Calculate the pH of 0.150 m Sodium Chlorite
Use this premium calculator to estimate the pH of sodium chlorite solutions by treating chlorite ion, ClO2–, as the conjugate base of chlorous acid, HClO2. The default setup is tuned for the common textbook case of a 0.150 concentration with chlorous acid pKa near 1.96 at 25 degrees Celsius.
Sodium Chlorite pH Calculator
Click the button to compute pH, pOH, Kb, hydroxide concentration, and the extent of hydrolysis for 0.150 sodium chlorite.
Visual Analysis
This chart compares the initial chlorite concentration with the calculated equilibrium concentrations of ClO2–, HOClO, and OH–. It helps you see why the pH stays only moderately basic even when sodium chlorite is fully dissociated as a salt.
ClO2– + H2O ⇌ HClO2 + OH–
How to Calculate the pH of 0.150 m Sodium Chlorite
Calculating the pH of a sodium chlorite solution is a classic acid-base equilibrium problem. Although sodium chlorite, NaClO2, is an ionic compound that dissociates completely in water, the chlorite ion itself is not neutral. It is the conjugate base of chlorous acid, HClO2, so once the salt dissolves, the chlorite ion reacts with water to produce a small amount of hydroxide. That hydrolysis step is what raises the pH above 7.
If you were asked to “calculate the pH of 0.150 m sodium chlorite,” the standard approach is to treat the given concentration as the concentration of ClO2– in solution, determine the base dissociation constant Kb from the acid dissociation constant Ka of chlorous acid, and then solve for the equilibrium hydroxide concentration. For the common textbook value pKa = 1.96, the pH comes out to about 7.57 at 25 degrees Celsius, assuming an ideal dilute aqueous solution.
Step 1: Identify the species that controls pH
Sodium chlorite dissociates essentially completely:
NaClO2 → Na+ + ClO2–
The sodium ion is a spectator ion from a strong base and does not significantly affect pH. The chlorite ion is the important species because it is the conjugate base of chlorous acid. In water it undergoes hydrolysis:
ClO2– + H2O ⇌ HClO2 + OH–
That means the solution should be basic, but only weakly basic, because chlorous acid itself is a relatively strong weak acid. A stronger parent acid implies a weaker conjugate base.
Step 2: Relate Ka and Kb
For a conjugate acid-base pair at 25 degrees Celsius:
Ka × Kb = Kw = 1.0 × 10-14
If chlorous acid has pKa = 1.96, then:
Ka = 10-1.96 = 1.10 × 10-2
Now solve for Kb:
Kb = Kw / Ka = (1.0 × 10-14) / (1.10 × 10-2) = 9.12 × 10-13
This is a very small base dissociation constant, which tells you that chlorite ion hydrolyzes only slightly.
Step 3: Set up the equilibrium table
Use an ICE table for the hydrolysis reaction:
- Initial: [ClO2–] = 0.150, [HClO2] = 0, [OH–] = 0
- Change: -x, +x, +x
- Equilibrium: 0.150 – x, x, x
Substitute into the equilibrium expression:
Kb = [HClO2][OH–] / [ClO2–] = x2 / (0.150 – x)
Because Kb is so small, x will be tiny compared with 0.150, so many instructors allow the approximation:
x ≈ √(KbC) = √[(9.12 × 10-13)(0.150)] = 3.70 × 10-7 M
That gives the hydroxide concentration.
Step 4: Convert hydroxide concentration into pOH and pH
Now calculate pOH:
pOH = -log(3.70 × 10-7) = 6.43
Then compute pH at 25 degrees Celsius:
pH = 14.00 – 6.43 = 7.57
So the expected answer for a 0.150 sodium chlorite solution is approximately pH = 7.57.
Why the pH is only slightly basic
Students sometimes expect every salt of a weak acid to produce a strongly basic solution. Sodium chlorite is a good reminder that the magnitude matters. Chlorous acid is much stronger than acids like acetic acid or hypochlorous acid, so its conjugate base is correspondingly much weaker. Even though the formal concentration is 0.150, the extent of hydrolysis is tiny. In fact, the equilibrium hydroxide concentration is only on the order of 10-7 M.
The percent hydrolysis is therefore very small:
% hydrolysis = (x / C) × 100 = (3.70 × 10-7 / 0.150) × 100 ≈ 0.000247%
This tiny value explains why the chlorite concentration barely changes from its starting concentration and why the square-root approximation is excellent here.
Molality versus molarity in the problem statement
The notation “0.150 m” formally means molality, not molarity. Molality is moles of solute per kilogram of solvent, while molarity is moles of solute per liter of solution. In many general chemistry exercises involving relatively dilute aqueous solutions, instructors often use these values interchangeably as an approximation because the numerical difference is small. At 0.150 concentration in water, the pH estimate using 0.150 M versus 0.150 m is typically very similar unless a highly precise density correction is required.
For careful work, especially in analytical chemistry or thermodynamics, you would account for:
- Solution density
- Activity coefficients
- Temperature dependence of Kw
- The exact thermodynamic dissociation constant used
For standard educational calculations, however, treating 0.150 m sodium chlorite as approximately 0.150 M is common and usually expected.
Exact quadratic solution versus approximation
The approximation x = √(KbC) works when x is much smaller than the initial concentration. That condition is overwhelmingly satisfied here. Still, an exact quadratic solution can be written as:
x2 + Kb x – KbC = 0
Solving gives:
x = [-Kb + √(Kb2 + 4KbC)] / 2
When you substitute Kb = 9.12 × 10-13 and C = 0.150, the result is effectively the same as the approximation to the digits usually reported in chemistry coursework. This is why the approximation is not just simpler, but also fully justified.
Comparison table: how chlorite differs from other conjugate bases
| Conjugate base in water | Parent acid | Representative pKa of parent acid | Estimated Kb at 25 degrees C | Expected basicity trend |
|---|---|---|---|---|
| ClO2– | HClO2 | 1.96 | 9.12 × 10-13 | Very weak base, slightly basic solution |
| CH3COO– | Acetic acid | 4.76 | 5.75 × 10-10 | Noticeably stronger weak base than chlorite |
| ClO– | HOCl | 7.53 | 3.39 × 10-7 | Much more basic than chlorite |
| CN– | HCN | 9.21 | 1.62 × 10-5 | Clearly basic solution |
This table shows the key principle: the stronger the parent acid, the weaker the conjugate base. Since chlorous acid is comparatively strong for a weak acid, chlorite ion is only weakly basic.
Numerical results for several sodium chlorite concentrations
Using pKa = 1.96 and 25 degrees Celsius, the approximate pH values below show how concentration affects the final answer. Because hydroxide scales with the square root of concentration for a weak base approximation, the pH increases slowly rather than dramatically.
| [NaClO2] concentration | Calculated [OH–] | pOH | pH | Percent hydrolysis |
|---|---|---|---|---|
| 0.0100 | 9.55 × 10-8 M | 7.02 | 6.98 to 7.00 region with water autoionization effects | 0.000955% |
| 0.0500 | 2.14 × 10-7 M | 6.67 | 7.33 | 0.000428% |
| 0.150 | 3.70 × 10-7 M | 6.43 | 7.57 | 0.000247% |
| 0.500 | 6.75 × 10-7 M | 6.17 | 7.83 | 0.000135% |
One interesting lesson from this data is that the weaker the base, the more cautiously you should interpret low-concentration pH values. At very small concentrations, the autoionization of water becomes comparable to the hydroxide produced by hydrolysis, so more rigorous treatment may be needed.
Common mistakes when solving sodium chlorite pH problems
- Using Ka directly to compute pH. Since the solution contains the conjugate base ClO2–, you need Kb, not Ka, unless you derive Kb from Ka first.
- Treating NaClO2 as a strong base. Sodium chlorite is a salt, not a hydroxide such as NaOH. It does not release OH– stoichiometrically.
- Ignoring the conjugate relationship. The reaction is hydrolysis of ClO2–, not direct ionization of NaClO2 as an acid or base by itself.
- Confusing molality and molarity. In precision work, these are not identical. In standard textbook calculations, they are often approximated as the same for dilute aqueous solutions.
- Over-rounding Kb. Since the pH is only slightly above neutral, a rough Kb can shift the final pH by a few hundredths.
Practical context and data quality
Sodium chlorite is important industrially, especially in applications related to chlorine dioxide generation and bleaching chemistry. However, classroom pH calculations assume an idealized aqueous equilibrium system. Real industrial formulations may include buffers, stabilizers, ionic strength effects, or oxidizing side chemistry that are not part of a simple general chemistry model. If you are doing process work rather than homework, you should use validated safety data, measured pH, and formulation-specific specifications.
For fundamental chemistry references and educational support, these authoritative sources are useful:
Final takeaway
To calculate the pH of 0.150 sodium chlorite, treat chlorite ion as a weak base. Convert the known pKa of chlorous acid into Ka, derive Kb using Kw, set up the hydrolysis equilibrium, solve for hydroxide concentration, and then calculate pOH and pH. Using pKa = 1.96 at 25 degrees Celsius, the result is approximately pH 7.57. That answer makes chemical sense: the solution is basic, but only slightly, because chlorite is the conjugate base of a relatively strong weak acid.
This calculator automates that workflow while still showing the chemistry behind it. If you adjust concentration, temperature, or pKa, you can immediately see how sensitive the final pH is to each assumption.