Calculate the pH of a 0.26 M CH3COONa Solution
Use this premium sodium acetate hydrolysis calculator to find pH, pOH, hydroxide concentration, and the base dissociation constant relationship for acetate in water.
Interactive Calculator
For sodium acetate, the acetate ion behaves as a weak base in water: CH3COO- + H2O ⇌ CH3COOH + OH-. Enter your values below or use the default 0.26 M setup.
Expert Guide: How to Calculate the pH of a 0.26 M CH3COONa Solution
To calculate the pH of a 0.26 M CH3COONa solution, 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 sodium ion is essentially neutral in water, the chemistry that controls pH comes from the acetate ion, CH3COO-. Acetate reacts with water to produce hydroxide ions, so the final solution is basic rather than acidic. This is why the pH of a sodium acetate solution is greater than 7 under typical room-temperature conditions.
In practical chemistry classes, this is one of the most important examples of salt hydrolysis. Students often memorize that salts of weak acids and strong bases form basic solutions, but the real skill is knowing how to convert that idea into a numerical pH result. For a 0.26 M solution of CH3COONa, the standard answer at 25°C using Ka = 1.8 × 10-5 for acetic acid and Kw = 1.0 × 10-14 is approximately pH = 9.08. The calculator above automates the process, but understanding each step is what makes the result trustworthy.
Step 1: Write the species present in solution
When sodium acetate dissolves, it dissociates almost completely:
The sodium ion does not significantly react with water, so it can be ignored in the acid-base calculation. The acetate ion does react:
This equilibrium generates hydroxide ions, which raise the pH above neutral.
Step 2: Convert Ka of acetic acid into Kb of acetate
The acid dissociation constant of acetic acid and the base dissociation constant of acetate are related by:
Using typical 25°C values:
So:
This tells you acetate is a weak base, not a strong one. That matters because the hydroxide concentration generated will be much smaller than the original 0.26 M acetate concentration.
Step 3: Set up the equilibrium expression
Let x be the amount of OH- formed at equilibrium. Then the ICE setup is:
- Initial acetate concentration = 0.26 M
- Initial hydroxide concentration from hydrolysis = 0
- Change = -x for CH3COO-, +x for CH3COOH, +x for OH-
- Equilibrium = 0.26 – x, x, x
The Kb expression becomes:
Since Kb is small and concentration is fairly large, x is much smaller than 0.26, so most textbook solutions use the approximation:
Then:
This x value is the hydroxide concentration:
Step 4: Calculate pOH and pH
Now compute pOH from hydroxide concentration:
Finally:
Therefore, the pH of a 0.26 M CH3COONa solution is approximately 9.08 at 25°C.
Why the Approximation Works So Well
The weak-base approximation assumes that x is small enough that 0.26 – x is essentially 0.26. You can check the validity by calculating the percent ionization:
That is extremely small, so the approximation is excellent. In fact, the exact quadratic result and the approximate result differ by such a tiny amount that they are usually identical to two decimal places in pH. This is why many instructors encourage the shortcut after students understand the equilibrium setup.
Common Mistakes When Solving CH3COONa pH Problems
- Treating sodium acetate as neutral. It is not neutral because acetate is the conjugate base of a weak acid.
- Using Ka directly instead of converting to Kb. Since acetate acts as a base, Kb is the equilibrium constant you need.
- Using the salt concentration as [OH-]. Only a tiny fraction hydrolyzes, so [OH-] is much smaller than 0.26 M.
- Mixing up pH and pOH. If you calculate hydroxide concentration first, you must find pOH before converting to pH.
- Forgetting temperature dependence. The common 14.00 relationship for pH + pOH assumes 25°C and a standard value of Kw.
Comparison Table: pH of Sodium Acetate at Different Concentrations
The table below uses Ka = 1.8 × 10-5 and Kw = 1.0 × 10-14 at 25°C. These values are representative of standard introductory chemistry calculations. The pH values are based on weak-base hydrolysis and show how pH increases modestly as concentration rises.
| CH3COONa Concentration (M) | Kb of Acetate | Approx. [OH-] (M) | pOH | pH |
|---|---|---|---|---|
| 0.010 | 5.56 × 10^-10 | 2.36 × 10^-6 | 5.627 | 8.373 |
| 0.050 | 5.56 × 10^-10 | 5.27 × 10^-6 | 5.278 | 8.722 |
| 0.100 | 5.56 × 10^-10 | 7.45 × 10^-6 | 5.128 | 8.872 |
| 0.260 | 5.56 × 10^-10 | 1.20 × 10^-5 | 4.920 | 9.080 |
| 0.500 | 5.56 × 10^-10 | 1.67 × 10^-5 | 4.778 | 9.222 |
Comparison Table: Acetic Acid Data Used in the Calculation
Different textbooks report slightly different acetic acid dissociation constants, often because of rounding or reference conditions. The resulting pH for the same 0.26 M sodium acetate solution changes only a little, but it is good scientific practice to report the constant you used.
| Reference Ka for CH3COOH | Calculated Kb for CH3COO- | Estimated pH of 0.26 M CH3COONa | Interpretation |
|---|---|---|---|
| 1.75 × 10^-5 | 5.71 × 10^-10 | 9.086 | Slightly more basic result |
| 1.80 × 10^-5 | 5.56 × 10^-10 | 9.080 | Common textbook value |
| 1.84 × 10^-5 | 5.43 × 10^-10 | 9.075 | Slightly lower pH |
Exact Quadratic vs Approximate Method
If you want maximum mathematical rigor, solve the full equilibrium expression without approximation:
Then solve for the physically meaningful positive root:
For this problem, the exact and approximate answers are almost identical because Kb is tiny compared with the concentration term. The calculator above lets you switch methods so you can compare them directly. That is useful for homework checking, lab reports, and exam review.
How This Connects to Buffer Chemistry
Sodium acetate also appears constantly in buffer problems. A solution containing both acetic acid and sodium acetate forms an acetate buffer. In that context, pH is often calculated with the Henderson-Hasselbalch equation rather than with simple hydrolysis. However, when the problem gives only CH3COONa dissolved in water, this is not a buffer by itself. It is a weakly basic salt solution, so hydrolysis is the correct approach.
This distinction matters in analytical chemistry, biochemistry, and laboratory practice. Acetate systems are common in pH-controlled experiments because acetic acid and acetate are inexpensive, safe to handle compared with many other systems, and effective in mildly acidic to near-neutral ranges when used together. But isolated sodium acetate in water will still push the pH upward because the acetate ion consumes protons from water indirectly by generating OH-.
Practical Interpretation of the Result
A pH of about 9.08 means the solution is mildly basic. It is nowhere near the alkalinity of strong bases such as sodium hydroxide at similar formal concentration, but it is definitely not neutral. In laboratory settings, this matters for indicator color changes, reaction rates, metal ion solubility, and biological compatibility. Even modestly basic conditions can influence ester hydrolysis, enzyme activity, and the stability of pH-sensitive reagents.
If your instructor asks whether the solution is acidic, basic, or neutral, the answer is basic. If the question asks for the dominant chemistry behind the pH, the answer is hydrolysis of the acetate ion. If the question asks why sodium ion is ignored, the reason is that Na+ is the cation of a strong base and does not significantly affect pH in aqueous solution.
Fast Summary for Exams
- CH3COONa dissociates into Na+ and CH3COO-.
- CH3COO- is the conjugate base of weak acid CH3COOH.
- Use Kb = Kw / Ka.
- Set [OH-] ≈ √(KbC) for a quick weak-base estimate.
- For C = 0.26 M and Ka = 1.8 × 10^-5, pH ≈ 9.08.
Authoritative References for Constants and Water Chemistry
- NIST Chemistry WebBook
- UC Davis chemistry content on acid-base properties of salts
- USGS pH and Water Science overview
Final Answer
Using standard 25°C constants, the pH of a 0.26 M CH3COONa solution is approximately 9.08. The result comes from acetate hydrolysis, which produces a small but measurable hydroxide concentration of about 1.20 × 10-5 M. That gives a pOH of about 4.92 and confirms that sodium acetate solutions are basic, not neutral.