Calculate the pH of 0.01 M Solution of Sodium Acetate
Use this interactive hydrolysis calculator to determine the pH of a sodium acetate solution from concentration, acid dissociation data, and temperature. By default, it is preloaded for a 0.01 M sodium acetate solution with acetic acid pKa = 4.76 at 25 C, which gives a mildly basic pH.
Example: 0.01 for a 0.01 M CH3COONa solution.
Temperature changes Kw, so pKw is adjusted automatically.
Sodium acetate is the conjugate base of acetic acid.
Default is pKa = 4.76. If you switch to Ka mode, enter about 1.8e-5.
The exact method is preferred. The approximation is shown for comparison and quick classroom checks.
Expert Guide: How to Calculate the pH of 0.01 M Solution of Sodium Acetate
To calculate the pH of 0.01 M solution of sodium acetate, you need to recognize that sodium acetate is a salt formed from a strong base, sodium hydroxide, and a weak acid, acetic acid. Because the anion acetate, CH3COO-, is the conjugate base of a weak acid, it reacts with water and produces hydroxide ions. That means the solution is basic, not neutral. This is a classic salt hydrolysis problem in general chemistry, analytical chemistry, and introductory biochemistry.
The hydrolysis reaction is: CH3COO- + H2O ⇌ CH3COOH + OH-
Since sodium ions are spectator ions, the chemistry that matters is entirely controlled by acetate. The basicity of acetate is related to the acidity of acetic acid through the relationship Kb = Kw / Ka. At 25 C, Kw is 1.0 x 10-14, and a commonly used value for acetic acid is Ka = 1.8 x 10-5, which corresponds to pKa ≈ 4.76.
Quick answer for 0.01 M sodium acetate at 25 C
For a 0.01 M solution of sodium acetate, the accepted classroom calculation gives a pH of about 8.37. This comes from finding Kb first:
- Ka of acetic acid = 1.8 x 10-5
- Kw = 1.0 x 10-14
- Kb = Kw / Ka = 1.0 x 10-14 / 1.8 x 10-5 = 5.56 x 10-10
Then use the weak base relation for acetate hydrolysis:
Kb = x2 / C
where C = 0.01 M and x = [OH-]. Solving gives:
- x = √(Kb x C) = √(5.56 x 10-10 x 0.01)
- x = √(5.56 x 10-12) ≈ 2.36 x 10-6 M
- pOH = -log(2.36 x 10-6) ≈ 5.63
- pH = 14.00 – 5.63 = 8.37
Why sodium acetate solutions are basic
Many students first assume that all salts make neutral solutions, but that is only true for salts formed from a strong acid and a strong base. Sodium acetate is different. The sodium ion, Na+, does not significantly react with water, but acetate does. Because acetic acid is weak, its conjugate base is strong enough to accept a proton from water to a measurable extent. That hydrolysis creates OH-, which pushes the pH above 7.
A useful memory rule is this:
- Strong acid + strong base salt → approximately neutral
- Strong acid + weak base salt → acidic
- Weak acid + strong base salt → basic
Sodium acetate belongs to the third category, so a basic pH is expected even before you calculate the exact number.
Step-by-step method
- Write the hydrolysis equation for acetate in water.
- Find Kb using Kb = Kw / Ka.
- Set the initial acetate concentration equal to the salt concentration.
- Let x be the amount hydrolyzed, which also equals [OH-] at equilibrium.
- Solve for x using either the approximation or the exact quadratic expression.
- Convert [OH-] to pOH, then convert pOH to pH using pH = pKw – pOH.
Exact equilibrium setup
If the initial concentration of acetate is 0.01 M, then the ICE setup is:
- Initial: [CH3COO-] = 0.01, [CH3COOH] = 0, [OH-] = 0
- Change: -x, +x, +x
- Equilibrium: [CH3COO-] = 0.01 – x, [CH3COOH] = x, [OH-] = x
Substitute into the base dissociation expression:
Kb = x2 / (0.01 – x)
Because Kb is very small, x is much smaller than 0.01, so many textbooks use 0.01 – x ≈ 0.01. This produces the same pH to two decimal places in this problem. The calculator above can show either the exact or approximate result.
Comparison table: sodium acetate concentration vs predicted pH at 25 C
The table below uses Ka = 1.8 x 10-5 for acetic acid and standard 25 C conditions. It shows how pH rises as sodium acetate concentration increases. This happens because more acetate is available to hydrolyze and generate hydroxide.
| NaCH3COO concentration (M) | Kb of acetate | Approx. [OH-] (M) | Predicted pH |
|---|---|---|---|
| 0.001 | 5.56 x 10-10 | 7.46 x 10-7 | 7.87 |
| 0.010 | 5.56 x 10-10 | 2.36 x 10-6 | 8.37 |
| 0.100 | 5.56 x 10-10 | 7.46 x 10-6 | 8.87 |
| 1.000 | 5.56 x 10-10 | 2.36 x 10-5 | 9.37 |
Temperature matters: use pKw, not always 14.00
A subtle but important point is that pH and pOH do not always sum to exactly 14.00 unless the temperature is 25 C. The ion-product constant of water, Kw, changes with temperature. Advanced students and laboratory analysts should use pKw = -log(Kw) at the selected temperature. That is why the calculator includes a temperature selector.
| Temperature | Kw of water | pKw | Practical implication |
|---|---|---|---|
| 20 C | 6.81 x 10-15 | 14.17 | Neutral pH is slightly above 7.00 |
| 25 C | 1.00 x 10-14 | 14.00 | Standard textbook reference point |
| 30 C | 1.47 x 10-14 | 13.83 | Neutral pH is slightly below 7.00 |
Common mistakes when solving this problem
- Using the Ka equation directly on sodium acetate instead of converting to Kb first.
- Forgetting that sodium acetate is basic because acetate is the conjugate base of a weak acid.
- Assuming pH = 7 because it is a salt.
- Using pH = 14 – pOH at temperatures other than 25 C without checking pKw.
- Mixing up acetic acid and acetate concentrations in the ICE table.
When the approximation is valid
In this example, the approximation is excellent. The hydroxide concentration is only around 2.36 x 10-6 M, while the initial acetate concentration is 1.00 x 10-2 M. That means the fraction hydrolyzed is tiny, roughly 0.024 percent. Because x is so much smaller than the starting concentration, replacing 0.01 – x with 0.01 introduces negligible error. In other words, the exact and approximate solutions are nearly identical.
Relation to buffers and the common ion effect
Sodium acetate becomes even more important when paired with acetic acid to form an acetate buffer. In a pure sodium acetate solution, hydrolysis controls the pH. But if acetic acid is also present, the Henderson-Hasselbalch equation often becomes the preferred tool:
pH = pKa + log([A-] / [HA])
In that setting, sodium acetate supplies the acetate component, while acetic acid supplies the proton donor. The common ion effect suppresses dissociation of acetic acid and stabilizes pH. This is why acetate systems are standard in laboratory buffers, biological sample preparation, and analytical methods.
Authoritative references for acid-base data and pH fundamentals
If you want to verify equilibrium constants and pH principles from trusted sources, these references are useful:
- NIST Chemistry WebBook (.gov) for acetic acid compound data and physical chemistry references.
- U.S. EPA pH overview (.gov) for practical interpretation of pH and aqueous chemistry.
- University of Wisconsin acid-base chemistry resource (.edu) for equilibrium concepts used in weak acid and conjugate base calculations.
Final takeaway
To calculate the pH of 0.01 M solution of sodium acetate, treat acetate as a weak base. Convert acetic acid strength into acetate basicity with Kb = Kw / Ka, solve for hydroxide concentration, and then convert to pH. Under standard 25 C conditions with acetic acid pKa = 4.76, the answer is approximately pH 8.37. That value makes chemical sense because sodium acetate is the salt of a weak acid and a strong base, so its aqueous solution is mildly basic.
Use the calculator above whenever you want a fast, transparent solution with full steps, equilibrium concentrations, and a chart. It is especially useful for homework checks, lab prep, exam revision, and comparing how concentration or temperature changes the final pH.