Calculate the pH of a 0.200 M NaClO2 Solution
Use this premium acid-base calculator to find the pH, pOH, hydroxide concentration, hydronium concentration, and chlorite hydrolysis behavior for sodium chlorite solutions at 25 degrees Celsius.
Calculated Results
Enter your values and click Calculate pH to see the full equilibrium breakdown for NaClO2.
Chart compares pH, pOH, and the logarithmic concentrations of hydronium and hydroxide in the computed equilibrium state.
Expert Guide: How to Calculate the pH of a 0.200 M NaClO2 Solution
Sodium chlorite, written as NaClO2, is the ionic salt of sodium ion and the chlorite ion, ClO2–. When NaClO2 dissolves in water, it dissociates essentially completely into Na+ and ClO2–. The sodium ion is a spectator ion with negligible acid-base behavior under normal conditions, but the chlorite ion is different. Because ClO2– is the conjugate base of chlorous acid, HClO2, it can react with water and produce a mildly basic solution. That is why the pH of a sodium chlorite solution is expected to be greater than 7.
To calculate the pH of a 0.200 M NaClO2 solution, the key idea is to treat ClO2– as a weak base. The relevant equilibrium is:
ClO2– + H2O ⇌ HClO2 + OH–
The base dissociation constant for chlorite is not usually listed first in textbooks, but it can be derived from the acid dissociation constant of chlorous acid using the familiar conjugate relationship:
Kb = Kw / Ka
At 25 degrees Celsius, the ion-product constant of water is approximately 1.0 × 10-14. A commonly used value for chlorous acid is Ka = 1.1 × 10-2. Substituting gives:
Kb = (1.0 × 10-14) / (1.1 × 10-2) = 9.09 × 10-13
This very small Kb shows that chlorite is a weak base. Even though the starting concentration is 0.200 M, only a tiny fraction of chlorite ions react with water to create hydroxide. That is why the pH is basic, but only modestly so.
Step-by-step setup
- Write the hydrolysis equilibrium for ClO2–.
- Determine Kb from the given or assumed Ka of HClO2.
- Construct an ICE table using initial chlorite concentration 0.200 M.
- Solve for x, where x is the equilibrium concentration of OH– produced.
- Find pOH from pOH = -log[OH–].
- Convert to pH using pH = 14.00 – pOH at 25 degrees Celsius.
ICE table for 0.200 M NaClO2
For the reaction ClO2– + H2O ⇌ HClO2 + OH–, the initial equilibrium table is:
- Initial: [ClO2–] = 0.200, [HClO2] = 0, [OH–] = 0
- Change: -x, +x, +x
- Equilibrium: 0.200 – x, x, x
The equilibrium expression is:
Kb = x2 / (0.200 – x)
Substitute Kb = 9.09 × 10-13:
9.09 × 10-13 = x2 / (0.200 – x)
Because Kb is extremely small relative to the starting concentration, the standard approximation 0.200 – x ≈ 0.200 is very good. That gives:
x2 = (9.09 × 10-13)(0.200) = 1.818 × 10-13
x = 4.26 × 10-7 M
Since x represents hydroxide concentration, we have:
[OH–] = 4.26 × 10-7 M
Now calculate pOH:
pOH = -log(4.26 × 10-7) = 6.37
Finally:
pH = 14.00 – 6.37 = 7.63
Final answer: the pH of a 0.200 M NaClO2 solution is approximately 7.63 when Ka for HClO2 is taken as 1.1 × 10-2 and the solution is assumed to be at 25 degrees Celsius.
Why the solution is only mildly basic
Many students expect the pH of any salt containing a conjugate base to be strongly basic. That is not always true. The strength of the basic solution depends on the strength of the parent acid. Chlorous acid is much stronger than acetic acid and many other weak acids, so its conjugate base is correspondingly much weaker. Since HClO2 has a relatively large Ka, chlorite has an extremely small Kb. The result is a pH only a little above neutral even at 0.200 M concentration.
This is a useful conceptual checkpoint. Salts of weak acids can produce basic solutions, but the magnitude of the basicity depends on how weak the original acid was. A conjugate base from a very weak acid can push pH high. A conjugate base from a relatively strong weak acid often raises pH only slightly.
Approximation versus quadratic method
For most classroom work, the square-root approximation is enough because x is tiny compared with 0.200 M. However, because the computed hydroxide concentration is only a few times larger than 1.0 × 10-7 M, some instructors prefer a more careful treatment, especially when comparing concentrations near neutrality. The calculator above includes both an approximation method and a quadratic method so you can see that they produce nearly identical pH values in this case.
The quadratic form comes directly from:
Kb(0.200 – x) = x2
or
x2 + Kb x – 0.200 Kb = 0
Solving the positive root gives essentially the same hydroxide concentration because Kb is so small. The difference is typically beyond the second decimal place in pH for this specific example.
Comparison of chlorine oxyacid conjugate bases
The table below helps place chlorite in context by comparing common chlorine oxyacids and the relative basicity of their conjugate bases. As acid strength increases, conjugate-base strength decreases.
| Acid | Formula | Typical pKa | Conjugate Base | Relative Base Strength |
|---|---|---|---|---|
| Hypochlorous acid | HClO | 7.5 | ClO– | Much stronger base than chlorite |
| Chlorous acid | HClO2 | 1.96 | ClO2– | Weak base |
| Chloric acid | HClO3 | About -1 | ClO3– | Extremely weak base |
| Perchloric acid | HClO4 | About -10 | ClO4– | Negligible basicity in water |
This trend explains why chlorite does not behave like hypochlorite. The parent acid HClO2 is vastly stronger than HClO, so ClO2– is vastly weaker as a base than ClO–.
Effect of concentration on pH
Another useful way to understand the chemistry is to see how the calculated pH changes as NaClO2 concentration changes, while keeping Ka = 1.1 × 10-2 and Kw = 1.0 × 10-14. Because this is weak-base hydrolysis, the pH rises gradually rather than sharply.
| NaClO2 Concentration (M) | Kb of ClO2– | Approximate [OH–] (M) | Approximate pOH | Approximate pH |
|---|---|---|---|---|
| 0.010 | 9.09 × 10-13 | 9.53 × 10-8 | 7.02 | 6.98 to 7.02 range with water effects |
| 0.050 | 9.09 × 10-13 | 2.13 × 10-7 | 6.67 | 7.33 |
| 0.100 | 9.09 × 10-13 | 3.02 × 10-7 | 6.52 | 7.48 |
| 0.200 | 9.09 × 10-13 | 4.26 × 10-7 | 6.37 | 7.63 |
| 1.000 | 9.09 × 10-13 | 9.53 × 10-7 | 6.02 | 7.98 |
The statistics in the table show a characteristic weak-base pattern: a tenfold increase in concentration does not create a one-unit increase in pH. Instead, pH changes more slowly because hydroxide concentration is proportional to the square root of both Kb and the formal concentration when the small-x approximation is valid.
Common mistakes students make
- Treating NaClO2 as neutral. Not all sodium salts are neutral. You must examine the anion.
- Using Ka directly instead of Kb. The reacting species in solution is ClO2–, so the correct equilibrium constant is Kb.
- Assuming the solution is strongly basic. Chlorite is only a weak base because HClO2 is a relatively strong weak acid.
- Forgetting the pOH step. Once [OH–] is found, convert to pOH, then to pH.
- Ignoring temperature. If temperature changes, Kw changes too, and pH shifts slightly.
When would a more advanced treatment matter?
For most introductory chemistry problems, the workflow shown above is sufficient. But in research, industrial water chemistry, or analytical chemistry, you may need to account for ionic strength, activity coefficients, temperature corrections, and decomposition side reactions. Sodium chlorite also has practical importance in oxidation chemistry and disinfection applications, so real systems can become more complicated than the ideal textbook equilibrium.
If you are using this calculation in a lab or process setting, remember that measured pH can differ from ideal values due to meter calibration, dissolved carbon dioxide, competing redox reactions, and sample contamination. In other words, the equilibrium math gives the theoretical pH of a pure aqueous solution under stated conditions.
Authoritative references for acid-base and pH fundamentals
For deeper study, review these authoritative sources:
- National Institute of Standards and Technology (NIST) for measurement standards and pH-related metrology resources.
- U.S. Environmental Protection Agency pH overview for applied pH context in aqueous systems.
- University of California, Berkeley Chemistry for academic chemistry instruction and equilibrium concepts.
Bottom line
If you are asked to calculate the pH of a 0.200 M NaClO2 solution, identify chlorite as the conjugate base of chlorous acid, derive Kb from Ka, solve the weak-base equilibrium, and convert hydroxide concentration into pH. Using Ka = 1.1 × 10-2 for HClO2 at 25 degrees Celsius leads to a pH of about 7.63. That answer is chemically sensible because chlorite is a weak base, so the solution is only mildly alkaline.