Calculate the pH of a Saturated Solution of CaOH
This calculator is designed for the common chemistry case usually meant by “CaOH,” namely saturated calcium hydroxide, Ca(OH)2. Enter a solubility product value, choose a preset, and instantly compute molar solubility, hydroxide concentration, pOH, and pH.
Saturated Calcium Hydroxide pH Calculator
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Enter or confirm your values, then click Calculate pH.
Expert Guide: How to Calculate the pH of a Saturated Solution of CaOH
When students, lab workers, or search users ask how to “calculate the pH of a saturated solution of CaOH,” they almost always mean calcium hydroxide, whose correct chemical formula is Ca(OH)2. The shorthand “CaOH” is common in casual searches, but calcium hydroxide is the chemically meaningful compound used in pH and solubility calculations. In water, calcium hydroxide behaves as a sparingly soluble ionic base. That means it does not dissolve infinitely, but the amount that does dissolve releases hydroxide ions and makes the solution strongly basic.
The key idea is that a saturated solution contains the maximum amount of dissolved Ca(OH)2 that can exist in equilibrium with undissolved solid at a specified temperature. Because the dissolved ions are limited by equilibrium, the pH is not found by assuming an arbitrary concentration. Instead, you calculate the concentration from the solubility product constant, or Ksp.
Why the pH is high even though calcium hydroxide is only sparingly soluble
Calcium hydroxide is not as soluble as sodium hydroxide or potassium hydroxide, but each formula unit that dissolves produces two hydroxide ions. Those hydroxide ions strongly increase basicity. As a result, a saturated solution of Ca(OH)2 typically has a pH around 12.3 to 12.4 at room temperature, depending on the Ksp value and assumptions used.
- It is a strong base with respect to the dissolved portion.
- Its total dissolved amount is limited by solubility equilibrium.
- Each mole of dissolved Ca(OH)2 produces 2 moles of OH–.
- That stoichiometric factor is why the pH is noticeably higher than many learners first expect.
The equilibrium you must use
For a saturated solution, begin with the dissolution equation:
Ca(OH)2(s) ⇌ Ca2+(aq) + 2OH–(aq)
The solubility product expression is:
Ksp = [Ca2+][OH–]2
If the molar solubility is represented by s, then:
- [Ca2+] = s
- [OH–] = 2s
Substitute these into the Ksp expression:
Ksp = s(2s)2 = 4s3
From this, solve for molar solubility:
s = (Ksp/4)1/3
Then compute hydroxide concentration:
[OH–] = 2s
Next, find pOH:
pOH = -log[OH–]
Finally, convert to pH:
pH = 14.00 – pOH at 25 degrees C, or more generally pH = pKw – pOH.
Step-by-step example at 25 degrees C
Suppose the Ksp of calcium hydroxide at 25 degrees C is 5.02 × 10-6. Here is the complete workflow.
- Write the Ksp equation: Ksp = 4s3
- Insert Ksp: 5.02 × 10-6 = 4s3
- Solve for s: s = (5.02 × 10-6 / 4)1/3
- Find s: s ≈ 0.01079 M
- Find hydroxide concentration: [OH–] = 2s ≈ 0.02158 M
- Find pOH: pOH = -log(0.02158) ≈ 1.666
- Find pH: pH = 14.000 – 1.666 ≈ 12.334
So the pH of a saturated calcium hydroxide solution is approximately 12.33 under these assumptions.
Common mistakes when calculating saturated Ca(OH)2 pH
Even strong students can make a few predictable errors with this problem. The most common issues are not algebraic, but conceptual.
- Using CaOH instead of Ca(OH)2: The actual compound is calcium hydroxide, not “CaOH” as a valid neutral formula.
- Forgetting the 2 in 2OH–: This is the biggest source of incorrect answers.
- Setting Ksp = s3 instead of 4s3: You must square the hydroxide term properly.
- Confusing pOH and pH: For basic solutions, calculate pOH first, then convert.
- Using 14 regardless of temperature: In more advanced work, use the appropriate pKw for the temperature given.
Comparison table: effect of Ksp on saturated solution pH
The exact pH changes a little depending on the Ksp value selected from different references or temperatures. The table below uses the same equilibrium model and shows how pH varies as Ksp changes.
| Ksp | Molar Solubility, s (M) | [OH–] (M) | pOH | pH at pKw = 14.00 |
|---|---|---|---|---|
| 4.68 × 10-6 | 0.01053 | 0.02106 | 1.677 | 12.323 |
| 5.02 × 10-6 | 0.01079 | 0.02158 | 1.666 | 12.334 |
| 5.50 × 10-6 | 0.01112 | 0.02224 | 1.653 | 12.347 |
| 6.00 × 10-6 | 0.01145 | 0.02289 | 1.640 | 12.360 |
How saturated calcium hydroxide compares with other bases
Another good way to understand the result is to compare saturated calcium hydroxide with common strong bases. Sodium hydroxide and potassium hydroxide dissolve readily to make very high hydroxide concentrations. Calcium hydroxide is less soluble, so while it is still strongly basic, the saturation point keeps its pH below the pH of a concentrated NaOH solution.
| Base | Representative Aqueous Case | Approximate [OH–] | Approximate pH | Comment |
|---|---|---|---|---|
| Ca(OH)2 | Saturated at 25 degrees C | 0.0216 M | 12.33 | Limited by Ksp equilibrium |
| NaOH | 0.10 M solution | 0.10 M | 13.00 | Highly soluble strong base |
| KOH | 0.10 M solution | 0.10 M | 13.00 | Highly soluble strong base |
| NH3 | 0.10 M solution | Much lower than 0.10 M | About 11.1 | Weak base equilibrium |
When you should use Ksp instead of direct concentration
Use a Ksp-based approach whenever the problem states that the solution is saturated, contains undissolved solid, or asks for equilibrium concentration in contact with the solid phase. If the problem instead gives you a known dissolved concentration of Ca(OH)2 and says the solution is fully dissolved and not limited by saturation, then you can often work directly from concentration and stoichiometry.
For example:
- Saturated solution: Use Ksp.
- 0.0050 M Ca(OH)2 prepared and fully dissolved: [OH–] = 0.0100 M directly.
- Buffer or mixed-ion system: Consider common ion effects and full equilibrium treatment.
How the common ion effect changes the answer
If calcium ions or hydroxide ions are already present from another source, the solubility of Ca(OH)2 decreases. This is the common ion effect. A saturated solution in pure water has one pH, but a saturated solution in a solution already containing CaCl2 or NaOH will behave differently. In those cases, you do not simply use s and 2s as the full equilibrium concentrations. Instead, you include the initial ion concentration in the equilibrium expression and solve accordingly.
Example concept
If a solution already contains OH–, then adding Ca(OH)2 will not dissolve as much before equilibrium is reached. This lowers the added solubility contribution and can shift the resulting pH from the pure-water saturated value.
Laboratory relevance of saturated calcium hydroxide
Saturated calcium hydroxide solution is often called limewater. It has important uses in chemistry teaching, environmental testing, and industrial processes. One famous application is the qualitative test for carbon dioxide. When CO2 is bubbled through limewater, the solution becomes cloudy due to calcium carbonate precipitation:
Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)
This simple reaction links acid-base chemistry, gas analysis, and precipitation equilibrium in one classic experiment.
Authoritative references for chemistry data and equilibrium concepts
For further reading, consult authoritative educational and government resources such as: LibreTexts Chemistry, U.S. Environmental Protection Agency, NIST Chemistry WebBook, UC Berkeley Chemistry
Final takeaway
To calculate the pH of a saturated solution of “CaOH,” interpret the compound correctly as Ca(OH)2, use the dissolution equilibrium for calcium hydroxide, solve for molar solubility from Ksp = 4s3, determine hydroxide concentration as 2s, and then convert from pOH to pH. At room temperature with a representative Ksp around 5.02 × 10-6, the pH is approximately 12.33. That result is high because even limited dissolution produces a meaningful hydroxide concentration. If your instructor or source provides a different Ksp or temperature, use those values and repeat the exact same method.
This calculator automates that workflow while still showing the chemistry behind the answer, making it useful for homework checks, lab preparation, and quick equilibrium analysis.