Wired Chemist

Background

Unlike covalent compounds, which can be identified using physical properties like boiling point and refractive index, ionic compounds are more appropriately identified with their chemical properties. In the qualitative analysis procedure, the chemical properties of an unknown substance are determined by systematically reacting the unknown with a number of different reagents. By predetermining what the particular reaction will produce if a specific ion is present, the ions that actually are in the solution can be identified. For example, if a reaction is known to produce a precipitate if ion A is present and a precipitate is formed when the reaction is run, then ion A may be present in solution (there may be, and usually are, other ions that will also precipitate with a particular reagent). If no precipitate is formed when the reaction is run, then ion A is clearly not present in the unknown solution and a different reaction will have to be run to determine what ions are present.

There are two general situations in which qualitative analysis is used - in the identification of a simple salt, or the identification of multiple cations in a solution.

Identifying a Simple Salt

The basic testing procedure for identifying a salt is as follows.

1. Appearance of compound

   The compound will most likely be in solid form. Note the color and shape of the crystals. Ionic compounds formed from the representative elements tend to be white or colorless, while ions of transition elements tend to be colored. The following is a table of the colors of metal ions in solution with NO₃^-.

   Ion	Color
   Co2+	rose
   Co3+	violet
   Cr3+	violet
   Cu2+	blue
   Fe2+	pale green, pale violet
   Fe3+	yellow-brown
   Mn2+	pale pink
   Ni2+	blue-green

2. Heating effect

   Heating a compound can cause a liquid to condense on the inside of the test tube. This is probably water, indicating that the compound is a hydrate. If a gas is given off, note the color and odor of the gas. The nitrate, carbonate, and sulfite ions may decompose, as illustrated by the reactions:

   2 Pb(NO₃)₂(s) + heat → 2 PbO(s) + O₂(g) + 4 NO₂(g, brown)
   CaCO₃(s) + heat → CaO(s) + CO₂(g, colorless, odorless)
   CaSO₃(s) + heat → CaO(s) + SO₂(g, colorless, pungent)

   Some bromides and iodides decompose to give Br₂(g, orange-brown) and I₂(g, purple).

3. Flame test

   Solutions of ions, when mixed with concentrated HCl and heated on a nickel/chromium wire in a flame, cause the flame to change to a color characteristic of the atom.Visible colors occur with the following ions:

   Sodium	Bright yellow (intense, peristent)
   Potassium	Pale violet (slight, fleeting)
   Calcium	Brick red (medium, fleeting)
   Strontium	Crimson (medium)
   Barium	Light green (slight)
   Lead	Pale bluish (slight, fleeting)
   Copper	Green or blue (medium, persistent)

4. Solubility in water

   Place one small spatula of the compound in 1 mL of water. If the compound is soluble this amount will dissolve after considerable stirring. If the compound is moderately soluble, some of this amount will dissolve. If the compound is insoluble, even a very small amount will not dissolve.

   General solubility rules:

   • All nitrates are soluble.
   • Practically all sodium, potassium, and ammonium salts are soluble.
   • All chlorides, bromides, and iodides are soluble except those of silver, mercury(I), and lead(II).
   • All sulfates are soluble except those of strontium, barium, and lead(II), which are insoluble, and those of calcium and silver which are moderately soluble.
   • All carbonates, sulfites, and phosphates are insoluble except those of sodium, potassium, and ammonium.
   • All sulfides are insoluble except those of the alkali metals, the alkaline earth metals, and ammonium.
   • All hydroxides are insoluble except those of the alkali metals. The hydroxides of calcium, strontium, and barium are moderately soluble. Ammonium hydroxide does not exist; ammonium hydroxide is a misnomer for aqueous ammonia, NH₃(aq).
5. Reaction with nitric acid

   Add nitric acid to the compound and observe any reaction that occurs. If the compound dissolved in water, it should dissolve in nitric acid. If it did not dissolve in water, but appears to be dissolving in nitric acid, it is undergoing a chemical reaction. In general, compounds that contain anions that are the conjugate bases of weak acids will react (unless the compounds are very insoluble). For example:

   CaCO₃(s) + 2 H⁺(aq) → Ca₂⁺(aq) + H₂O(l) + CO₂(g, colorless, odorless)
   NiS(s) + 2 H⁺(aq) → Ni₂⁺(aq) + H₂S(g, colorless, rotten egg smell)
   Ca₃(PO₄)₂(s) + 6 H⁺(aq) → 3 Ca₂⁺(aq) + 2 H₃PO₄(aq)

   The remaining tests must be perfomed on a solution of the compound.
   If the compound is insoluble in water, dissolve it in nitric acid. Otherwise, dissolve in water.

6. Reaction with sodium hydroxide

   Add NaOH dropwise to the solution, stir or shake the solution, and observe any reaction (if the compound was dissolved in nitric acid, the first several drops will neutralize the acid so be sure to check the pH with litmus paper). Look for a precipitate (refer to the solubility rules for hydroxides). If a precipitate forms, continue adding NaOH. Some metal hydroxides are amphoteric and will form a complex ion and redissolve. See Figures 1, 2, and 3 for an example of this reaction. The following ions are amphoteric:

   Species	Acidic Solution	Slightly Basic Solution	Basic Solution
   Al3+	Al3+(aq)	Al(OH)3(s)	Al(OH)4-(aq)
   Cr3+	Cr3+(aq)	Cr(OH)3(s)	Cr(OH)4-(aq)
   Pb2+	Pb2+(aq)	Pb(OH)2(s)	Pb(OH)42-(aq)
   Zn2+	Zn2+(aq)	Zn(OH)2(s)	Zn(OH)42-(aq)
   Sn4+	Sn4+(aq)	Sn(OH)4(s)	Sn(OH)62-(aq)

   

   Figure 1. Al(NO₃)₃ in solution.

   

   Figure 2. Al(OH)₃ precipitates with the addition of NaOH.

   

   Figure 3. When excess NaOH is added, the precipitate redissolves as the Al(OH)₄^- complex ion is formed.

7. Reaction with ammonia

   Add NH₃ dropwise to the solution, stir or shake the solution, and observe any reaction. If a metal hydroxide precipitate forms, continue adding ammonia. Some metal hydroxides fom a complex ion and redissolve. See Figures 4, 5, and 6 for an example of this reaction. The following ions form ammonia complexes:

   Acid Solution	Basic Solution	Solution with Excess NH3	Color of Complex
   Ni2+(aq)	Ni(OH)2(s)	Ni(NH3)62+(aq)	violet
   Cu2+(aq)	Cu(OH)2(s)	Cu(NH3)42+(aq)	blue
   Zn2+(aq)	Zn(OH)2(s)	Zn(NH3)62+(aq)	colorless
   Ag+(aq)	Ag2O(s)	Ag(NH3)2+(aq)	colorless
   Cd2+(aq)	Cd(OH)2(s)	Cd(NH3)42+(aq)	colorless

   

   Figure 4. Cu(NO₃)₂ in solution.

   

   Figure 5. Cu(OH)₂ precipitates with the addition of NH₃.

   

   Figure 6. With excess NH₃, copper forms the Cu(NH₃)₄^2- complex ion.

8. Reaction with hydrochloric acid

   Add HCl dropwise until solution tests acidic to litmus paper and observe any reaction. A precipitate will form with any cation that forms an insoluble chloride (refer to the solubility rules). For example:

   Pb²⁺ + 2Cl- → PbCl₂(s)

9. Reaction with sulfuric acid

   Add H₂SO₄ dropwise until solution is acidic and observe any reaction. A precipitate will form with any cation that forms an insoluble sulfate (refer to the solubility rules). For example:

   Ba²⁺ + SO₄^2- → BaSO₄(s)

10. Reaction with silver nitrate

    Add HNO₃ dropwise until solution is acidic (unless of course it was dissolved in nitric acid), then add a few drops of AgNO₃ and observe any reaction. A precipitate will form with certain cations that form insoluble silver compounds, but because of the acidic environment, some insoluble silver salts (e.g. salts containing CO₃^2-, S^2-, and PO₄^3- ions) are "destroyed." Cl^-, Br^-, and I^- form insoluble compounds, while SO₄^2- forms a moderately insoluble compound.

    Ag⁺ + Cl^- → AgCl(s)

11. Reaction with barium nitrate

    Add HNO₃ dropwise until solution is acidic, boil the solution for two minutes, then test with litmus paper. Continue adding and boiling until solution remains acidic after boiling. Cool the solution and add a few drops of Ba(NO₃)₂ and observe any reaction. A precipitate will form with anions that form an insoluble barium compound (except the ones destroyed by acid as in the above test).

12. Specific Tests

    Sometimes the above tests can not definitively confirm the presence of a specific ion. In these cases, it is necessary to do specific tests for a particular ion.

    Example Unknown Salts

    Sample 1 had the following characteristics:

    • Visual test: white crystalline powder
    • Heat test: brown gas given off
    • Flame test: no color identified
    • Solubility in water: soluble
    • Nitric acid: soluble
    • Hydroxide: formed an insoluble white precipitate, then dissolved with excess NaOH
    • Ammonia: formed an insoluble white precipitate which did not redissolve in excess NH₃
    • Hydrochloric acid: formed an insoluble white precipitate

      Analysis of observations:

      • The brown gas given off during the heat test indicates presence of the NO₃^- ion, since the NO₃^- ion reacts to form brown NO₂ gas as shown below:

        2 Pb(NO₃)₂(s) + heat → 2 PbO(s) + O₂(g) + 4 NO₂(g)

      • The white precipitate which formed and then redissolved with the addition of sodium hydroxide indicates the presence of an amphoteric cation. The possibilities are Al, Cr, Pb, Zn, and Sn. The reaction occurs as follows:

        Cr³⁺(aq) + 3 OH^-(aq) → Cr(OH)₃(s)
        Cr(OH)₃(s) + OH^-(aq) → Cr(OH)₄^-(aq)

      • Since the cation does not form an ammonia complex, it eliminates Zn from the list of possible cations established above.
      • There are only three cations which form precipitates with hydrochloric acid, Pb²⁺, Ag⁺, and Hg₂²⁺. Of these three, only Pb²⁺ is amphoteric. The reaction with chloride occurs as follows:

        Pb₂⁺ + 2 Cl- → PbCl₂(s)

    Conclusion: Sample 1 is Pb(NO₃)₂

    Sample 2 had the following characteristics:

    • Visual test: white crystalline powder
    • Heat test: no reaction
    • Flame test: green color observed
    • Specific barium test (precipitation with CrO₄^2-): formed an insoluble yellow precipitate
    • Solubility in water: insoluble
    • Nitric acid: produced odorless, colorless gas

      Analysis of observations:

      • The green color observed during the flame test indicates the presence of barium. Although the remaining tests could be done to confirm the presence of barium, none is specific for just the Ba²⁺ ion. Therefore, the specific barium test is used.
      • The yellow precipitate produced by the barium test confirms the presence of barium. The reaction occurs as follows:

        Ba²⁺(aq) + CrO₄^2-(aq) → BaCrO₄(s)

      • The odorless, colorless gas produced by the addition of nitric acid is CO₂.
        This indicates the presence of the CO₃^2- ion.

    Conclusion: Sample 2 is BaCO₃

    Sample 3 had the following characteristics:

    • Visual test: yellow/brown solid
    • Heat test: no reaction
    • Flame test: no reaction
    • Solubility in water: soluble
    • Nitric acid: soluble
    • Sodium hydroxide: formed an insoluble white precipitate which did not redissolve in excess NaOH
    • Ammonia: formed an insoluble white precipitate which did not redissolve in excess NH₃
    • Hydrochloric acid: no reaction
    • Specific iron(III) test (formation of a thiocyanate complex): solution turned blood red
    • Silver nitrate: formed an insoluble white precipitate
    • Specific chloride, bromide, iodide test: precipitate dissolved with addition of 6 M ammonia

      Analysis of observations:

      • The color of the solid is characteristic of the Fe³⁺ ion.
      • Since the compound forms an insoluble precipitate with hydroxide, it eliminates the alkali metals, calcium, strontium, and barium. Since the cation is not amphoteric, it also eliminates aluminum, chromium, lead, zinc, and tin.
      • The cation does not form an ammonia complex, which eliminates nickel, copper(II), silver, and cadmium.
      • Since there is no reaction with Cl^-, Hg₂²⁺ is also eliminated. Of the ions which have not been eliminated (Fe³⁺, Mn²⁺, Bi²⁺), Fe³⁺ is the most likely. A specific ion test is needed to confirm the presence of Fe³⁺.
      • The red solution produced by the iron(III) test confirms the presence of iron(III). The reaction occurs as follows:

        Fe³⁺(aq) + SCN^-(aq) → Fe(SCN)²⁺(aq)

      • The precipitate formed by the addition of silver nitrate indicates the presence of either chloride, bromide, or iodide. Although the colors are different (AgCl white, AgBr cream, AgI yellow), they are difficult to distinguish and a specific ion test is needed to determine which one is present.
      • The specific ion test indicated that the anion present is Cl^-. The reaction occurs as follows:

        Ag⁺(aq) + Cl^-(aq) → AgCl(s)
        AgCl(s) + 2 NH₃(aq) → Ag(NH₃)₂⁺(aq) + Cl^-(aq)

    Conclusion: Sample 3 is FeCl₃

    Specific Tests

    1. Cation Identification Tests
       1. Generally Soluble Cations
          • Ammonium Ions:

            Take a small amount of the material to be tested and place it in a 50-mL beaker. Add 6 M NaOH and smell cautiously. The odor of ammonia indicates the presence of ammonium ions. If you do not smell ammonia, warm the beaker and again smell the emitted vapors. The liberated ammonia will also change the color of a moistened strip of red litmus paper held at the entrance of the test tube.

            NH₄⁺(aq) + OH^-(aq) → NH₃(g) + H₂O

            This test is very reliable. It should be performed whenever the generally soluble cations, NH₄⁺, Na⁺, and K⁺, are suspected.

          • Sodium Ions:

            The most common method of identification of Na⁺ is the flame test. Sodium imparts a brilliant, long lasting, yellow flame that masks colors from other ions. The test may be performed on a small sample of the unknown treated with concentrated HCl or a few drops of solution unknown treated with concentrated HCl. The flame should be bright and it should last as long as that of 0.1 M NaCl. Sodium is a common impurity and traces will be found in almost any unknown. You must learn to distinguish between an unknown that has sodium ion as the cation and an unknown that has sodium ion as an impurity.

          • Potassium Ions:

            The most common method of identification of K⁺ is the flame test. The test may be performed on a small sample of the unknown treated with concentrated HCl or a few drops of solution unknown treated with concentrated HCl. The violet flame is not intense but it is clearly visible in the absence of sodium ions. Cobalt glass filters yellow light from sodium impurities and allows the violet flame to be seen. Do not confuse the glowing wire for the potassium flame.

       2. Cations That Form Insoluble Chlorides
          • Silver Ions:

            Although Ag⁺, Pb²⁺, and Hg₂²⁺ all form insoluble white chlorides, Ag⁺ is the only one of these cations that forms an ammonia complex. Therefore, AgCl dissolves readily in aqueous NH₃. When the resulting solution is acidified with HNO₃, AgCl reprecipitates.

            AgCl(s) + 2NH₃(aq) → Ag(NH₃)²⁺(aq) + Cl^-(aq)
            Ag(NH₃)²⁺(aq) + Cl^-(aq) + 2H₃O⁺ → AgCl(s) + 2NH₄⁺ + 2H₂O

            Add 3 M HCl dropwise to the solution being tested. If a white precipitate is formed, centrifuge and remove the supernatant liquid. Add 6 M NH₃ solution to the precipitate. If the precipitate dissolves, add 6 M HNO₃. Formation of a white precipitate indicates Ag⁺.

          • Lead Ions:

            Although PbCl₂ is insoluble at room temperature, its solubility is increased dramatically at higher temperatures; it dissolves readily in boiling water. Pb²⁺ also forms an insoluble white sulfate, which dissolves in a solution containing acetate ion due to the formation of the weak electrolyte, Pb(CH₃COO)₂. The addition of chromate ion to this lead acetate solution yields a precipitate of yellow lead chromate.

            Pb²⁺(aq) + SO₄^2-(aq) → PbSO₄(s)
            PbSO₄(s) + 2CH₃COO^-(aq) → Pb(CH₃COO)₂(aq) + SO₄^2-(aq)
            Pb(CH₃COO)₂(aq) + CrO₄^2-(aq) → PbCrO₄(s) + 2CH₃COO^-(aq)

            To the solution to be tested add 3 M HCl dropwise. (A large excess of HCl must be avoided because of the formation of the soluble chloro complex, PbCl₄^2-.) Centrifuge and remove the supernatant from the white precipitate (PbCl₂). Add hot water to the precipitate and stir. If the precipitate dissolves, Pb²⁺ is indicated. Add 3 M H₂SO₄ to the hot solution. Centrifuge and remove the supernatant liquid from the white precipitate (PbSO₄). To the precipitate add 3 M NH₄(CH₃COO) and stir. If the white precipitate was PbSO₄, it will dissolve. To confirm, add a few drops of 0.5 M K₂CrO₄ to the resulting solution. A yellow precipitate of PbCrO₄ indicates the presence of Pb²⁺.

          • Mercury(I) Ions:

            When Hg₂Cl₂ is treated with aqueous NH₃ a reaction occurs in which free mercury and amidochloromercury(II) are formed.

            Hg₂Cl₂(s) + NH₃(aq) → Hg(l) + HgNH₂Cl(s) + HCl(aq)

            The HgNH₂Cl is a white solid, while the Hg in a finely divided state appears black. The resultant mixture is gray to black.

            Add 3 M HCl to the solution to be tested for Hg₂²⁺. If a white precipitate forms, centrifuge and remove the supernatant liquid. To the precipitate, add 6 M NH₃ and stir. The appearance of a gray to black precipitate is positive for Hg₂²⁺.

       3. Cations That Form Insoluble Sulfates

          Identification tests for Pb²⁺ and Ag⁺ (Ag₂SO₄ is moderately soluble) are described above (Cations that form Insoluble Chlorides). Ba²⁺, Sr²⁺, and Ca²⁺ form moderately soluble sulfates.

          The alkaline earth ions Mg²⁺, Ca²⁺, Sr²⁺, and Ba²⁺ are one of the best examples of a periodic relationship among the elements of a family. Solubilities of their compounds are graduated nicely and the separations (qualitatively) can be accomplished readily. Flame tests are very important.

          • Barium Ions:

            Barium ions can be identified by precipitation of its insoluble yellow BaCrO₄ salt. If Ca²⁺ or Sr²⁺ are present they will also precipitate in the presence of high concentrations of CrO₄^2-. However, the chromates of Ca²⁺ and Sr²⁺ are moderately soluble; their precipitation can be prevented by addition of acetic acid. This weak acid provides sufficient hydronium ions to lower the CrO₄^2- concentratiion enough to keep CaCrO₄ and SrCrO₄ in solution but to allow the BaCrO₄ to precipitate.

            2CrO₄^2-(aq) + 2H⁺(aq) → Cr₂O₇^2-(aq) + H₂O

            The flame test on the solid chromate is important for confirmation.

            To about 1 mL of solution add 10 drops of 6 M CH₃COOH. Then add a few drops of 0.5 M K₂CrO₄ solution. The appearance of a yellow precipitate indicates the presence of Ba²⁺. To confirm, dissolve the precipitate in concentrated HCl and flame test.

          • Strontium Ions:

            Strontium can be identified, in the absence of calcium, by precipitating its sulfate. To the solution add 0.1 M H₂SO₄ dropwise. The formation of a finely-divided, crystalline, white precipitate indicates the presence of Sr²⁺. (Ba²⁺ must be absent, of course.) To confirm, dissolve the precipitate in concentrated HCl and flame test.

          • Calcium Ions:

            If Ba²⁺ and Sr²⁺ are absent, Ca²⁺ may be precipitated as the oxalate from neutral or alkaline solutions. Test the acidity of the solution with litmus paper. If it is acidic, add 3 M NH₃ until basic. Then add 0.2 M (NH₄)₂C₂O₄ solution. The formation of a white precipitate indicates the presence of Ca²⁺. Confirm by adding a few drops of concentrated HCl and flame testing.

       4. Cations That Form Ammonia Complexes
          • Cadmium Ions:

            Cadmium forms a yellow precipitate with sulfide ion either from a neutral solution containing free Cd²⁺ or from an ammoniacal solution of Cd(NH₃)₄²⁺. Since most sulfides are insoluble, and many of them are black, the presence of other metal ions may make it difficult to detect the yellow color of CdS. Therefore, separations must be as complete as possible before testing for Cd²⁺.

            Cd²⁺(aq) + S^2-(aq) → CdS(s)
            Cd(NH₃)₄²⁺(aq) + S^2-(aq) → CdS(s) + 4NH₃

            To a solution of Cd²⁺ or to a solution thought to contain Cd(NH₃)₄²⁺ add 0.1 M Na₂S solution dropwise. The formation of a yellow precipitate confirms the presence of Cd²⁺.

          • Copper(II) Ions:

            The very distinct deep blue color of the copper ammonia complex can be used to identify Cu²⁺. This identification can be carried out in the presence of other cations which form either colorless ammonia complexes or white precipitates. Thus, Zn²⁺, Cd²⁺, Al³⁺, among others, will not interfere.

            In relatively dilute solutions the color of the ammonia complex may not be intense enough to give an unqualified identification, and some other test for confirmation must be used. Cu²⁺ forms a very insoluble reddish-brown hexacyanoferrate(II).

            2Cu²⁺(aq) + Fe(CN)₆^4-(aq) → Cu₂Fe(CN)₆(s)

            Other cations that react with this reagent to form highly colored precipitates must be absent (Co²⁺ and Fe³⁺ for example). Acidify the test solution with acetic acid. Then add a few drops of 0.1 M potassium hexacyanoferrate(II) solution (K₄Fe(CN)₆). A red-brown precipitate confirms the presence of Cu²⁺.

          • Nickel(II) Ions:

            Nickel(II) is one of the easiest cations to identify. Ni²⁺ forms a red precipitate with dimethylglyoxime in a buffered acid solution. Palladium(II) is the only other cation which forms a precipitate with this reagent. However, a few other cations can interfere. Cobalt(II) preferentially forms a dark brown solution with dimethylglyoxime, and excess reagent must be used in its presence.

            Acidify the solution to be tested with 6 M CH₃COOH. Then add about one mL of 0.2 M NaOOCCH₃ solution. Add dimethylglyoxime solution dropwise. A bright red precipitate is positive for Ni²⁺.

          • Zinc Ions:

            Zinc forms one of the few insoluble white sulfides. It is precipitated from a solution of the ammonia complex. Small traces of cations that form dark colored sulfides will obviously interfere.

            Add an excess of 3 M NH₃ to the test solution, so that any zinc present is in the form of Zn(NH₃)₄²⁺. Then add a few drops of 0.1 M Na₂S solution. A white precipitate indicates the presence of Zn²⁺.

       5. Cations That Form Amphoteric Hydroxides
          • Aluminum Ions:

            Aluminum is generally identified by making use of the amphoteric property of its hydroxide and the red color of the "lake" AlOH₃ forms with the reagent, aluminon. Aluminon is a dye (an organic molecule, usually fairly large, that absorbs visible light). As the Al(OH)₃ precipitates the dye is adsorbed on the Al(OH)₃ particles. The adsorption of the dye is called "laking." Aluminum is a fairly common impurity and care must be taken that trace quantities are not reported. Since most laboratory manipulations are carried out in glass containers, silica gel, which physically resembles aluminum hydroxide, is also a common impurity.

            Adjust the pH of about 1 mL of the test solution (with 3 M NaOH and 3 M HNO₃) to precipitate the hydroxide. Centrifuge the mixture. Remove the mother liquor with a capillary pipet and wash the precipitate with distilled water. Centrifuge the mixture. Remove the mother liquor with a capillary pipet and wash the precipitate with distilled water. Centrifuge the mixture and remove the mother liquor with a capillary pipet. These repeated washings remove other ions from the precipitate. Dissolve the precipitate in 3 M HNO₃. If any precipitate does not dissolve in the nitric acid, remove the supernatant to a clean test tube and discard the residue. Add two drops of aluminon reagent (avoid any excess). Add 3 M NH₃(aq) until the solution is basic. Centrifuge. A red, gelatinous precipitate (sometimes called a red lake) indicates Al³⁺.

            Any precipitate that remains after the addition of the nitric acid is probably silica gel, SiO₂•xH₂O. Silica gel is present in many solutions; it is leached from glass containers. Any silica gel present must be removed before the addition of the aluminon and the ammonia because silica gel will also give a red lake.

            Do not confuse traces of red-brown ferric hydroxide for the red lake. Other precipitates will also form colors with the reagent. The supernatant liquid will be an intense blue-purple color if too much reagent has been added. This color has nothing to do with the presence of aluminum. The color of the reagent is sensitive to changes in pH, (the reagent is an acid-base indicator).

          • Chromium(III) Ions:

            Chromium can be taken through a series of colored tests which leaves no doubt as to its identity. Chromium(III) forms a steel green hydroxide which dissolves in excess strong base to give a deeply green colored solution of the hydroxy complex. Treating this complex with 3% hydrogen peroxide gives the yellow solution of the chromate ion, which upon acidification with dilute nitric acid gives the orange color of dichromate. Treatment of the cold solution of dichromate with 3% hydrogen peroxide gives the intense blue color of a peroxide of chromium. (The actual composition of this peroxide is not known, but it is believed to have the empirical formula CrO₅.) This peroxide readily decomposes to the pale violet color of the original hydrated chromium(III) ion. In low concentrations of dichromate the blue color is fleeting, and attention must be focused on the test tube during the addition of the hydrogen peroxide to avoid missing the color change.

            Cr(OH)₄^- (green) --(H₂O₂)--(OH^-)→ CrO₄^2- (yellow)
            CrO₄^2- --(H⁺)→ Cr₂O₇^2- (orange)
            Cr₂O₇^2- --(H₂O₂)--(HNO₃)→ CrO₅ (blue)→ Cr(H₂O)₆³⁺ (violet)

            The following color changes are all indicative of Cr³⁺. Add an excess of 6 M NaOH to about one mL of test solution. To this green solution add 10 drops of 3% H₂O₂. Heat the test tube in the water bath until the excess H₂O₂ is destroyed as indicated by the cessation of bubbles. Acidify the yellow solution with 3 M HNO₃. Cool the resulting orange solution in an ice bath. To the cooled solution add a drop or two of 3% H₂O₂ and observe the immediate fleeting blue color.

          • Tin(IV) Ions:

            Sn⁴⁺ is most conveniently identified by reduction of Sn⁴⁺ to Sn²⁺ with iron. The Sn²⁺ solution is treated with HgCl₂ solution, whereupon Sn²⁺ is oxidized to Sn⁴⁺ and, simultaneously, HgCl₂ is reduced to Hg₂Cl₂ (a silky, white precipitate). The Hg₂Cl₂ is further reduced by Sn²⁺ to Hg, which appears black.

            Sn⁴⁺(aq) + Fe(s) → Sn²⁺(aq) + Fe²⁺(aq)
            Sn²⁺(aq) + 2HgCl₂(aq) → Sn⁴⁺(aq) + Hg₂Cl₂(s) + 2Cl^-(aq)
            Sn²⁺(aq) + Hg₂Cl₂(s) → Sn⁴⁺(aq) + 2Hg(l) + 2Cl^-(aq)

            Add some concentrated HCl to the solution to be tested for Sn⁴⁺. Place an iron brad (or small iron wire) in this solution and heat in a water bath for 5 minutes. Take the clear solution (filter if necessary) and add HgCl₂ solution dropwise. The appearance of a silky, white precipitate, which then turns black, confirms the presence of tin.

       6. Other Cations
          • Manganese(II) Ions:

            Manganese is easily identified by oxidation of Mn²⁺ to purple MnO₄^- using sodium bismuthate (NaBiO₃). Heat must be avoided to prevent the decomposition of permanganate ion to brown, insoluble manganese dioxide. Chloride ion must be absent, because it reduces permanganate ion to either manganese dioxide or manganese(II) depending upon the conditions.

            Acidify the test solution with 3 M HNO₃. Add solid NaBiO₃ and stir. Centrifuge. If the supernatant has the characteristic purple color of MnO₄^-, Mn²⁺ was present.

            2Mn²⁺ + 14H⁺ + 5NaBiO₃ → 5Bi³⁺ + 5Na⁺ + 7H₂O + 2MnO₄

          • Bismuth(III) Ions:

            Bismuth(III) forms a highly insoluble hydroxide which upon treatment with the hydroxy complex of tin(II) is immediately converted to free bismuth, a black precipitate.

            3Sn(OH)₄^2-(aq) + 2Bi(OH)₃(s) → 2Bi(s) + 3Sn(OH)₆^2-(aq)

            Precipitate Bi³⁺ from the test solution with 3 M NaOH and centrifuge the precipitate. Then, to a solution of tin(II) chloride add with stirring 6 M sodium hydroxide until the precipitate of tin(II) hydroxide which first forms just redissolves. This solution is then added dropwise to the precipitate of bismuth(III) hydroxide. The rapid formation of a black color confirms bismuth.

          • Iron(III) Ions:

            The Fe³⁺ ion is readily identified in a dilute nitric acid solution through the blood red color of its thiocyanate complex. A large excess of reagent should be avoided.

            Fe³⁺(aq) + SCN^-(aq) → Fe(SCN)²⁺(aq)

            Acidify the solution with 3 M HNO₃. Then add a few drops of 0.1 M NH₄SCN solution. The solution turns red if Fe³⁺ is present.

    2. Anion Identification Tests
       1. • Carbonate Ions:

            The most characteristic reaction of carbonate is the formation of carbon dioxide upon treatment with acid:

            CO₃^2-(aq) + 2H⁺(aq) → CO₂(g) + H₂O(l)

            The colorless, odorless carbon dioxide can be identified by bubbling it through a saturated solution of barium hydroxide, with which it forms a white precipitate of barium carbonate.

            CO₂(g) + Ba²⁺(aq) + 2OH^-(aq) → BaCO₃(s) + H₂O(l)

            Assemble a gas-liberation apparatus from a small test tube and a section of bent tubing. Dissolve or suspend a portion of your compound in a small amount of water and place it in the small test tube. Add about 0.5 mL of 6 M HCl and quickly fit the tube into the small test tube, allowing the gas liberated to bubble into a 6" test tube of saturated Ba(OH)₂ solution. The formation of a white precipitate in the large test tube (if the gas liberated is odorless) is a positive test for carbonate. It is imperative to test the gas-liberation apparatus by adding HCl to Na₂CO₃.

          • Sulfide Ions:

            When treated with nonoxidizing acids (HCl, CH₃COOH) sulfides react to liberate H₂S gas (rotten-egg odor). If the sulfide is very insoluble liberation of the gas may require concentrated acids (indeed some sulfides, HgS, CuS, are so insoluble that dissolution requires special treatment). The gas is generally identified by its odor and its precipitation of colored sulfides of various metal ions. Sulfides or hydrogen sulfide also are oxidized to elemental sulfur and sulfate by oxidizing agents such as permanganate, nitric acid, sulfuric acid, Fe(III), etc.

            3H₂S(aq) + 2H⁺(aq) + 2NO₃^-(aq) → 2NO(g) + 4H₂O + 3S(s)

            Acidify a sample with 6 M hydrochloric acid and warm. Cautiously smell the gas evolved and also test the gas with a piece of filter paper moistened with lead acetate solution. A foul smelling gas which turns lead acetate paper black constitutes a positive sulfide test.

          • Sulfate Ions:

            Sulfate is conveniently identified by precipitation of BaSO₄. Other insoluble barium salts contain anions of weak acids (CO₃^2-, SO₃^2-and PO₄^3-). Precipitation of these anions is prevented by acidifying the solution.

            Acidify the test solution with 6 M HCl, and add a few drops of 0.2 M BaCl₂ solution. A white precipitate indicates the presence of SO₄^2-.

          • Nitrate Ions:

            The most notable feature of the chemistry of the nitrate ion is its oxidizing ability as illustrated by the following reactions:

            3Fe²⁺(aq) + 4H⁺(aq) + NO₃^-(aq) → NO(g) + 2H₂O + 3Fe³⁺(aq)

            In the last reaction the nitrogen oxide reacts with excess Fe²⁺ to give the brown complex ion Fe(NO)²⁺. It is the formation of this brown complex that is used to identify NO₃^- (called the brown ring test).

            Acidify about 2 mL of the test solution with 3 M H₂SO₄ and then dissolve one-half spatula full of solid FeSO₄.7H₂O in the acidified solution. Cool the solution and then carefully introduce about 0.5 mL of concentrated H₂SO₄ by allowing it to flow down the side of the tilted test tube. Allow the solution to sit undisturbed so that the sulfuric acid forms a definite layer. The formation of a brown color at the interface of the layer constitutes a positive test for nitrate.

          • Phosphate Ions:

            The precipitation usually used to identify phosphate is the formation of yellow ammonium molybdophosphate from ammonium molybdate in acidic solution.

            12MoO₄^2- + 3NH₄⁺ + PO₄^3- + 24H⁺ → (NH₄)₃[P(Mo₁₂O₄₀)] + 12H₂O

            Acidify the sample with concentrated nitric acid and add several drops in excess. Then treat the solution with ammonium molybdate reagent and warm. The formation of a yellow crystalline precipitate confirms the presence of phosphate.

          • Chloride, Bromide, and Iodide Ions:

            All three of these anions form insoluble silver salts. Although the precipitates are of different colors (AgCl white, AgBr cream, AgI yellow) the colors are difficult to distinguish, and confirmatory tests are necessary.

            Silver chloride, the most soluble of the three, dissolves readily in 6 M NH₃ solution because of formation of the ammonia complex. Furthermore, when the solution of the ammonia complex is acidified, AgCl reprecipitates. Neither AgBr nor AgI will dissolve readily in 6 M NH₃, a much higher concentration of NH₃ being required to form the complex.

            Cl^-(aq) + Ag⁺(aq) → AgCl(s)
            AgCl(s) + 2NH₃(aq) → Ag(NH₃)²⁺(aq) + Cl^-(aq)
            Ag(NH₃)²⁺(aq) + Cl^-(aq) + 2H⁺(aq) → AgCl(s) + 2NH₄⁺(aq)

            Bromide and iodide are usually identified by oxidation to the free elements with chlorine. The elements thus formed are extracted into carbon tetrachloride and identified by their color.

            2Br^-(aq) + Cl₂(g) → Br₂(g) + 2Cl^-(aq)
            2I^- (aq) + Cl₂(g) → I₂(s) + 2Cl^-(aq)

            • Chloride:

              Acidify the test solution with 3 M HNO₃. Then add several drops of 0.1 M AgNO₃. If a white precipitate forms, centrifuge and remove the supernatant. To the precipitate add 6 M NH₃ with stirring. If the precipitate dissolves, add 6 M HNO₃ to the solution. A white precipitate will form if the original test solution contained Cl^-.

            • Bromide and Iodide:

              Acidify the sample with several drops of 6 M HCl and add 4-5 drops of carbon tetrachloride. Then add about 0.5 mL of chlorine water and shake. Appearance of an orange-brown carbon tetrachloride layer indicates the presence of bromide. Formation of a purple layer indicates iodide.