Calculate the pH of a 0.15 M Solution of CH3COONa
This premium calculator finds the pH of aqueous sodium acetate by modeling acetate ion hydrolysis. Adjust concentration, acetic acid Ka, and temperature assumptions, then view the full result set with an interactive chart.
Sodium Acetate pH Calculator
For CH3COONa, the acetate ion acts as a weak base in water: CH3COO- + H2O ⇌ CH3COOH + OH-. The tool below uses Kb = Kw / Ka and solves the hydrolysis equilibrium.
Expert guide: how to calculate the pH of a 0.15 M solution of CH3COONa
If you need to calculate the pH of a 0.15 M solution of CH3COONa, the key idea is that sodium acetate is the salt of a weak acid and a strong base. In water, sodium acetate dissociates essentially completely into sodium ions and acetate ions. The sodium ion is a spectator ion, but the acetate ion reacts with water to produce hydroxide. That makes the final solution basic, with a pH greater than 7 at ordinary laboratory temperatures.
Written chemically, sodium acetate dissolves as CH3COONa → Na+ + CH3COO-. The acetate ion then undergoes hydrolysis according to CH3COO- + H2O ⇌ CH3COOH + OH-. Because hydroxide ions are produced, the pH rises above neutrality. This is why solutions of sodium acetate are commonly used in introductory acid-base chemistry as examples of basic salt solutions.
The exact problem here asks for the pH of a 0.15 M sodium acetate solution. At 25 C, acetic acid has a commonly used acid dissociation constant of about Ka = 1.8 × 10-5. Since acetate is the conjugate base of acetic acid, we first convert Ka into Kb using the water relation Ka × Kb = Kw. With Kw = 1.0 × 10^-14 at 25 C, we obtain the base dissociation constant of acetate.
Step by step calculation
- Start with the acid constant of acetic acid: Ka = 1.8 × 10^-5.
- Use Kb = Kw / Ka.
- At 25 C, Kb = (1.0 × 10^-14) / (1.8 × 10^-5) = 5.56 × 10^-10.
- For the hydrolysis equilibrium, let x = [OH-] formed.
- Then Kb = x^2 / (0.15 – x).
- Since Kb is very small relative to concentration, the standard approximation gives x ≈ √(Kb × C).
- So [OH-] ≈ √(5.56 × 10^-10 × 0.15) = 9.13 × 10^-6 M.
- Now compute pOH = -log(9.13 × 10^-6) ≈ 5.04.
- Finally, pH = 14.00 – 5.04 = 8.96.
Therefore, the pH of a 0.15 M solution of CH3COONa is approximately 8.96 at 25 C, using standard textbook constants. If you solve the equilibrium expression exactly with the quadratic formula, you get essentially the same answer because the hydrolysis amount is tiny compared with the initial acetate concentration.
Why sodium acetate makes the solution basic
Students often wonder why a salt can change pH at all. The answer depends on the acid and base from which the salt was formed. Sodium acetate comes from acetic acid, a weak acid, and sodium hydroxide, a strong base. The conjugate base of a weak acid retains measurable basicity in water. In contrast, the conjugate base of a strong acid is negligibly basic. Since acetate is the conjugate base of a weak acid, it removes a proton from water to a small extent, generating hydroxide ions.
- Na+ does not appreciably affect pH in water.
- CH3COO- is a weak base and hydrolyzes.
- OH- production pushes the solution into the basic range.
- The higher the salt concentration, the more basic the solution becomes, though not in a linear way.
This type of calculation is part of a larger acid-base framework in which salts can be acidic, basic, or nearly neutral depending on the strengths of their parent acids and bases. Sodium chloride is approximately neutral, ammonium chloride is acidic, and sodium acetate is basic.
Exact equilibrium setup using an ICE table
An ICE table is a structured way to write the equilibrium quantities. For acetate hydrolysis:
Reaction: CH3COO- + H2O ⇌ CH3COOH + OH-
Initial: [CH3COO-] = 0.15, [CH3COOH] = 0, [OH-] = 0
Change: -x, +x, +x
Equilibrium: 0.15 – x, x, x
Expression: Kb = x^2 / (0.15 – x)
For exact work, solve the quadratic form x^2 + Kb x – KbC = 0, where C = 0.15. The physically meaningful solution is x = (-Kb + √(Kb^2 + 4KbC)) / 2. Because Kb is on the order of 10-10 and C is 10-1, the quadratic collapses numerically to the same result as the square root approximation.
Comparison table: sodium acetate pH as concentration changes
The table below shows how the predicted pH changes with concentration at 25 C using Ka = 1.8 × 10^-5 and the weak base approximation. These values are realistic classroom and laboratory estimates.
| Sodium acetate concentration (M) | Kb of acetate | Calculated [OH-] (M) | pOH | Predicted pH at 25 C |
|---|---|---|---|---|
| 0.010 | 5.56 × 10^-10 | 2.36 × 10^-6 | 5.63 | 8.37 |
| 0.050 | 5.56 × 10^-10 | 5.27 × 10^-6 | 5.28 | 8.72 |
| 0.100 | 5.56 × 10^-10 | 7.45 × 10^-6 | 5.13 | 8.87 |
| 0.150 | 5.56 × 10^-10 | 9.13 × 10^-6 | 5.04 | 8.96 |
| 0.500 | 5.56 × 10^-10 | 1.67 × 10^-5 | 4.78 | 9.22 |
| 1.000 | 5.56 × 10^-10 | 2.36 × 10^-5 | 4.63 | 9.37 |
Notice that increasing concentration does raise the pH, but not as dramatically as many learners first expect. Since the hydroxide concentration scales roughly with the square root of concentration for a weak base salt, a tenfold increase in salt concentration does not create a tenfold increase in pH.
Comparison table: effect of water autoionization constant with temperature
Temperature affects Kw and therefore also changes pOH and pH. The next table keeps the sodium acetate concentration fixed at 0.15 M and uses the same Ka for acetic acid, while varying the water constant with temperature. The values shown are standard rounded instructional values.
| Temperature | Kw | pKw | Kb = Kw / Ka | Approximate pH for 0.15 M CH3COONa |
|---|---|---|---|---|
| 20 C | 6.81 × 10^-15 | 14.167 | 3.78 × 10^-10 | 9.04 |
| 25 C | 1.00 × 10^-14 | 14.000 | 5.56 × 10^-10 | 8.96 |
| 30 C | 1.48 × 10^-14 | 13.830 | 8.22 × 10^-10 | 8.88 |
The pH trend can look surprising at first. As temperature rises, Kw increases and neutral water itself no longer has a pH of exactly 7. Because pH and neutrality both shift with temperature, the numeric pH of the sodium acetate solution may move in ways that seem counterintuitive if you only think in terms of room-temperature rules.
Common mistakes to avoid
- Using Ka directly in the ICE table for acetate hydrolysis. Once you are working with CH3COO-, you need Kb, not Ka.
- Treating CH3COONa as a strong base. Sodium acetate is a salt, not NaOH. Its basicity is weak and comes from hydrolysis.
- Forgetting the pOH step. Hydrolysis gives [OH-], so you find pOH first, then convert to pH.
- Assuming pH 7 is always neutral. That is strictly true only at 25 C.
- Ignoring units and significant figures. Concentration should be in molarity, and constants should be entered in scientific notation carefully.
When the square root approximation is valid
The approximation x ≈ √(KbC) is valid when the equilibrium change x is much smaller than the initial concentration C. A common chemistry check is the 5 percent rule. In this problem, the hydrolyzed amount is around 9.13 × 10^-6 M, while the initial concentration is 0.15 M. The ratio is only about 0.0061%, which is far below 5 percent. That means the approximation is excellent.
This is why many textbook solutions go directly from Kb = x^2 / (C – x) to Kb ≈ x^2 / C. However, in software or advanced coursework, solving the exact quadratic is easy and avoids approximation concerns. The calculator above supports both methods so you can compare them.
Practical interpretation of the result
A pH near 8.96 means the sodium acetate solution is mildly basic. It is nowhere near as alkaline as a strong base solution of the same concentration, but it is definitely above neutral. This matters in laboratory work, especially in buffer preparation, biological media, food chemistry, and titration systems. Sodium acetate often appears in acetate buffer systems, where the pH depends on the ratio of acetic acid to acetate. In the special case of sodium acetate alone, the pH is controlled by hydrolysis rather than by the Henderson-Hasselbalch ratio formula.
Acetate systems also matter in environmental and biochemical contexts. Acetate is a common conjugate base in aqueous chemistry, and understanding its behavior helps students connect equilibrium constants, conjugate acid-base pairs, and the pH scale in a coherent way.
Authoritative references for acid-base constants and aqueous chemistry
For deeper study, consult reliable educational and government sources. The following references provide sound background on water chemistry, equilibrium, and acid-base behavior:
- LibreTexts Chemistry for general equilibrium and weak acid-base theory.
- U.S. Environmental Protection Agency (.gov) for water chemistry and pH background used in environmental contexts.
- University of California, Berkeley Chemistry (.edu) for foundational chemical principles and instructional resources.
Final answer
Using the standard values Ka = 1.8 × 10^-5 for acetic acid and Kw = 1.0 × 10^-14 at 25 C, the acetate ion has Kb = 5.56 × 10^-10. For a 0.15 M sodium acetate solution, the hydroxide concentration is approximately 9.13 × 10^-6 M, the pOH is about 5.04, and the resulting pH is about 8.96.
These values are standard educational estimates. Slight differences can occur depending on the exact Ka selected for acetic acid, the temperature, ionic strength corrections, and rounding conventions.