Calculate the pH of a 0.25 M CH3COONa Solution
Use this premium sodium acetate pH calculator to determine the alkalinity of a 0.25 M CH3COONa solution from first principles. The tool estimates Kb from the acetic acid Ka, calculates hydroxide concentration, and returns pOH and pH with a visual concentration trend chart.
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How to calculate the pH of a 0.25 M CH3COONa solution
When you need to calculate the pH of a 0.25 M CH3COONa solution, you are dealing with a classic weak base hydrolysis problem. CH3COONa is sodium acetate, the sodium salt of acetic acid. Because sodium hydroxide is a strong base and acetic acid is a weak acid, the acetate ion left in solution behaves as a weak base. That means the final pH is above 7, not neutral.
Many students initially make the mistake of treating sodium acetate like a neutral salt because it dissociates completely in water. It is true that the ionic dissociation is complete:
However, complete dissociation is only the first step. The acetate ion then reacts with water:
This equilibrium generates hydroxide ions, so the solution becomes basic. The amount of hydroxide formed depends on the base dissociation constant of acetate, Kb. Since acetate is the conjugate base of acetic acid, Kb is found from the well-known relation:
Step by step setup for 0.25 M sodium acetate
- Start with the concentration of sodium acetate: C = 0.25 M.
- Use the acetic acid dissociation constant. A common 25 degrees C value is Ka = 1.74 × 10-5.
- Use Kw = 1.0 × 10-14 at 25 degrees C.
- Calculate the base constant of acetate: Kb = Kw / Ka.
- Solve for the hydroxide concentration [OH-].
- Calculate pOH, then pH.
Numerical calculation
Using the standard values:
For a weak base, a common approximation is:
Substitute the values:
Now compute pOH:
Finally, at 25 degrees C:
So the pH of a 0.25 M CH3COONa solution is approximately 9.08 under typical textbook conditions. If you use the exact quadratic method instead of the approximation, the answer is almost identical because the degree of hydrolysis is very small compared with the starting concentration.
Why sodium acetate gives a basic pH
The chemistry behind sodium acetate is a good illustration of conjugate acid-base behavior. Acetic acid is weak, which means it does not donate protons completely in water. Its conjugate base, acetate, therefore retains a measurable tendency to accept a proton from water. When it does, hydroxide ions are produced. By contrast, sodium ions from NaOH are spectators and do not affect pH under normal conditions.
This is why salts can be acidic, basic, or neutral depending on the acid and base from which they are formed:
- Strong acid + strong base usually gives a neutral salt.
- Weak acid + strong base usually gives a basic salt.
- Strong acid + weak base usually gives an acidic salt.
CH3COONa belongs firmly in the second category. The acetate ion hydrolyzes, and the solution becomes basic.
Exact method versus approximation
If you want more rigor, you can solve the equilibrium exactly. Let x be the hydroxide concentration produced by hydrolysis:
The equilibrium expression is:
Rearrange into a quadratic equation:
Solving gives nearly the same x value as the square root approximation. In this problem, x is far smaller than 0.25, so replacing 0.25 – x with 0.25 is excellent. This is one reason textbook pH estimates for sodium acetate are so reliable.
Comparison table: sodium acetate pH at different concentrations
The pH of sodium acetate changes with concentration because the hydrolysis equilibrium depends on the starting acetate ion concentration. The following values use Ka = 1.74 × 10-5 and Kw = 1.0 × 10-14 at 25 degrees C. The pH values are approximate but representative of standard general chemistry calculations.
| CH3COONa concentration (M) | Estimated [OH-] (M) | Estimated pOH | Estimated pH |
|---|---|---|---|
| 0.010 | 2.40 × 10-6 | 5.62 | 8.38 |
| 0.050 | 5.36 × 10-6 | 5.27 | 8.73 |
| 0.100 | 7.58 × 10-6 | 5.12 | 8.88 |
| 0.250 | 1.20 × 10-5 | 4.92 | 9.08 |
| 0.500 | 1.70 × 10-5 | 4.77 | 9.23 |
| 1.000 | 2.40 × 10-5 | 4.62 | 9.38 |
How sodium acetate compares with related solutions
Students also often ask how sodium acetate compares with acetic acid or with a sodium chloride solution. This comparison helps clarify why sodium acetate has a pH above 7 while the acid itself has a pH below 7.
| Solution | Main acid-base species | Typical 0.25 M pH range | Why |
|---|---|---|---|
| Acetic acid, CH3COOH | Weak acid | About 2.7 to 2.9 | Releases H+ only partially, but enough to make solution acidic |
| Sodium acetate, CH3COONa | Conjugate base of a weak acid | About 9.0 to 9.1 | Acetate hydrolyzes and produces OH- |
| Sodium chloride, NaCl | Spectator ions only | Near 7.0 | Neither ion appreciably hydrolyzes in water |
Common mistakes when calculating the pH of 0.25 M CH3COONa
- Using Ka directly instead of Kb. Because acetate is a base, you must use Kb. If only Ka is given, convert it using Kb = Kw / Ka.
- Treating the salt as neutral. Sodium acetate is not neutral in water because acetate hydrolyzes.
- Forgetting the pOH step. The hydrolysis gives OH-, so you compute pOH first and then convert to pH.
- Using the wrong concentration in the ICE table. The starting concentration for acetate is the sodium acetate molarity, 0.25 M.
- Ignoring temperature effects. If temperature differs substantially from 25 degrees C, both Ka and Kw may change, and the pH will shift slightly.
When the approximation is valid
The weak base approximation works well when the fraction ionized is small. In this problem, [OH-] is on the order of 10-5 M while the starting concentration is 0.25 M. That means the hydrolyzed fraction is tiny, far less than 5 percent. Therefore, using:
is fully justified. In a laboratory or exam setting, this shortcut can save time without sacrificing accuracy.
Broader context: buffers and acetate chemistry
Sodium acetate is especially important because it is commonly paired with acetic acid to make acetate buffers. A buffer resists pH changes when small amounts of acid or base are added. In that case, the Henderson-Hasselbalch equation becomes useful. But for a pure sodium acetate solution with no added acetic acid, the hydrolysis approach used in this calculator is the correct one.
The acetate system is widely used in chemistry, biology, and analytical labs because it is inexpensive, reasonably safe to handle, and has well-characterized equilibrium constants. Knowing how to calculate the pH of sodium acetate solutions helps with titration analysis, buffer preparation, and general equilibrium reasoning.
Authoritative references for equilibrium constants and water chemistry
For reliable reference material, consult: NIST, LibreTexts Chemistry, U.S. Environmental Protection Agency, U.S. Geological Survey, and UC Berkeley Chemistry.
Specific authoritative educational and government sources relevant to acid-base chemistry include EPA pH overview, USGS pH and water science, and UC Berkeley Department of Chemistry.
Final answer
Under standard 25 degrees C assumptions with Ka for acetic acid = 1.74 × 10-5 and Kw = 1.0 × 10-14, the pH of a 0.25 M CH3COONa solution is approximately 9.08. This value comes from acetate ion hydrolysis, which generates a small but significant concentration of hydroxide ions in water.
If you want to explore how the pH changes with concentration or with a slightly different Ka value, use the calculator above. It updates the result instantly and draws a chart so you can compare 0.25 M sodium acetate with nearby concentrations.