Water solubility for ionic substances reflects a competition between the crystal’s lattice attraction and the stabilizing interactions formed when separated ions become hydrated. Salts that remain as free ions in water behave as electrolytes, while salts with very low dissolution equilibria typically appear as precipitates in aqueous mixtures.
Which of the following would not be water soluble
A room-temperature comparison is assumed (about 25°C) with typical laboratory concentrations where “water soluble” means dissolving to give a clear solution dominated by aqueous ions.
- NaNO3
- K2SO4
- AgCl
- NH4Cl
Answer
AgCl would not be water soluble under ordinary conditions; it is only sparingly soluble and commonly forms a precipitate in water-based reactions.
Assumptions used for the solubility comparison
The solvent is liquid water, temperature is near 25°C, and no strong complexing agents are present (for example, no added NH3 to complex Ag+). “Not water soluble” is interpreted as “does not dissolve appreciably to produce a clear ionic solution.”
Key solubility patterns in water
The following empirically reliable patterns separate most soluble salts from common precipitates in general chemistry:
- Alkali metal salts (Li+, Na+, K+, …): soluble.
- Ammonium salts (NH4+): soluble.
- Nitrates (NO3−): soluble.
- Most chlorides (Cl−): soluble, with notable exceptions including AgCl, PbCl2, and Hg2Cl2.
- Most sulfates (SO42−): soluble, with notable exceptions including BaSO4, PbSO4, and (to a lesser extent) CaSO4.
Evaluation of each option
| Option | Ions in water (if soluble) | Relevant solubility pattern | Expected behavior |
|---|---|---|---|
| NaNO3 | Na+(aq) and NO3−(aq) | Alkali metal salts and nitrates are soluble. | Clear solution; strong electrolyte. |
| K2SO4 | 2 K+(aq) and SO42−(aq) | Alkali metal salts are soluble; most sulfates are soluble. | Clear solution; strong electrolyte (solubility is substantial). |
| AgCl | Only small amounts of Ag+(aq) and Cl−(aq) | Chlorides are usually soluble, with AgCl as a classic exception. | Solid persists; precipitate forms easily. |
| NH4Cl | NH4+(aq) and Cl−(aq) | Ammonium salts are soluble; most chlorides are soluble. | Clear solution; strong electrolyte. |
Equilibrium description for a sparingly soluble salt
The dissolution of silver chloride is an equilibrium:
\[ \mathrm{AgCl(s)\rightleftharpoons Ag^+(aq) + Cl^-(aq)} \]
The solubility product constant formalism captures the low extent of dissolution:
\[ K_{sp} = [\mathrm{Ag^+}]\,[\mathrm{Cl^-}] \]
A small \(K_{sp}\) corresponds to equilibrium concentrations of ions that are too low to dissolve much solid, so AgCl typically remains as a solid phase in water.
Connection to precipitation and net ionic equations
Mixing sources of Ag+(aq) and Cl−(aq) commonly produces a precipitate because the ionic product \([\mathrm{Ag^+}][\mathrm{Cl^-}]\) readily exceeds \(K_{sp}\). A representative net ionic equation is:
\[ \mathrm{Ag^+(aq) + Cl^-(aq)\rightarrow AgCl(s)} \]
Visualization: hydration-dominated dissolution vs lattice-dominated precipitation
Common pitfalls
Chlorides are often remembered as “soluble,” but AgCl is one of the standard exceptions. Sulfates are often remembered as “soluble,” but BaSO4 and PbSO4 are classic insoluble exceptions; confusion between the chloride and sulfate exception lists is a frequent source of errors.