Calculate pH of 0.137 M Sodium Hexanoate NaC6H11O2
Use this premium weak-base hydrolysis calculator to estimate the pH, pOH, hydroxide concentration, and base dissociation constant for aqueous sodium hexanoate solutions.
Calculator
Enter the sodium hexanoate concentration and the acid pKa for hexanoic acid. The tool assumes complete dissociation of the sodium salt and weak-base hydrolysis of the hexanoate ion.
Results
Click Calculate pH to solve for the pH of 0.137 M sodium hexanoate NaC6H11O2.
Expert Guide: How to Calculate the pH of 0.137 M Sodium Hexanoate NaC6H11O2
Calculating the pH of a salt derived from a weak acid is a classic acid-base equilibrium problem in general chemistry, analytical chemistry, and solution chemistry. In this case, the compound is sodium hexanoate, written as NaC6H11O2. When dissolved in water, sodium hexanoate dissociates essentially completely into sodium ions and hexanoate ions:
The sodium ion is a spectator ion in this context, so the chemistry that matters is the behavior of the hexanoate ion, C6H11O2-. Hexanoate is the conjugate base of hexanoic acid, a weak carboxylic acid. Since it is the conjugate base of a weak acid, it reacts with water and produces hydroxide:
That reaction is the reason the solution becomes basic. If the concentration of sodium hexanoate is 0.137 M, then the initial concentration of hexanoate ion is also approximately 0.137 M, assuming ideal behavior and full dissociation of the ionic salt. The goal is to determine the resulting hydroxide concentration and convert it into pOH and pH.
Step 1: Identify the correct equilibrium constant
For weak base hydrolysis, the relevant constant is Kb, not Ka. However, reference data for organic acids are usually listed as Ka or pKa values for the parent weak acid. For hexanoic acid, a commonly used pKa at 25°C is about 4.85. To convert this to Ka:
The relationship between Ka and Kb for a conjugate acid-base pair is:
At 25°C, Kw = 1.0 × 10^-14, so:
This small Kb value tells you that hexanoate is a weak base, but not a negligible one. Even though the base is weak, a 0.137 M solution contains enough dissolved ions to create a measurable basic pH.
Step 2: Set up the ICE table
The hydrolysis reaction is:
An ICE table for this equilibrium looks like this:
- Initial: [C6H11O2-] = 0.137 M, [HC6H11O2] = 0, [OH-] = 0
- Change: -x, +x, +x
- Equilibrium: 0.137 – x, x, x
Plugging these into the Kb expression:
Since Kb is much smaller than the initial concentration, many textbooks allow the approximation 0.137 – x ≈ 0.137. That gives:
Because x is the hydroxide concentration:
Then:
So the pH of a 0.137 M sodium hexanoate solution is approximately 8.99 at 25°C when using pKa = 4.85. The quadratic method gives virtually the same answer because x is tiny relative to 0.137 M.
Why the solution is basic
Many students initially assume that a sodium salt should be neutral because sodium compounds like sodium chloride are neutral in water. The difference is the anion. Chloride is the conjugate base of a strong acid, so it does not appreciably react with water. Hexanoate, however, is the conjugate base of a weak acid, so it does react with water and generates OH-. That shift toward hydroxide raises the pH above 7.
Comparison table: sodium salts and their expected acid-base behavior
| Salt | Parent acid | Acid strength of parent acid | Expected aqueous behavior | Typical pH trend |
|---|---|---|---|---|
| Sodium chloride, NaCl | HCl | Strong acid | Essentially no hydrolysis | Near neutral, around 7 |
| Sodium acetate, CH3COONa | Acetic acid | Weak acid, pKa about 4.76 | Basic due to acetate hydrolysis | Above 7 |
| Sodium hexanoate, NaC6H11O2 | Hexanoic acid | Weak acid, pKa about 4.85 | Basic due to hexanoate hydrolysis | About 8.99 at 0.137 M |
| Ammonium chloride, NH4Cl | NH4+ | Conjugate acid of weak base | Acidic due to ammonium ion | Below 7 |
Exact versus approximate calculation
The weak-base approximation is usually valid when the percent ionization is small, often under 5%. For 0.137 M sodium hexanoate, the estimated hydroxide concentration is only about 9.85 × 10^-6 M. The percent hydrolysis is therefore:
That is far below 5%, so the approximation is excellent. Still, a premium calculator should be able to perform the full quadratic solution:
Solving for the positive root gives:
Because Kb² is tiny compared with 4KbC, the answer is almost identical to the square-root shortcut. The quadratic method is useful when concentrations are lower, equilibrium constants are larger, or your instructor specifically asks for an exact solution.
Data table: numerical summary for 0.137 M sodium hexanoate
| Quantity | Value | How obtained |
|---|---|---|
| Salt concentration | 0.137 M | Given problem statement |
| pKa of hexanoic acid | 4.85 | Reference value at 25°C |
| Ka | 1.41 × 10^-5 | 10^-4.85 |
| Kb | 7.08 × 10^-10 | Kw / Ka |
| [OH-] | 9.85 × 10^-6 M | √(KbC) or quadratic |
| pOH | 5.01 | -log[OH-] |
| pH | 8.99 | 14.00 – pOH |
| Percent hydrolysis | 0.0072% | [OH-] / initial concentration × 100 |
Common mistakes when solving this problem
- Treating sodium hexanoate as neutral. The sodium cation does not control the pH here. The hexanoate anion does.
- Using Ka directly instead of Kb. You need the base hydrolysis constant for the hexanoate ion.
- Forgetting the relationship Ka × Kb = Kw. This is the key bridge between the weak acid and its conjugate base.
- Using 14.00 for every temperature without thought. If temperature changes, Kw changes, and pH + pOH may not equal exactly 14.00.
- Ignoring units and logarithms. pOH and pH are logarithmic quantities and must be handled carefully.
How concentration affects the pH
For weak base salts such as sodium hexanoate, higher concentration generally means a somewhat higher pH, but the relationship is not linear. The hydroxide concentration scales approximately with the square root of concentration:
That means if you quadruple the concentration, hydroxide roughly doubles, not quadruples. This is why pH changes more slowly than many learners expect. For example, dropping the concentration by a factor of 100 does not reduce the pH by 100-fold. Instead, it changes the logarithm of the hydroxide concentration, which moves the pH more moderately.
Practical interpretation of a pH near 9
A pH of about 8.99 indicates a mildly basic solution. This is nowhere near the caustic strength of a strong base like sodium hydroxide, but it is clearly basic enough to matter in analytical methods, extraction chemistry, and buffer-related calculations. Solutions in this range can affect indicator color, extraction efficiency, and the protonation state of weak acids.
When a more advanced model is needed
The standard classroom solution assumes ideal dilute behavior, full salt dissociation, negligible water autoionization compared with hydrolysis, and a single equilibrium. In real laboratory systems, several complications can matter:
- Activity coefficients become important at higher ionic strength.
- Temperature changes alter both pKa and Kw.
- Mixed solvents can dramatically shift acid-base behavior.
- Impurities or dissolved carbon dioxide can perturb measured pH.
- Very dilute solutions may require a stricter treatment of water autoionization.
For ordinary coursework and most standard aqueous exercises, though, the weak-base hydrolysis model is the correct and expected approach.
Recommended reference sources
If you want deeper background on pH, acid-base constants, and laboratory standards, these authoritative references are useful:
- NIST pH standards and measurement guidance
- U.S. EPA overview of pH and aqueous chemistry significance
- University of Washington chemistry resources and instructional materials
Final answer
Using a concentration of 0.137 M sodium hexanoate and a typical pKa of 4.85 for hexanoic acid at 25°C, the calculated solution pH is:
The path to that answer is straightforward: identify hexanoate as a weak base, convert pKa to Ka, convert Ka to Kb, solve for hydroxide concentration, then calculate pOH and pH. If you use the calculator above, you can also explore how the result changes with different assumed pKa values or temperatures.