How to Calculate pH of CH3COONa
Use this premium sodium acetate pH calculator to estimate the pH of a CH3COONa solution from concentration, pKa of acetic acid, and pKw. The tool applies salt hydrolysis for the conjugate base of a weak acid and displays the full chemistry behind the answer.
CH3COONa pH Calculator
Enter your values and click Calculate pH to see the hydrolysis result for CH3COONa.
pH vs Concentration Trend
The chart shows how the pH of sodium acetate changes with concentration using your selected pKa and pKw values.
Expert Guide: How to Calculate pH of CH3COONa
To calculate the pH of CH3COONa, you need to recognize what this compound actually is in water. CH3COONa is sodium acetate, a salt formed from a strong base (NaOH) and a weak acid (CH3COOH, acetic acid). When sodium acetate dissolves, the sodium ion is essentially a spectator ion, while the acetate ion reacts with water. That hydrolysis produces a small amount of OH–, which makes the solution basic. This is why a sodium acetate solution normally has a pH greater than 7.
Students often make the mistake of treating CH3COONa as if it were a strong base. It is not. It is a basic salt, and its alkalinity comes from the equilibrium behavior of acetate ions in water. The most reliable way to solve the problem is to calculate the base dissociation constant of acetate from the acid dissociation constant of acetic acid, then use the salt concentration to estimate the hydroxide ion concentration.
Step 1: Write the Dissociation and Hydrolysis Reactions
In water, sodium acetate dissociates as follows:
The acetate ion then undergoes hydrolysis:
This second equation is the one that controls the pH. Since OH– is produced, the solution becomes basic.
Step 2: Find Kb from Ka or pKa
Because acetate is the conjugate base of acetic acid, its base dissociation constant is related to the acid dissociation constant of acetic acid by:
If pKa is given instead of Ka, convert with:
At 25 degrees C, acetic acid has a pKa close to 4.76, so:
With Kw = 1.0 × 10-14:
Step 3: Set Up the Hydrolysis Expression
If the initial concentration of sodium acetate is C, then the initial acetate concentration is also approximately C because sodium acetate dissociates completely. Let x be the amount that reacts with water:
- Initial [CH3COO–] = C
- Change = -x
- Equilibrium [CH3COO–] = C – x
- Equilibrium [CH3COOH] = x
- Equilibrium [OH–] = x
The equilibrium expression is:
If x is small compared with C, you can use the common approximation:
Since x equals [OH–], you then calculate:
Worked Example: 0.10 M CH3COONa
Suppose you want to calculate the pH of a 0.10 M sodium acetate solution at 25 degrees C.
- Use pKa = 4.76 for acetic acid.
- Calculate Ka = 10-4.76 ≈ 1.74 × 10-5.
- Find Kb = 10-14 / 1.74 × 10-5 ≈ 5.75 × 10-10.
- Use the approximation x = √(Kb × C).
So the pH of a 0.10 M CH3COONa solution is about 8.88.
Shortcut Formula for Sodium Acetate
When the hydrolysis approximation is valid, there is a convenient shortcut. Starting from the weak base relation and converting to pH form gives:
Since pKb = pKw – pKa, you can also write:
For 0.10 M sodium acetate:
This shortcut is extremely useful in exams, but it should only be used when the approximation is justified.
When to Use the Approximation and When to Use the Quadratic Equation
For most typical sodium acetate concentrations used in laboratory calculations, the hydrolysis is weak enough that x is much smaller than C. In those cases, the square root approximation gives a very accurate answer. However, if the solution is extremely dilute, the approximation becomes less reliable and the full quadratic approach is better.
The exact hydrolysis equation is:
Solving for the positive root:
That value of x is the exact equilibrium [OH–] from acetate hydrolysis. The calculator above can use either method.
Comparison Table: Typical pH Values for CH3COONa Solutions at 25 Degrees C
The table below uses pKa = 4.76 and pKw = 14.00. These values are representative for room-temperature calculations in many chemistry courses and laboratory settings.
| CH3COONa Concentration (M) | Calculated [OH-] (M) | pOH | pH | Interpretation |
|---|---|---|---|---|
| 1.00 | 2.40 × 10^-5 | 4.62 | 9.38 | Clearly basic |
| 0.10 | 7.58 × 10^-6 | 5.12 | 8.88 | Mildly basic |
| 0.010 | 2.40 × 10^-6 | 5.62 | 8.38 | Basic, but less alkaline |
| 0.0010 | 7.58 × 10^-7 | 6.12 | 7.88 | Slightly basic |
| 0.00010 | 2.40 × 10^-7 | 6.62 | 7.38 | Approaching neutral range |
Hydrolysis Strength and Percent Conversion Data
Another useful way to understand sodium acetate is to compare how much of the acetate actually hydrolyzes. Even though the pH is basic, only a tiny fraction of acetate ions react with water in ordinary solutions. This is why the weak-base approximation works so well.
| CH3COONa Concentration (M) | [OH-] from Hydrolysis (M) | Fraction Hydrolyzed, x/C | Percent Hydrolyzed | Approximation Quality |
|---|---|---|---|---|
| 1.00 | 2.40 × 10^-5 | 2.40 × 10^-5 | 0.0024% | Excellent |
| 0.10 | 7.58 × 10^-6 | 7.58 × 10^-5 | 0.0076% | Excellent |
| 0.010 | 2.40 × 10^-6 | 2.40 × 10^-4 | 0.024% | Excellent |
| 0.0010 | 7.58 × 10^-7 | 7.58 × 10^-4 | 0.076% | Very good |
| 0.00010 | 2.40 × 10^-7 | 2.40 × 10^-3 | 0.24% | Still good |
Why CH3COONa Gives a Basic Solution
The sodium ion, Na+, does not significantly react with water because it comes from the strong base sodium hydroxide. The acetate ion, CH3COO–, is different. It is the conjugate base of acetic acid, so it can accept a proton from water. This proton transfer generates hydroxide ion. In acid-base terms, any salt of a weak acid and strong base tends to produce a basic aqueous solution.
- Strong acid + strong base salt: usually neutral
- Weak acid + strong base salt: usually basic
- Strong acid + weak base salt: usually acidic
- Weak acid + weak base salt: depends on Ka and Kb
Common Mistakes in Sodium Acetate pH Calculations
- Using the concentration directly as [OH–]. Sodium acetate is not a strong base, so [OH–] is not equal to the salt concentration.
- Using Ka instead of Kb in the hydrolysis equation. Acetate behaves as a base in water, so you must use Kb for the hydrolysis step.
- Forgetting to convert from pOH to pH. Once you calculate [OH–], you usually get pOH first, not pH directly.
- Ignoring temperature effects. The common values pKa = 4.76 and pKw = 14.00 are temperature dependent. For high-precision work, use temperature-appropriate constants.
- Applying the approximation in very dilute systems without checking. In very low concentrations, a quadratic treatment can be safer.
Fast Exam Strategy for CH3COONa Problems
If you are solving an exam problem and the concentration is not extremely low, this is the fastest reliable workflow:
- Identify CH3COONa as the salt of a weak acid and strong base.
- Conclude the solution is basic.
- Use pKa of acetic acid to find Ka.
- Compute Kb = Kw / Ka.
- Use [OH–] ≈ √(KbC).
- Calculate pOH and then pH.
This approach is accepted in most introductory chemistry and analytical chemistry settings unless the problem explicitly asks for an exact equilibrium treatment.
Authoritative Chemistry References
For deeper study of acid-base equilibria, weak acids, conjugate bases, and thermodynamic constants, review these authoritative educational and government resources:
- NIST Chemistry WebBook for trusted chemical data and equilibrium-related reference values.
- LibreTexts Chemistry is popular, but if you need only .edu or .gov references, compare with university-hosted equilibrium notes such as University of Wisconsin Department of Chemistry.
- Michigan State University Chemistry for acid-base equilibrium teaching materials and general chemistry support.
Final Takeaway
Calculating the pH of CH3COONa is fundamentally a weak base hydrolysis problem. The acetate ion reacts with water to create OH–, so the pH is above 7. The core method is simple: convert pKa of acetic acid to Ka, compute Kb of acetate, estimate [OH–] from the salt concentration, then convert pOH to pH. For a typical 0.10 M sodium acetate solution, the pH is about 8.88. Once you understand that CH3COONa behaves as the conjugate base of a weak acid, the calculation becomes systematic and easy to reproduce.