Calculate the pH of a 0.575 m Sodium Acetate Solution
Use this premium calculator to estimate the pH of aqueous sodium acetate from hydrolysis chemistry. By default, the calculator is preloaded for a 0.575 m sodium acetate solution at 25 degrees Celsius using the standard acetic acid dissociation constant.
Interactive pH Calculator
Click the button to compute the pH of a 0.575 m sodium acetate solution. The default answer will be slightly basic, because acetate is the conjugate base of acetic acid.
Expert Guide: How to Calculate the pH of a 0.575 m Sodium Acetate Solution
To calculate the pH of a 0.575 m sodium acetate solution, you need to recognize that sodium acetate is not an acid itself. Instead, it is a salt formed from a strong base, sodium hydroxide, and a weak acid, acetic acid. When sodium acetate dissolves in water, the sodium ion has essentially no effect on pH, while the acetate ion behaves as a weak base. That means the solution becomes basic because acetate reacts with water to produce hydroxide ions.
This is a classic weak base hydrolysis problem. The main challenge is deciding which equilibrium constant to use and how to turn the final hydroxide concentration into pH. If you understand conjugate acid base pairs, this entire calculation becomes very systematic. The acetate ion is the conjugate base of acetic acid, so its base dissociation constant, Kb, is related to the acid dissociation constant, Ka, by the equation Kb = Kw / Ka. At 25 degrees Celsius, Kw is 1.0 × 10-14. For acetic acid, a commonly used Ka value is 1.8 × 10-5. Those numbers give acetate a small but meaningful basic character.
Step 1: Write the Dissociation and Hydrolysis Reactions
Sodium acetate dissociates almost completely in water:
Then the acetate ion hydrolyzes water:
This second equation is the one that controls the pH. Because OH– is produced, the pH must end up above 7 at 25 degrees Celsius.
Step 2: Convert Ka to Kb
Most reference data list acetic acid by its acid constant, Ka, rather than listing the base constant for acetate directly. That is not a problem. Use the conjugate relationship:
Substituting the standard values gives:
This Kb is small, which tells you that acetate is a weak base. Even so, at a concentration of 0.575 m, there is enough acetate present to push the solution noticeably into the basic range.
Step 3: Set Up an ICE Table
For the hydrolysis reaction, use an ICE table. If we approximate the 0.575 m value as 0.575 mol per liter for a typical instructional pH calculation, the setup looks like this:
- Initial [CH3COO–] = 0.575
- Initial [CH3COOH] = 0
- Initial [OH–] ≈ 0
Let x be the amount of acetate that reacts:
- Change in [CH3COO–] = -x
- Change in [CH3COOH] = +x
- Change in [OH–] = +x
At equilibrium:
- [CH3COO–] = 0.575 – x
- [CH3COOH] = x
- [OH–] = x
Step 4: Apply the Equilibrium Expression
The base equilibrium expression is:
Substitute the ICE table terms:
Because Kb is very small relative to the concentration, x will be much smaller than 0.575. That allows the common weak base approximation:
So the equilibrium equation simplifies to:
Solving gives:
Step 5: Convert Hydroxide Concentration to pOH and pH
Now take the negative logarithm of hydroxide concentration:
Then calculate pH:
So the pH of a 0.575 m sodium acetate solution is approximately 9.25 under standard classroom assumptions at 25 degrees Celsius.
Why Sodium Acetate Makes Water Basic
Students often wonder why a salt can change pH at all. The answer depends on the parent acid and base. Sodium acetate comes from sodium hydroxide, which is a strong base, and acetic acid, which is a weak acid. Strong base cations such as Na+ generally do not hydrolyze water appreciably, but the conjugate base of a weak acid does. Acetate still has enough proton affinity to pull a proton from water, generating OH–. That is exactly why the pH rises.
By contrast, if you dissolved sodium chloride in water, neither Na+ nor Cl– would hydrolyze in a meaningful way, and the pH would stay close to neutral. The chemistry of sodium acetate is different because acetate retains measurable basic behavior.
Molality Versus Molarity in This Problem
The prompt uses 0.575 m, which technically indicates molality, not molarity. Molality means moles of solute per kilogram of solvent. Molarity means moles of solute per liter of solution. In rigorous physical chemistry, those are different concentration units and should not be interchanged casually. However, in many introductory acid base calculations, especially for moderately dilute aqueous solutions, instructors and textbook problems often treat molality and molarity as approximately similar for practical estimates.
For a more exact treatment, you would need solution density to convert 0.575 m into an actual molarity. Since density is not provided here, the standard instructional approach is to proceed with 0.575 as the effective concentration in the equilibrium expression. That is exactly what the calculator on this page does by default. If you know the exact molarity from experimental data, you can switch the unit and enter it directly.
Approximation Check
Any time you use x ≈ √(KbC), it is smart to verify that the approximation is justified. Here, x was about 1.79 × 10-5, while the initial concentration was 0.575. The ratio is:
That is far below the common 5% guideline, so the approximation is exceptionally good. The quadratic solution and the square root approximation give practically identical pH values for this case.
Comparison Table: pH of Sodium Acetate at Different Concentrations
The table below uses Ka = 1.8 × 10-5 and Kw = 1.0 × 10-14 at 25 degrees Celsius. Values are calculated using the weak base approximation. This helps place 0.575 m in context.
| Sodium acetate concentration | Kb of acetate | Estimated [OH-] | Estimated pOH | Estimated pH |
|---|---|---|---|---|
| 0.010 | 5.56 × 10-10 | 2.36 × 10-6 | 5.63 | 8.37 |
| 0.050 | 5.56 × 10-10 | 5.27 × 10-6 | 5.28 | 8.72 |
| 0.100 | 5.56 × 10-10 | 7.45 × 10-6 | 5.13 | 8.87 |
| 0.575 | 5.56 × 10-10 | 1.79 × 10-5 | 4.75 | 9.25 |
| 1.000 | 5.56 × 10-10 | 2.36 × 10-5 | 4.63 | 9.37 |
Reference Data Table for the Chemistry Used
The next table summarizes the core constants and interpretations that support this calculation. These are standard values commonly used in general chemistry at 25 degrees Celsius.
| Quantity | Typical value | Meaning in this problem |
|---|---|---|
| Ka of acetic acid | 1.8 × 10-5 | Measures how strongly acetic acid donates H+ |
| pKa of acetic acid | 4.74 to 4.76 | Logarithmic form often used in buffer calculations |
| Kw of water | 1.0 × 10-14 | Relates H+ and OH– at 25 degrees Celsius |
| Kb of acetate | 5.56 × 10-10 | Shows acetate is a weak base |
| Neutral pH at 25 degrees Celsius | 7.00 | Benchmark used to classify the sodium acetate solution as basic |
Common Mistakes to Avoid
- Using Ka directly instead of converting to Kb. Because acetate is acting as a base, you need Kb for the hydrolysis equilibrium.
- Forgetting that sodium is a spectator ion. Na+ does not control the pH here.
- Assuming the pH is acidic because the salt contains a known acid fragment. Acetate is the conjugate base of acetic acid, so it raises pH.
- Stopping at pOH. Once [OH–] is known, you still need to calculate pOH and then pH.
- Ignoring the unit issue entirely. 0.575 m is molality, not molarity, so high precision work needs density data for exact conversion.
When Would You Need a More Advanced Model?
The simple calculation is ideal for most classroom and routine lab estimates, but there are cases where a more advanced treatment matters. If the sodium acetate solution is highly concentrated, if temperature differs significantly from 25 degrees Celsius, or if ionic strength effects become important, then activity corrections may be needed. In analytical chemistry, pH measurements can shift slightly from the idealized values predicted by simple concentration-based equations. That does not mean the basic method is wrong. It means it is an approximation built on ideal solution behavior.
For most educational applications, though, the result near pH 9.25 is exactly what you should expect. The approximation is strongly justified, and the chemistry interpretation is sound.
Authoritative Learning Resources
- USGS: pH and Water
- Purdue University: Acid Base Equilibrium Help
- University of Wisconsin: Weak Acids and Weak Bases
Final Answer
If you calculate the pH of a 0.575 m sodium acetate solution using standard general chemistry assumptions at 25 degrees Celsius, with acetic acid Ka = 1.8 × 10-5, the solution has a pH of approximately 9.25. The reason is that acetate acts as a weak base and generates hydroxide ions by hydrolyzing water.