Metal compounds and qualitative analysis
Qualitative analysis identifies ions from controlled chemical changes—not from color alone.
Learning goals. Write net ionic precipitation equations; use solubility and amphoterism; connect observations to Cu2+, Fe2+, Fe3+, Al3+, and Zn2+; distinguish preliminary and confirmatory tests; record uncertainty and work safely.
1. From ions to precipitates
Many hydrated metal ions react with OH− to form sparingly soluble hydroxides. The general pattern is Mn+(aq) + nOH−(aq) → M(OH)n(s). A net ionic equation omits spectator ions but must conserve atoms and total charge.
| Ion + OH− | Observation | Equation |
|---|---|---|
| Cu2+ | blue precipitate | Cu2+ + 2OH− → Cu(OH)2(s) |
| Fe2+ | pale green precipitate; may brown in air | Fe2+ + 2OH− → Fe(OH)2(s) |
| Fe3+ | red-brown precipitate | Fe3+ + 3OH− → Fe(OH)3(s) |
| Al3+ | white precipitate | Al3+ + 3OH− → Al(OH)3(s) |
| Zn2+ | white precipitate | Zn2+ + 2OH− → Zn(OH)2(s) |
2. Excess reagent reveals amphoterism
Al(OH)3 and Zn(OH)2 react with both acids and strong bases. In excess aqueous OH−, they form soluble hydroxo complexes: Al(OH)3 + OH− → [Al(OH)4]−; Zn(OH)2 + 2OH− → [Zn(OH)4]2−. Their white precipitates dissolve, whereas the listed Cu and Fe hydroxides remain under this test.