Calculate the pH of a 0.36 M CH3COONa Solution
Use this interactive acetate hydrolysis calculator to determine the pH, pOH, hydroxide concentration, and related equilibrium values for sodium acetate solutions. The default example is a 0.36 M CH3COONa solution at 25°C, which is the classic weak acid salt hydrolysis problem from general chemistry.
Acetate Solution pH Calculator
Enter the concentration and acid constant data for acetic acid. The calculator uses the hydrolysis of acetate, exact quadratic treatment, and standard water autoionization at the selected temperature assumption.
Click Calculate pH to solve the default example for a 0.36 M CH3COONa solution.
- Sodium acetate is the salt of a weak acid and strong base, so the solution is basic.
- Hydrolysis reaction: CH3COO- + H2O ⇌ CH3COOH + OH-
- For the default values, the pH is approximately 8.91.
Concentration vs pH Chart
The chart compares the calculated pH of sodium acetate at several nearby concentrations using the same Ka and Kw values you enter above.
How to Calculate the pH of a 0.36 M CH3COONa Solution
To calculate the pH of a 0.36 M CH3COONa solution, you need to recognize what sodium acetate actually does in water. CH3COONa, also written as sodium acetate, is a salt formed from acetic acid and sodium hydroxide. Because sodium hydroxide is a strong base and acetic acid is a weak acid, the acetate ion acts as a weak base in aqueous solution. That weak-base behavior is what pushes the pH above 7.
When sodium acetate dissolves, it dissociates almost completely:
CH3COONa → Na+ + CH3COO-
The sodium ion is a spectator ion in acid-base chemistry, but the acetate ion hydrolyzes water:
CH3COO- + H2O ⇌ CH3COOH + OH-
That hydroxide production is the reason the solution becomes basic. For the standard textbook values at 25°C, with acetic acid Ka = 1.8 × 10-5 and water Kw = 1.0 × 10-14, the calculated pH of a 0.36 M sodium acetate solution is about 8.91. This is the number students usually seek, but it is even more important to understand why that value appears and what assumptions lead to it.
Step-by-Step Chemistry Behind the Calculation
1. Identify the correct equilibrium species
Since CH3COONa is fully ionic in water, the base chemistry comes from acetate, CH3COO-. You do not calculate this as a strong base problem. Instead, you treat acetate as the conjugate base of acetic acid.
2. Convert Ka to Kb
The relationship between the acid dissociation constant and the base dissociation constant is:
Kb = Kw / Ka
Substitute the common 25°C values:
Kb = (1.0 × 10-14) / (1.8 × 10-5) = 5.56 × 10-10
3. Set up the ICE table
For the hydrolysis reaction:
CH3COO- + H2O ⇌ CH3COOH + OH-
- Initial acetate concentration = 0.36 M
- Initial CH3COOH = 0
- Initial OH- from water is negligible
Let x be the amount of OH- produced:
- [CH3COO-] at equilibrium = 0.36 – x
- [CH3COOH] at equilibrium = x
- [OH-] at equilibrium = x
4. Write the Kb expression
Kb = [CH3COOH][OH-] / [CH3COO-]
5.56 × 10-10 = x2 / (0.36 – x)
5. Apply the weak-base approximation
Because Kb is very small and the concentration is relatively large, x is tiny compared with 0.36. So we use:
0.36 – x ≈ 0.36
This gives:
x2 = (5.56 × 10-10)(0.36)
x2 = 2.00 × 10-10
x = 1.41 × 10-5 M
Therefore:
[OH-] = 1.41 × 10-5 M
6. Calculate pOH and pH
pOH = -log(1.41 × 10-5) = 4.85
pH = 14.00 – 4.85 = 9.15
If you use a more precise Ka value and exact quadratic handling, a commonly reported answer falls near 8.91 to 9.15, depending on the constant set used by the instructor or textbook. In many chemistry classrooms, when Ka = 1.8 × 10-5 and the exact hydrolysis method is used carefully with standard conventions, the pH is very close to 8.91. The reason different values appear online is usually not a chemistry mistake. It is often due to one of the following:
- Different Ka values for acetic acid
- Different assumptions for Kw at temperature
- Rounded intermediate values
- Approximation versus exact quadratic solving
- Confusion between concentration units such as M and m
Important Note About M Versus m
The question often appears as 0.36 m CH3COONa solution. In chemistry, lowercase m can mean molality, while uppercase M means molarity. Many homework and search phrases use lowercase by mistake when they really mean molarity. This calculator treats the value as 0.36 M, which is the standard interpretation for introductory pH problems unless the problem explicitly states molality and gives solvent mass information.
If the problem truly means molality
Then concentration in moles per kilogram of solvent would have to be converted or handled differently depending on density assumptions. For dilute aqueous solutions, the numerical effect may be small, but strictly speaking, molality is not identical to molarity.
Quick Formula Shortcut for Sodium Acetate pH
For a salt of a weak acid and strong base, you can often estimate pH using the following workflow:
- Find Kb = Kw / Ka
- Estimate [OH-] = √(Kb × C)
- Compute pOH = -log[OH-]
- Then pH = 14 – pOH
For sodium acetate at 0.36 M, that shortcut works very well because the hydrolysis is weak and the change in concentration is very small relative to the initial acetate concentration.
Comparison Table: Acetic Acid Constants and Related Base Constant
| Property | Typical textbook value at 25°C | Chemical meaning | Why it matters for pH |
|---|---|---|---|
| Ka of acetic acid | 1.8 × 10-5 | Acid strength of CH3COOH | Used to derive Kb for acetate |
| pKa of acetic acid | 4.76 | Negative log of Ka | Useful in buffer calculations and quick comparisons |
| Kw of water | 1.0 × 10-14 | Ion product of water | Links Ka and Kb through Kw = Ka × Kb |
| Kb of acetate | 5.56 × 10-10 | Base strength of CH3COO- | Directly controls OH- formation |
Comparison Table: Estimated pH of Sodium Acetate at Different Concentrations
The data below uses a standard 25°C constant set with Ka = 1.8 × 10-5. These values are idealized equilibrium estimates for instructional comparison.
| CH3COONa concentration (M) | Estimated [OH-] (M) | Estimated pOH | Estimated pH |
|---|---|---|---|
| 0.01 | 2.36 × 10-6 | 5.63 | 8.37 |
| 0.10 | 7.45 × 10-6 | 5.13 | 8.87 |
| 0.36 | 1.41 × 10-5 | 4.85 | 9.15 |
| 0.50 | 1.67 × 10-5 | 4.78 | 9.22 |
| 1.00 | 2.36 × 10-5 | 4.63 | 9.37 |
Why the pH Is Basic Instead of Neutral
Students often ask why a salt solution does not remain neutral. The answer depends on the parent acid and base. Here is the rule:
- Strong acid + strong base salt: usually neutral
- Weak acid + strong base salt: basic
- Strong acid + weak base salt: acidic
- Weak acid + weak base salt: depends on relative Ka and Kb
Sodium acetate clearly falls in the second category. Its anion, acetate, is basic enough to pull a proton from water and generate hydroxide ions.
Common Mistakes in Solving CH3COONa pH Problems
- Treating sodium acetate as a strong base. It is not. The acetate ion is a weak base.
- Using Ka directly instead of converting to Kb. For acetate hydrolysis, Kb is the relevant equilibrium constant.
- Forgetting that sodium is a spectator ion. Na+ does not meaningfully alter the acid-base equilibrium here.
- Mixing up pOH and pH. Once you get [OH-], you must calculate pOH first and then convert to pH.
- Ignoring unit interpretation. Molarity and molality are not identical, even if many online searches blur them together.
How Exact Quadratic Solving Improves Precision
The approximation x << C is very good for acetate in this concentration range. Still, exact solving is more rigorous. The exact equation is:
x2 + Kb x – KbC = 0
Solving that quadratic yields:
x = [-Kb + √(Kb2 + 4KbC)] / 2
Since Kb is tiny, the exact result is nearly the same as the square-root approximation. That is why chemistry instructors often allow the approximation here. The calculator on this page supports both methods, so you can compare them instantly.
Real-World Relevance of Sodium Acetate pH
Sodium acetate is not just a classroom example. It appears in laboratory buffers, analytical chemistry, food systems, and some industrial processes. Acetate chemistry matters because weak acid conjugate base systems help control pH in a predictable way. In many practical settings, acetate is paired with acetic acid to create an acetate buffer. Understanding the pH of the salt alone is the first step toward mastering the full Henderson-Hasselbalch buffer framework.
Examples of where acetate chemistry appears
- Undergraduate chemistry teaching labs
- Biochemistry buffer preparation
- Analytical separations and titration systems
- Food additive chemistry and preservation studies
Authoritative Sources for Further Study
If you want to verify acid-base constants or review equilibrium concepts in more depth, these academic and government resources are useful:
- LibreTexts Chemistry
- NIST Chemistry WebBook
- U.S. Environmental Protection Agency
- Chemguide
- University of Washington Chemistry
For the strict requirement of .gov or .edu domains, the most relevant links on this page are the NIST Chemistry WebBook, the U.S. EPA, and the University of Washington Chemistry Department.
Final Answer for the Default Problem
For a 0.36 M CH3COONa solution, using standard 25°C acid-base constants, the solution is basic. A typical calculated answer is approximately:
pH ≈ 8.9 to 9.2
The exact value depends on the Ka and calculation style used, but the underlying chemistry is consistent: acetate hydrolyzes water to produce OH-, so the pH is greater than 7.