Calculate the Solubility of Zn(OH)2 at pH 9.35
Use this premium calculator to estimate the molar solubility of zinc hydroxide in a buffered solution at a defined pH. The default setup uses a common textbook Ksp value of 3.0 × 10-17 at 25 degrees C and applies the standard common ion approximation: s = Ksp / [OH–]2.
How to Calculate the Solubility of Zn(OH)2 at pH 9.35
To calculate the solubility of zinc hydroxide, Zn(OH)2, at pH 9.35, you start with the equilibrium between solid zinc hydroxide and its dissolved ions. The governing reaction is Zn(OH)2(s) ⇌ Zn2+ + 2OH–. Because the solid phase has constant activity, it does not appear in the equilibrium expression. The solubility product becomes Ksp = [Zn2+][OH–]2. If the pH is fixed by a buffer or by the surrounding solution chemistry, then the hydroxide concentration is known, and the dissolved zinc concentration can be solved directly.
At pH 9.35, the pOH is 14.00 – 9.35 = 4.65, assuming standard 25 degrees C acid-base relations. The hydroxide concentration is therefore [OH–] = 10-4.65 ≈ 2.24 × 10-5 M. If you adopt a representative Ksp value of 3.0 × 10-17 for Zn(OH)2, then the molar solubility is:
That means the dissolved zinc concentration, under the assumptions of the simple Ksp model, is about 5.98 × 10-8 M. Converting that to mass concentration gives approximately 0.00594 mg/L as Zn(OH)2 or about 0.00391 mg/L as Zn2+. This is a very low concentration, and it illustrates a key principle in solubility equilibria: when hydroxide is already present in solution, the common ion effect suppresses the dissolution of metal hydroxides.
Why pH matters so much
Metal hydroxides are especially sensitive to pH because hydroxide appears directly in the Ksp expression. For Zn(OH)2, the hydroxide term is squared. That means even modest changes in pH create substantial shifts in calculated solubility. When pH rises by 1 unit, [OH–] increases by a factor of 10, and for a hydroxide with a squared hydroxide term, the predicted solubility can drop by a factor of about 100, if complexation is ignored. This is why zinc removal by precipitation often depends strongly on pH control.
At lower pH, hydroxide concentration falls, so the equilibrium can tolerate more dissolved Zn2+. At higher pH, the common ion effect pushes equilibrium toward the solid phase, reducing dissolved zinc. However, zinc chemistry is not perfectly monotonic across all conditions, because zinc is amphoteric. In strongly basic media, species such as zincate complexes can form, increasing total dissolved zinc again. That is one reason why real treatment systems often rely on both thermodynamic modeling and plant data.
Step by step method for pH 9.35
- Write the dissolution equilibrium: Zn(OH)2(s) ⇌ Zn2+ + 2OH–.
- Write the solubility product expression: Ksp = [Zn2+][OH–]2.
- Convert pH to pOH using pOH = 14 – pH.
- Calculate hydroxide concentration: [OH–] = 10-pOH.
- Solve for dissolved zinc concentration: [Zn2+] = Ksp / [OH–]2.
- Convert molar concentration to mg/L if needed.
pOH = 4.65
[OH–] = 2.24 × 10-5 M
Solubility = 5.98 × 10-8 M
Solubility = 0.00594 mg/L as Zn(OH)2
Solubility = 0.00391 mg/L as Zn2+
Important Assumptions Behind the Calculator
This type of calculation is widely taught in general chemistry, analytical chemistry, and environmental engineering because it captures the first order dependence of solubility on pH. Still, it depends on several assumptions. First, the pH is treated as fixed, which is appropriate for a buffered system or a large body of water where dissolution of a small amount of solid does not significantly change pH. Second, activity effects are ignored, meaning concentrations are used in place of activities. Third, the calculation assumes that zinc exists primarily as free Zn2+ rather than as hydroxo-complexes, carbonate complexes, or organic ligand complexes. In very dilute, moderately basic systems, that can be a useful approximation. In more complex waters, it becomes less exact.
For Zn(OH)2, another important issue is amphoterism. Zinc hydroxide can dissolve in strongly acidic and strongly basic conditions. At higher pH, complex ions such as Zn(OH)3– and Zn(OH)42- may become significant. The simple Ksp treatment only tracks free Zn2+ coupled to OH– and therefore tends to underestimate total dissolved zinc in conditions where complexation matters. That is why sophisticated water chemistry software, such as geochemical speciation models, can predict a U shaped solubility profile rather than a continuously decreasing one.
When the simple answer is usually acceptable
- Educational problems that specifically request a Ksp based solution.
- Buffered laboratory systems near neutral to moderately basic pH.
- Preliminary engineering screening calculations.
- Quick comparison of how pH shifts affect precipitation tendency.
When you need a more advanced model
- High ionic strength process streams.
- Solutions containing carbonate, phosphate, ammonia, citrate, EDTA, or strong organic ligands.
- Strongly alkaline solutions where zincate formation is important.
- Regulatory or design work requiring defensible site-specific predictions.
Comparison Data: Predicted Zn(OH)2 Solubility Versus pH
The table below uses Ksp = 3.0 × 10-17 and the simplified equation s = Ksp / [OH–]2. These are calculated values, not direct measurements, but they represent the standard equilibrium estimate under the fixed pH assumption.
| pH | pOH | [OH–] (M) | Predicted Solubility (M) | Predicted Solubility (mg/L as Zn2+) |
|---|---|---|---|---|
| 8.00 | 6.00 | 1.00 × 10^-6 | 3.00 × 10^-5 | 1.961 |
| 8.50 | 5.50 | 3.16 × 10^-6 | 3.00 × 10^-6 | 0.196 |
| 9.00 | 5.00 | 1.00 × 10^-5 | 3.00 × 10^-7 | 0.0196 |
| 9.35 | 4.65 | 2.24 × 10^-5 | 5.98 × 10^-8 | 0.00391 |
| 10.00 | 4.00 | 1.00 × 10^-4 | 3.00 × 10^-9 | 0.000196 |
| 11.00 | 3.00 | 1.00 × 10^-3 | 3.00 × 10^-11 | 0.00000196 |
Notice the pattern: from pH 8.00 to pH 9.00, predicted solubility drops by two orders of magnitude. By pH 10.00, the simple Ksp estimate becomes extremely small. This is a hallmark of metal hydroxide precipitation systems. The calculation for pH 9.35 falls neatly within this trend and shows why even a modestly alkaline pH can strongly suppress dissolved free zinc.
Mass-Based Interpretation of the pH 9.35 Result
Many students and practitioners prefer mg/L because it aligns with water quality reporting. The next table converts the pH 9.35 molar solubility result using the molar masses of Zn(OH)2 and Zn. Zinc hydroxide has a molar mass of about 99.38 g/mol, while elemental zinc contributes 65.38 g/mol per mole of dissolved zinc species. These conversions are useful depending on whether you are discussing the precipitate formula or the dissolved metal itself.
| Basis | Molar Mass Used | Concentration at pH 9.35 | Interpretation |
|---|---|---|---|
| As dissolved Zn(OH)2 equivalent | 99.38 g/mol | 0.00594 mg/L | Mass concentration expressed as the hydroxide formula unit |
| As dissolved Zn2+ | 65.38 g/mol | 0.00391 mg/L | Mass concentration expressed as zinc only |
| Molar concentration | Not applicable | 5.98 × 10^-8 mol/L | Direct equilibrium result from Ksp |
Practical Meaning in Water Treatment and Environmental Chemistry
The ability to calculate the solubility of Zn(OH)2 at pH 9.35 matters in several real-world settings. In industrial wastewater treatment, zinc is often removed by raising pH so that zinc hydroxide precipitates. In environmental monitoring, dissolved zinc levels depend strongly on water chemistry, especially pH and the presence of ligands. In electroplating, mining, and metal finishing processes, operators may use pH adjustment as part of compliance strategies. The calculation you see on this page provides a fast estimate of the dissolved free zinc level expected if equilibrium with Zn(OH)2 controls the system.
At the same time, field systems often show departures from the simple prediction. Carbonate alkalinity can generate zinc carbonate or mixed precipitates. Organic ligands can keep zinc in solution even when pH is favorable for hydroxide precipitation. Suspended solids, kinetic limitations, redissolution, and measurement artifacts can also affect observed concentrations. So while the pH 9.35 equilibrium estimate is chemically meaningful, it should be interpreted as a thermodynamic baseline, not an automatic guarantee of measured plant performance.
Key takeaways
- At pH 9.35, Zn(OH)2 is predicted to be only sparingly soluble under a simple Ksp model.
- Using Ksp = 3.0 × 10-17, the solubility is about 5.98 × 10-8 M.
- This equals approximately 0.00594 mg/L as Zn(OH)2 or 0.00391 mg/L as Zn2+.
- Increasing pH generally decreases free Zn2+ until amphoteric dissolution and complexation become important.
- For rigorous design, use activity corrections and speciation modeling when needed.
Authoritative Reference Links
For deeper study of zinc chemistry, pH effects, and aqueous environmental behavior, review these authoritative resources:
- PubChem, Zinc Hydroxide, U.S. National Library of Medicine
- U.S. EPA, pH and acidification overview
- U.S. Geological Survey, water resources and chemistry context
Final Answer Summary
If you need the short answer to the question, “calculate the solubility of Zn(OH)2 at pH 9.35,” a standard Ksp calculation with Ksp = 3.0 × 10-17 gives a molar solubility of about 5.98 × 10-8 M. That corresponds to roughly 0.00594 mg/L as Zn(OH)2 or 0.00391 mg/L as Zn2+. This result assumes a fixed pH, ideal behavior, and no significant zinc complexation beyond the basic solubility product model.