1. General Methods

1. Preparing for Qualitative Analysis
   1. Reactions are generally done in small test tubes with volume of 5 mL or less.
   2. Solutions need to be fully soluble and should be relatively dilute, typically ~0.1 M.
   3. Reagents should be slowly added drop-wise and observed carefully.
   4. Several common "test solutions" are required to establish solubility rules or identify an unknown ion. These contain ions known to react specifically with certain chemical species (cations or anions).
      1. Common solutions include CaNO₃, BaCl₂, (NH₄)₂MoO₄, HCl, AgNO₃, and NaOH, and other solutions as needed.
2. Mixing
   1. Mix solutions by tapping or swirling the test tube in a vertical direction. Use a cork or stopper to prevent splashing the solution.
   2. Remove the cork or stopper, then gently heat the solutions with a water bath or cool flame to induce a reaction. Point the test tube away from any individuals in the lab.
3. Observation and Recovery
   1. Separate the supernatant (non-reacting solution) and precipitate using centrifugation. If more precipitate forms when additional test ion is added, the reaction is incomplete. Continue adding test ion until no more precipitate forms.
   2. Wash the precipitate using centrifugation and pouring or decanting off the supernatant. Add more water and repeat the process for a total of three washings.
   3. Wash large quantities of precipitate by vacuum filtration and recover the dried precipitate from the filter paper.
   4. Note the formation of a precipitate as well as the properties of the precipitate such as color, thickness (gelatinous, cloudy, fine), and crystal formation.
4. Safety and Waste
   1. Always wear safety eyewear while performing qualitative analysis experiments. Gloves may also be necessary based on the reagents used and products formed.
   2. Proper waste disposal methods must be followed closely. Harmful products can be formed when multiple reactants are combined in one container.

2. Anion Analysis

1. Identifying phosphate, carbonate, chloride, and sulfide ions; PO₄^3-, CO₃^2-, Cl^-, S^2-
   1. Phosphate
      1. Add a solution containing phosphate, PO₄^3-, to another solution containing calcium ions, Ca²⁺. The formation of a white precipitate indicates the formation of calcium phosphate, Ca₃(PO₄)₂.
      2. Since many cations form insoluble products with calcium, a more specific reaction is possible. Add H⁺ (acid) to Ca₃(PO₄)₂ to dissolve the solid and form HPO₄^2-. Then combine the HPO₄^2- with ammonium molybdate, (NH₄)₂MoO₄. A positive test yields the yellow precipitate ammonium phosphomolybdate, NH₄)₃PO₄(MoO₃)₁₂(s). The net ionic reactions are as follows:
         3 Ca²⁺(aq) + 2 PO₄^3-(aq) → Ca₃PO₄(s)
         Ca₃PO₄(s) + 2 H⁺(aq) → 3 Ca²⁺ + 2 HPO₄^2-(aq)
         HPO₄^2-(aq) + 12 (NH₄)₂MoO₄(aq) + 23 H⁺(aq) → (NH₄)₃PO₄ (MoO₃)₁₂(s) + 21 NH₄⁺(aq) + 12 H₂O(l) 
   2. Carbonate
      1. Carbonate salts are generally insoluble except in the presence of Group 1 and ammonium cations. Add a few drops of calcium chloride, CaCl₂, to the carbonate-containing solution. In solutions with high carbonate concentrations, a white precipitate forms and indicates the possible formation of calcium phosphate, CaCO₃. The reaction has many interferences, including other anions like phosphate.
         Ca²⁺(aq) + CO₃^2-(aq) → CaCO₃(s) 
      2. Add H⁺ (acid) to a solution containing carbonate, CO₃^2-. The formation of bubbles indicates presence of CO₂, signifying CO₃^2- as a reactant. Carbonate ion behaves as a base in the presence of strong acid to form carbon dioxide gas and water.
         CO₃^2-(aq) + H⁺(aq) → CO₂(g) + H₂O(l)
   3. Chloride
      1. Add silver nitrate to a chloride-containing solution. The formation of a white precipitate indicates the formation of AgCl(s):
         Ag⁺(aq) + Cl^-(aq) → AgCl(s)
   4. Sulfide
      1. Add a copper chloride solution to a solution containing sulfide. The formation of a black precipitate indicates the formation of copper sulfide, CuS. In general, solutions containing sulfide ions, S^2-, react with metal ions to yield an insoluble metal sulfide.
         S^2- + Cu²⁺ → CuS(s). 
         The value of the solubility product, K_sp = 6.3 x 10^-36, indicates the high degree of insolubility of the product.

3. Cation Analysis

1. All alkali metals (group 1) and some alkaline earth metals (group 2) are soluble except under specific conditions.
2. Nearly all Group 3–13 metals are considered insoluble in the presence of sulfide, carbonate, phosphate, and hydroxide. The color and type of precipitate will vary.
   1. Place a chromium solution in a hydroxide solution. A green precipitate will be observed. The general reaction of a +2 metal with a hydroxide is shown below:
      M²⁺ + OH^- → M(OH)₂(s)
   2. It is not possible to differentiate most metal ions based on solubility alone with some notable exceptions:
      1. The addition of silver, Ag⁺, mercury, Hg₂²⁺, or lead, Pb²⁺ to chloride, bromide, or iodide results in precipitate formation.
      2. The addition of strontium, Sr²⁺, barium, Ba²⁺, mercury, Hg₂²⁺, or lead, Pb²⁺ results in a precipitate in the presence of sulfate.
      3. Ba²⁺ forms a yellow solid in the presence of CrO₄^2-, BaCrO₄(s). This is pigment used in oil-based paint commonly known as "barium yellow".
3. Limited insolubility of metal ions allows for other qualitative tests to positively identify each metal. While some form precipitates, others undergo unique color changes in the presence of chelating ions or molecules. Cation identifications include nickel, iron, aluminum, and zinc; Ni²⁺, Fe³⁺, Al³⁺, Zn²⁺.
   1. Add nickel (II) in the presence of dimethylglyoxime (H₂dmg) to form the rose-red precipitate Ni(H₂dmg):
      Ni²⁺(aq) + 2 H₂dmg(aq) → Ni(Hdmg)₂(s) + 2 H⁺(aq) 
   2. Add Iron (III) to thiocyanate ion, SCN^- to form the blood-red [FeNCS]²⁺] complex:
      Fe³⁺(aq) + SCN^-(aq) → [FeNCS]²⁺(aq)
   3. Aluminum ion
      1. Combine aluminum (III) with pyrocatechol violet in a pH 6 ammonium acetate buffer solution to form a blue solution.
      2. Aluminum (III) is also precipitated in the presence of weak base to form the gelatinous-white Al(OH)₃(s) compound. Addition of more base causes the compound to form the clear and colorless [Al(OH)₄]^-(aq) soluble complex.
   4. Zinc ion
      1. Add zinc (II) to a small amount of base to form a white precipitate. Then add more base to re-dissolve the precipitate and form the soluble [Zn(OH)₄]^2- complex.
      2. Add zinc (II) to potassium hexacyanoferrate, K₄[Fe(CN)₆] to form the light green precipitate K₂Zn₃[Fe(CN)₆]₂(s):
         3 Zn²⁺(aq) + 2 K₄[Fe(CN)₆](aq) → K₂Zn₃[Fe(CN)₆]₂(s) + 6 K⁺(aq)