Introduction to Qualitative Analysis
Learning Outcomes
- Apply the safety rules in the chemistry laboratory through proper and safe handling of chemicals and chemical equipment.
- Identify and use common equipment and measuring devices in the chemistry laboratory.
- Properly perform the technique of filtration, quantitative transfer of materials, pipetting and use of the Bunsen burner.
- Identify chemical and physical properties and changes.
- Identify unknown substances through qualitative analysis.
One of the most useful indicators in qualitative analysis is the formation of a precipitate — an insoluble solid that forms when two aqueous solutions react. When a precipitate appears, it signals that a chemical reaction has occurred, specifically a double displacement reaction (also called a double replacement or metathesis reaction). The general form of a double displacement reaction is:
Here, the cations and anions exchange partners. If one of the new combinations forms an insoluble compound, it appears as a solid precipitate, which may have a characteristic color or texture.
Predicting whether a precipitate will form depends on solubility rules, which summarize the tendency of ionic compounds to dissolve in water:
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Always soluble: All salts of Group 1 metals (Li⁺, Na⁺, K⁺, etc.) and ammonium (NH₄⁺) are soluble.
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Always soluble: All nitrates (NO₃⁻), acetates (C₂H₃O₂⁻), and most perchlorates (ClO₄⁻) are soluble.
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Chlorides, bromides, iodides: Soluble except those containing Ag⁺, Pb²⁺, or Hg₂²⁺.
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Sulfates (SO₄²⁻): Soluble except those of Pb²⁺, Ba²⁺, Sr²⁺, and Ca²⁺.
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Carbonates (CO₃²⁻), phosphates (PO₄³⁻), chromates (CrO₄²⁻): Insoluble except when paired with Group 1 metals or NH₄⁺.
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Hydroxides (OH⁻): Insoluble except for Group 1 metals, NH₄⁺, and slightly soluble for Ca²⁺, Sr²⁺, and Ba²⁺.
By applying these rules, you can predict which combinations of ions will produce a precipitate.
Suppose you mix aqueous solutions of silver nitrate (AgNO₃) and sodium chloride (NaCl). The cations (Ag⁺ and Na⁺) will exchange partners with the anions (NO₃⁻ and Cl⁻):
From the solubility rules, most chlorides are soluble, but AgCl is an exception — it is insoluble in water and appears as a white precipitate. Sodium nitrate, on the other hand, is soluble and stays dissolved in the solution. This white precipitate of AgCl provides clear visual evidence that a reaction has occurred and that silver ions were present in the original solution.
In your experiment, you will systematically mix known cation solutions with various anions and look for precipitates, noting their color and texture. Each cation produces a distinct pattern of reactions — a kind of chemical “fingerprint.” By comparing the reaction pattern of your unknown sample to those of the known cations, you will be able to identify the unknown. Careful observation and accurate recording of results are critical for making a correct identification.
Solubility Rules Quick Reference Chart
| Ion or Compound Type | Soluble Compounds (No Precipitate) | Insoluble Compounds (Precipitate Likely) |
|---|---|---|
| Group 1 Cations (Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺) | All compounds soluble | — |
| Ammonium (NH₄⁺) | All compounds soluble | — |
| Nitrates (NO₃⁻) | All nitrates soluble | — |
| Acetates (C₂H₃O₂⁻) | All acetates soluble | — |
| Chlorides (Cl⁻), Bromides (Br⁻), Iodides (I⁻) | Soluble except with Ag⁺, Pb²⁺, Hg₂²⁺ | AgCl, PbCl₂, Hg₂Cl₂ (and similar Br⁻, I⁻ salts) |
| Sulfates (SO₄²⁻) | Soluble except with Pb²⁺, Ba²⁺, Sr²⁺, Ca²⁺ | PbSO₄, BaSO₄, SrSO₄, CaSO₄ |
| Carbonates (CO₃²⁻), Phosphates (PO₄³⁻), Chromates (CrO₄²⁻) | Soluble only with Group 1 or NH₄⁺ | Most others are insoluble |
| Hydroxides (OH⁻) | Soluble with Group 1, NH₄⁺, and slightly soluble with Ca²⁺, Sr²⁺, Ba²⁺ | Most others insoluble |
| Sulfides (S²⁻) | Soluble with Group 1, Group 2, and NH₄⁺ | Most others insoluble |
