Calculate the pH of a 0.10 M NaCH3CO2 Solution
This interactive calculator finds the pH of sodium acetate solution by treating acetate as a weak base in water. Enter concentration, choose the acetic acid Ka value, and get pH, pOH, Kb, hydroxide concentration, and a chart showing how pH changes with concentration.
Sodium Acetate pH Calculator
Result
For a 0.10 M sodium acetate solution using Ka = 1.8 × 10^-5, the solution is mildly basic because acetate reacts with water to form OH-.
Expert Guide: How to Calculate the pH of a 0.10 M NaCH3CO2 Solution
If you need to calculate the pH of a 0.10 M NaCH3CO2 solution, you are working with a classic weak-base hydrolysis problem. NaCH3CO2 is sodium acetate, the sodium salt of acetic acid. In water, sodium acetate dissociates completely into sodium ions and acetate ions. The sodium ion is essentially a spectator ion in this context, but the acetate ion matters a great deal because it is the conjugate base of a weak acid. That means acetate can react with water and generate hydroxide ions, making the solution basic.
The most important conceptual point is that sodium acetate is not itself a strong base like sodium hydroxide. Instead, it behaves as a basic salt because one of its ions, CH3CO2-, accepts a proton from water. This creates acetic acid and hydroxide:
Because hydroxide ions are produced, the pH ends up above 7. At 25 degrees C, a 0.10 M sodium acetate solution has a pH of about 8.87 when you use the common acetic acid dissociation constant Ka = 1.8 × 10^-5. Below, you will see exactly where that answer comes from, why it works, and how to avoid the most common mistakes.
Step 1: Recognize that sodium acetate is the salt of a weak acid
Acetic acid, CH3COOH, is a weak acid. Its sodium salt, NaCH3CO2, dissociates completely in water:
The acetate ion is the conjugate base of acetic acid. Conjugate bases of weak acids are basic in water. Therefore, once sodium acetate dissolves, the equilibrium that controls pH is the reaction of acetate with water. This is why you do not calculate pH directly from an acid Ka expression. Instead, you first convert Ka into Kb for the acetate ion.
Step 2: Use the relationship between Ka, Kb, and Kw
The acid dissociation constant and base dissociation constant of a conjugate pair are linked by:
At 25 degrees C, Kw = 1.0 × 10^-14. If acetic acid has Ka = 1.8 × 10^-5, then the base dissociation constant of acetate is:
Kb = (1.0 × 10^-14) / (1.8 × 10^-5)
Kb = 5.56 × 10^-10
This small Kb value tells you acetate is a weak base, which means the hydroxide concentration produced is much smaller than the initial 0.10 M acetate concentration.
Step 3: Set up the ICE table
The equilibrium reaction is:
Start with 0.10 M acetate and assume essentially zero acetic acid and hydroxide from the salt itself.
| Species | Initial (M) | Change (M) | Equilibrium (M) |
|---|---|---|---|
| CH3CO2- | 0.10 | -x | 0.10 – x |
| CH3COOH | 0 | +x | x |
| OH- | 0 | +x | x |
Now apply the equilibrium expression:
5.56 × 10^-10 = x^2 / (0.10 – x)
Step 4: Solve for hydroxide concentration
Because Kb is small, many instructors allow the approximation 0.10 – x ≈ 0.10. Then:
x^2 = 5.56 × 10^-11
x = 7.45 × 10^-6 M
So the hydroxide concentration is approximately 7.45 × 10^-6 M. If you solve the quadratic exactly, you get nearly the same value. This confirms the approximation is excellent here because x is much smaller than 0.10 M.
Step 5: Convert hydroxide concentration to pOH and pH
Once you know [OH-], the rest is straightforward:
pOH = -log(7.45 × 10^-6)
pOH ≈ 5.13
pH = 14.00 – 5.13 = 8.87
That is the final answer for the standard textbook version of this problem:
Why the pH is not 7 and not strongly basic
Students sometimes expect every sodium salt to be neutral because many common sodium salts, such as NaCl, are neutral in water. Sodium acetate is different because acetate is the conjugate base of a weak acid. On the other hand, it is not strongly basic because acetate does not dissociate like OH- from a strong base. It only partially reacts with water. As a result, the pH is above 7 but still modest, typically in the high 8 range for a 0.10 M solution.
Comparison table: how sodium acetate differs from other sodium salts
| Salt | Parent acid/base | Expected aqueous behavior | Typical pH trend at moderate concentration |
|---|---|---|---|
| NaCl | Strong acid + strong base | Essentially neutral | Near 7 |
| NaCH3CO2 | Weak acid + strong base | Basic by anion hydrolysis | Above 7, often around 8 to 9 |
| NH4Cl | Strong acid + weak base | Acidic by cation hydrolysis | Below 7 |
| NaOH | Strong base | Strongly basic | Much greater than 7 |
Approximation versus exact solution
For many weak acid and weak base calculations, the shortcut x = √(KC) is accurate enough if the percent ionization is low. In this problem:
- Kb = 5.56 × 10^-10
- C = 0.10 M
- x = √(5.56 × 10^-10 × 0.10) = 7.45 × 10^-6 M
The percentage of acetate that hydrolyzes is:
That is far below 5%, so the approximation is clearly valid. Exact quadratic solving gives virtually the same pH to two decimal places. In practical classroom settings, both methods are acceptable as long as your instructor has not specifically required the quadratic formula.
How concentration affects pH
As sodium acetate concentration increases, pH rises, but not in a simple one-to-one linear way. Because weak-base hydrolysis depends on the square root relationship in the approximation, pH changes more gradually than many students initially expect. The table below shows representative values using Ka = 1.8 × 10^-5 at 25 degrees C.
| NaCH3CO2 concentration (M) | Kb of acetate | Approximate [OH-] (M) | Approximate pH |
|---|---|---|---|
| 0.001 | 5.56 × 10^-10 | 7.45 × 10^-7 | 7.87 |
| 0.010 | 5.56 × 10^-10 | 2.36 × 10^-6 | 8.37 |
| 0.10 | 5.56 × 10^-10 | 7.45 × 10^-6 | 8.87 |
| 1.00 | 5.56 × 10^-10 | 2.36 × 10^-5 | 9.37 |
This table reveals a useful pattern: every tenfold increase in concentration raises the pH by roughly 0.5 units for this weak-base system. That happens because [OH-] is proportional to the square root of concentration, not directly proportional to concentration itself.
Common mistakes when solving this problem
- Treating sodium acetate as a strong base. It is not. Only hydroxides of strong bases directly release OH- in full stoichiometric amounts.
- Using Ka directly to compute pH. You need Kb for acetate, so convert using Kb = Kw / Ka.
- Forgetting that sodium is a spectator ion. Na+ does not significantly alter pH in this calculation.
- Using the wrong equilibrium reaction. The hydrolysis of CH3CO2- with water is the correct equilibrium to write.
- Skipping the pOH step. Weak-base problems usually give [OH-] first, so you must calculate pOH and then convert to pH.
Real-world relevance of sodium acetate solution pH
Sodium acetate appears in analytical chemistry, biological buffers, textile applications, food processing, and deicing formulations. In many of these settings, understanding the pH of acetate-containing solutions matters because acetate can participate in buffering systems with acetic acid. While a pure sodium acetate solution is basic, a mixture of sodium acetate and acetic acid forms a buffer that can resist pH changes. This makes sodium acetate more important than it may first appear in introductory chemistry problems.
Authoritative references for acid-base data and equilibrium concepts
For deeper study, consult authoritative educational and scientific sources. Useful references include the National Institute of Standards and Technology, acid-base resources from chemistry educational materials hosted by universities and colleges, and chemistry course notes from major universities such as Princeton University. You can also explore federal science resources through PubChem at the U.S. National Institutes of Health.
Final takeaway
To calculate the pH of a 0.10 M NaCH3CO2 solution, identify acetate as a weak base, convert acetic acid Ka into acetate Kb, solve for hydroxide ion concentration, and then convert pOH to pH. Using Ka = 1.8 × 10^-5 at 25 degrees C gives Kb = 5.56 × 10^-10, [OH-] ≈ 7.45 × 10^-6 M, pOH ≈ 5.13, and pH ≈ 8.87. That value is the standard accepted result in most general chemistry settings.
If you want a fast answer, use the calculator above. If you want a deep understanding, follow the full hydrolysis setup and equilibrium logic outlined here. Either way, the chemistry tells the same story: sodium acetate makes water mildly basic because acetate is the conjugate base of a weak acid.