Solubility chart: predicting dissolution and precipitation
A solubility chart (often presented as “solubility rules” or a “solubility table”) classifies ionic compounds as soluble (remain dissolved as aqueous ions) or insoluble (form a solid precipitate) in water. The chart is a fast decision tool for precipitation reactions and net ionic equations.
Problem
Use a solubility chart to decide whether a precipitate forms when 50.0 mL of 0.10 M AgNO3(aq) is mixed with 50.0 mL of 0.10 M NaCl(aq). If a precipitate forms, write the net ionic equation.
1) Identify the ions in solution
Silver nitrate dissociates in water:
Sodium chloride dissociates in water:
2) Use the solubility chart to predict the possible precipitate
When two ionic solutions are mixed, new cation–anion pairings are possible: \(\mathrm{Ag^+}\) can pair with \(\mathrm{Cl^-}\), and \(\mathrm{Na^+}\) can pair with \(\mathrm{NO_3^-}\). A solubility chart is then used to check which product (if any) is insoluble.
| Common solubility chart (rules) | How it applies here |
|---|---|
| Nitrates \(\mathrm{(NO_3^-)}\) are soluble. | \(\mathrm{NaNO_3}\) is soluble \(\rightarrow\) remains aqueous (no precipitate from \(\mathrm{Na^+}\) and \(\mathrm{NO_3^-}\)). |
| Alkali-metal salts (Group 1, e.g., \(\mathrm{Na^+}\)) are soluble. | Any \(\mathrm{Na^+}\) salt stays dissolved; \(\mathrm{NaCl}\) and \(\mathrm{NaNO_3}\) are soluble. |
| Most chlorides \(\mathrm{(Cl^-)}\) are soluble, except those of \(\mathrm{Ag^+}\), \(\mathrm{Pb^{2+}}\), \(\mathrm{Hg_2^{2+}}\). | \(\mathrm{AgCl}\) is insoluble \(\rightarrow\) a solid precipitate is predicted. |
Conclusion from the solubility chart: AgCl forms as a precipitate.
3) Confirm by a quick mole check (stoichiometry)
A solubility chart predicts which solid can form. A short stoichiometric check confirms that both needed ions are present in nonzero amounts.
- Moles of \(\mathrm{Ag^+}\) provided by \(\mathrm{AgNO_3}\): \[ n(\mathrm{Ag^+}) = 0.10\,\mathrm{mol\,L^{-1}} \times 0.0500\,\mathrm{L} = 0.00500\,\mathrm{mol} \]
- Moles of \(\mathrm{Cl^-}\) provided by \(\mathrm{NaCl}\): \[ n(\mathrm{Cl^-}) = 0.10\,\mathrm{mol\,L^{-1}} \times 0.0500\,\mathrm{L} = 0.00500\,\mathrm{mol} \]
- Precipitation reaction uses a \(1{:}1\) ratio: \[ \mathrm{Ag^+(aq) + Cl^-(aq) \rightarrow AgCl(s)} \] Equal moles means neither ion is initially limiting; precipitation can proceed until one is consumed.
4) Write the net ionic equation
Spectator ions (ions that remain aqueous and unchanged) are \(\mathrm{Na^+}\) and \(\mathrm{NO_3^-}\). Removing them from the full ionic equation gives the net ionic equation:
5) Optional: predict the appearance of the mixture
Since \(\mathrm{AgCl}\) is insoluble, the mixture turns cloudy as a fine solid forms and can settle over time or be separated by filtration.
Visualization: solubility chart decision flow (mixing two salts)
Final result
A solubility chart predicts that AgCl is insoluble, so a precipitate forms when \(\mathrm{AgNO_3(aq)}\) and \(\mathrm{NaCl(aq)}\) are mixed. The net ionic equation is \(\mathrm{Ag^+(aq) + Cl^-(aq) \rightarrow AgCl(s)}\).