weak base in water. Use this interactive calculator to estimate pH, pOH, hydroxide concentration, and percent hydrolysis from concentration and acid dissociation data for hypochlorous acid.
Default values correspond to a common classroom setup at 25 C. Because KClO is the salt of a strong base and a weak acid, the resulting solution is basic.
Results
Enter or keep the default values, then click Calculate pH to view the worked answer for a 0.75 M potassium hypochlorite solution.
How to calculate the pH for a 0.75 M solution of potassium hypochlorite
To calculate the pH for a 0.75 M solution of potassium hypochlorite, you first identify what species in solution actually controls the acid-base chemistry. Potassium hypochlorite, written as KClO, is an ionic compound. In water it dissociates essentially completely into potassium ions, K+, and hypochlorite ions, ClO–. The potassium ion is the conjugate acid of a strong base and does not significantly affect pH, but the hypochlorite ion is the conjugate base of hypochlorous acid, HOCl, which is a weak acid. That means ClO– hydrolyzes in water to produce hydroxide ions, OH–, and this is why the solution becomes basic.
The key equilibrium is:
ClO- + H2O ⇌ HOCl + OH-
Since hydroxide is produced, the pH will be greater than 7. The calculation is a classic weak base equilibrium problem. Instead of being given a base dissociation constant directly, chemistry texts often provide the acid dissociation constant, Ka, for hypochlorous acid. That is not a problem, because Ka and Kb are connected by the familiar relation:
Ka × Kb = Kw
At 25 C, the ion product of water, Kw, is usually taken as 1.0 × 10-14. If we use a representative Ka for HOCl of 3.0 × 10-8, then:
Kb = 1.0 × 10^-14 / 3.0 × 10^-8 = 3.33 × 10^-7
Step by step solution for 0.75 M KClO
- Write the hydrolysis equation for hypochlorite:
ClO- + H2O ⇌ HOCl + OH-
- Set the initial concentration of ClO– equal to the salt concentration:
[ClO-]initial = 0.75 M
- Compute Kb from the Ka of HOCl:
Kb = Kw / Ka = 1.0 × 10^-14 / 3.0 × 10^-8 = 3.33 × 10^-7
- Let x represent the amount of OH– formed at equilibrium. Then:
Kb = x^2 / (0.75 – x)
- Use the weak base approximation if x is very small relative to 0.75:
x ≈ √(Kb × C) = √(3.33 × 10^-7 × 0.75)
- That gives:
[OH-] ≈ 5.00 × 10^-4 M
- Now convert to pOH:
pOH = -log(5.00 × 10^-4) ≈ 3.30
- Finally calculate pH:
pH = 14.00 – 3.30 = 10.70
Why the approximation works well
In weak acid and weak base problems, students are often taught to check whether the approximation is valid. Here the estimated hydroxide concentration is about 5.00 × 10-4 M. Compare that with the starting hypochlorite concentration of 0.75 M. The fraction reacted is only:
(5.00 × 10^-4 / 0.75) × 100 ≈ 0.067%
That is far below 5%, so the approximation is excellent. If you solve the quadratic exactly, the result differs by only a tiny amount in the third or fourth decimal place. This is why both introductory chemistry and analytical chemistry courses often accept the square root approximation in this kind of problem.
Important chemistry behind potassium hypochlorite
Potassium hypochlorite belongs to the hypochlorite family, which is closely related to sodium hypochlorite, the active component in many bleach solutions. In aqueous chemistry, the relevant conjugate acid-base pair is HOCl and ClO–. Hypochlorous acid is a weak acid and also a powerful oxidizing and disinfecting species. The acid-base distribution between HOCl and ClO– is very important in water treatment, sanitation chemistry, and environmental analysis because disinfecting efficiency often depends strongly on pH. At lower pH, a larger fraction exists as HOCl. At higher pH, more exists as ClO–.
For the pH problem in this calculator, however, we are starting with potassium hypochlorite itself. Because it provides the conjugate base directly, the solution shifts toward hydroxide formation. That basicity is the direct reason the pH lands around 10.7 instead of near neutral.
Comparison table: pH as concentration changes
The table below shows how the pH changes for potassium hypochlorite solutions of several concentrations when Ka(HOCl) is fixed at 3.0 × 10-8 and Kw is 1.0 × 10-14. These values come from the same weak base relation used in the calculator.
| KClO concentration (M) | Kb for ClO– | Estimated [OH–] (M) | Estimated pOH | Estimated pH |
|---|---|---|---|---|
| 0.010 | 3.33 × 10-7 | 5.77 × 10-5 | 4.24 | 9.76 |
| 0.050 | 3.33 × 10-7 | 1.29 × 10-4 | 3.89 | 10.11 |
| 0.100 | 3.33 × 10-7 | 1.83 × 10-4 | 3.74 | 10.26 |
| 0.750 | 3.33 × 10-7 | 5.00 × 10-4 | 3.30 | 10.70 |
| 1.000 | 3.33 × 10-7 | 5.77 × 10-4 | 3.24 | 10.76 |
Comparison table: HOCl acid strength values from common references and the effect on pH
Depending on the source, literature values for the acid dissociation constant of hypochlorous acid can vary slightly because of temperature, ionic strength, and reporting conventions. Even a small shift in Ka changes Kb and therefore changes the calculated pH slightly. The table below illustrates the sensitivity.
| Assumed Ka for HOCl | Calculated Kb for ClO– | Estimated [OH–] for 0.75 M KClO | Estimated pH |
|---|---|---|---|
| 2.9 × 10-8 | 3.45 × 10-7 | 5.09 × 10-4 | 10.71 |
| 3.0 × 10-8 | 3.33 × 10-7 | 5.00 × 10-4 | 10.70 |
| 3.5 × 10-8 | 2.86 × 10-7 | 4.64 × 10-4 | 10.67 |
Common mistakes students make
- Treating KClO as a strong base. It is not a strong base like NaOH. It is a salt that contains the weak base ClO–.
- Using Ka directly without converting to Kb. Since ClO– is the base in solution, the equilibrium constant needed is Kb.
- Forgetting that K+ is a spectator ion. Potassium does not significantly hydrolyze in water.
- Subtracting from 7 instead of 14. Once you find pOH, use pH = 14 – pOH at 25 C.
- Ignoring the temperature assumption. If temperature changes substantially, Kw changes too, and the exact pH relation shifts.
When should you use the exact quadratic solution?
For the specific case of 0.75 M KClO, the approximation works extremely well. Still, there are cases where the exact quadratic is preferred:
- When the concentration is very low
- When the weak acid or weak base is comparatively stronger
- When your instructor explicitly requires an exact equilibrium solution
- When you are doing lab-quality calculations and want to minimize rounding error
The exact equilibrium expression is:
Kb = x^2 / (C – x)
Rearranging leads to:
x^2 + Kb x – Kb C = 0
Then the physically meaningful root is:
x = (-Kb + √(Kb^2 + 4KbC)) / 2
That x value is the equilibrium hydroxide concentration. The calculator above supports both methods, so you can compare them directly.
Real world relevance of hypochlorite chemistry
Hypochlorite chemistry matters well beyond the classroom. Water treatment operators, sanitation engineers, and environmental chemists all need to understand how chlorine species behave across different pH ranges. The balance between HOCl and ClO– affects oxidation potential, microbial inactivation, and byproduct formation. Although the current problem asks for the pH of a prepared KClO solution, the same acid-base framework is central in practical disinfection chemistry.
For reliable reference material on chlorine chemistry and water quality, consult these authoritative resources:
- U.S. Environmental Protection Agency drinking water regulations and contaminants
- National Institutes of Health PubChem database
- Chemistry LibreTexts educational reference
Quick summary
- Potassium hypochlorite dissociates into K+ and ClO–.
- ClO– is a weak base because it is the conjugate base of HOCl.
- Use Kb = Kw / Ka to obtain the base dissociation constant.
- For 0.75 M KClO and Ka(HOCl) = 3.0 × 10-8, Kb = 3.33 × 10^-7.
- The calculated pH is approximately 10.70 at 25 C.