Calculate pH of 0.11 M Solution of CH3COONa Solution
Use this interactive sodium acetate pH calculator to find the pH, pOH, hydroxide concentration, and hydrolysis extent for CH3COONa in water. The calculator uses acetate hydrolysis and can solve the equilibrium with an exact quadratic method for premium accuracy.
Results
Enter values and click Calculate pH to see the equilibrium analysis for CH3COONa.
How to calculate pH of 0.11 M solution of CH3COONa solution
When you need to calculate the pH of a 0.11 M solution of CH3COONa solution, you are really solving the hydrolysis of the acetate ion in water. Sodium acetate, written as CH3COONa, is a salt formed from a strong base, NaOH, and a weak acid, CH3COOH. Because the cation Na+ is essentially neutral in water, the chemistry is controlled by the acetate anion, CH3COO-. That acetate ion acts as a weak base and reacts with water to generate a small amount of OH-.
This reaction makes the solution basic, so the pH of sodium acetate is greater than 7 at standard conditions. For a 0.11 M sodium acetate solution, the pH is usually about 8.89 at 25 C when the acetic acid dissociation constant is taken as 1.8 x 10^-5 and Kw is 1.0 x 10^-14. The exact value can vary slightly depending on the Ka source, temperature, ionic strength, and rounding convention used.
Step by step chemistry behind the calculation
The cleanest route is to convert the acid constant of acetic acid into the base constant of acetate. Since acetate is the conjugate base of acetic acid, the relationship is:
Using typical 25 C constants:
- Ka for acetic acid = 1.8 x 10^-5
- Kw for water = 1.0 x 10^-14
So:
Now let the initial acetate concentration be 0.11 M. If x is the amount of OH- produced at equilibrium, then:
- [CH3COO-]initial = 0.11
- [CH3COOH]initial = 0
- [OH-]initial = 0, for the hydrolysis setup
At equilibrium:
- [CH3COO-] = 0.11 – x
- [CH3COOH] = x
- [OH-] = x
Substitute into the Kb expression:
Because Kb is very small, x is much smaller than 0.11, so many textbook problems use the approximation:
This gives:
- pOH = -log(7.82 x 10^-6) ≈ 5.11
- pH = 14.00 – 5.11 ≈ 8.89
Why CH3COONa gives a basic solution
Students often ask why sodium acetate is basic if it is called a salt. The answer comes from the parent acid and base used to form it. Acetic acid is weak, but sodium hydroxide is strong. In water, Na+ does not appreciably react. Acetate, however, can accept a proton from water. This proton acceptance leaves behind OH-, increasing the basicity of the solution. The basicity is mild, not extreme, because acetate is still only a weak base.
This is the central concept to remember: salts of strong bases and weak acids usually produce basic solutions. By contrast, salts of strong acids and weak bases typically give acidic solutions, while salts of strong acids and strong bases are generally neutral.
Exact method versus approximation
For dilute weak equilibria, the approximation x << C is often sufficient. But a premium calculator should also support the exact quadratic solution. Starting from:
Rearranging gives:
The physically meaningful root is:
For 0.11 M sodium acetate, the exact and approximate answers are nearly identical because the degree of hydrolysis is tiny. The percent hydrolysis is only a few thousandths of a percent. That means the initial concentration barely changes during equilibrium establishment.
| Method | [OH-] (M) | pOH | pH at 25 C | Difference |
|---|---|---|---|---|
| Approximation, x = sqrt(KbC) | 7.82 x 10^-6 | 5.107 | 8.893 | Reference |
| Exact quadratic | 7.82 x 10^-6 | 5.107 | 8.893 | Less than 0.001 pH unit |
Important constants and what they mean
If you want consistent pH calculations, you should know the key constants involved:
- Ka: acid dissociation constant of acetic acid. Common textbook values range around 1.75 x 10^-5 to 1.8 x 10^-5 at 25 C.
- Kb: base dissociation constant of acetate, found using Kb = Kw / Ka.
- Kw: ion product of water, equal to 1.0 x 10^-14 at 25 C, but temperature dependent.
- pH and pOH: related by pH + pOH = 14 at 25 C. At other temperatures the sum changes slightly if you use a more rigorous treatment.
Temperature matters because Kw changes with temperature, and equilibrium constants can shift as well. In practical educational settings, however, 25 C values are usually assumed unless the problem explicitly states otherwise.
| Quantity | Typical 20 C Value | Typical 25 C Value | Typical 30 C Value | Why it matters |
|---|---|---|---|---|
| Kw | 6.81 x 10^-15 | 1.00 x 10^-14 | 1.47 x 10^-14 | Changes the calculated Kb and pH |
| Neutral pH | About 7.08 | 7.00 | About 6.92 | Neutrality shifts with temperature |
| Ka of acetic acid | Near 1.8 x 10^-5 | Near 1.8 x 10^-5 | Near 1.8 x 10^-5 | Source dependent, slight variation in references |
Worked example for 0.11 M CH3COONa
Here is the complete worked example in compact form:
- Write the hydrolysis reaction: CH3COO- + H2O ⇌ CH3COOH + OH-
- Use Ka = 1.8 x 10^-5 and Kw = 1.0 x 10^-14
- Calculate Kb = 5.56 x 10^-10
- Set initial concentration C = 0.11 M
- Use x ≈ sqrt(KbC) = sqrt(5.56 x 10^-10 x 0.11)
- Find x = [OH-] ≈ 7.82 x 10^-6 M
- Calculate pOH = 5.11
- Calculate pH = 14.00 – 5.11 = 8.89
That is the standard chemistry answer expected in general chemistry and introductory analytical chemistry courses.
What if your textbook uses a different Ka?
Some textbooks and data tables use Ka = 1.75 x 10^-5 instead of 1.8 x 10^-5. That small difference slightly changes Kb and therefore slightly changes pH. The resulting pH may shift by only a few thousandths to hundredths of a unit. In grading, that is usually acceptable if the method is correct and the assumptions are clearly stated.
Common mistakes when solving sodium acetate pH problems
- Using Ka directly without converting to Kb. For a salt of a weak acid, the base hydrolysis constant is needed.
- Treating sodium acetate as a strong base. It is basic, but only weakly basic.
- Forgetting that Na+ is a spectator ion. The sodium ion does not control pH here.
- Mixing up pH and pOH. Since OH- is produced, pOH is found first.
- Ignoring temperature assumptions. Kw is not fixed at all temperatures.
- Rounding too early. Keep at least three significant digits until the final pH step.
How the chart in this calculator helps
The interactive chart above visualizes equilibrium composition after the sodium acetate solution hydrolyzes. It typically compares the remaining acetate concentration, the small amount of acetic acid formed, and the hydroxide concentration generated. This is useful because many learners understand equilibrium better when they see the dramatic scale difference between the dominant species and the very small hydrolysis products.
In a 0.11 M solution, nearly all of the acetate remains as acetate. Only a tiny fraction converts to acetic acid and hydroxide. That is why the pH is only mildly basic rather than strongly alkaline.
Real world significance of sodium acetate pH
Sodium acetate is widely used in laboratory buffer preparation, food processing, textile work, and biochemical systems. Its chemistry becomes especially important when paired with acetic acid to create acetate buffers. In buffer design, understanding the pH of sodium acetate alone is the first step toward understanding how the acetate-acetic acid conjugate pair resists pH change.
Acetate systems are also common in educational laboratories because they provide a clear and safe demonstration of weak acid and weak base equilibrium principles. Measuring the pH of sodium acetate and comparing it to theory is a classic exercise in verifying equilibrium models.
Authoritative references for acid-base and pH concepts
If you want to validate the equilibrium principles or review broader pH concepts, these sources are useful:
- U.S. Environmental Protection Agency: alkalinity and acid-base context
- National Center for Biotechnology Information: pH fundamentals in aqueous systems
- University of Wisconsin: acid-base equilibrium tutorial
Quick answer summary
To calculate the pH of a 0.11 M solution of CH3COONa solution, treat acetate as a weak base. Compute Kb from the Ka of acetic acid, solve for hydroxide concentration, find pOH, and then convert to pH. At 25 C with Ka = 1.8 x 10^-5, the final pH is about 8.89. That value shows sodium acetate is a mildly basic salt, exactly as expected for the conjugate base of a weak acid.