topic12B

School Science Lessons
Topic 12B Reactions of metallic elements and compounds
2012-05-17 SPwp
Please send comments to: J. Elfick@uq.edu.au

Table of contents
12.0.0 Reactions of metallic elements and compounds
12.1.0 Aluminium, Reactions of aluminium, Al
12.2.1 Antimony, Reactions of antimony, Sb
12.2.2 Arsenic, Reactions of arsenic and arsenic compounds, As
12.2.3 Barium, Reactions of barium compounds, Ba
12.2.4 Bismuth, Reactions of bismuth compounds, Bi
12.3.1 Cadmium, Reactions of cadmium compounds, Cd
12.4.1 Calcium, Reactions of calcium and its compounds, Ca
12.5.0 Chromium, Reactions of chromium compounds, Cr
Cobalt, Co
Cobalt compounds
Copper, Cu, natural copper
Copper compounds
12.8.0 Iron, Reactions of iron, Fe
12.9.0 Lead, Reactions of lead, Pb
12.9.3 Lithium, Reactions of lithium with water, Li
12.10.1 Magnesium, Reactions of magnesium, Mg
12.10.2 Magnesium, Reactions of magnesium compounds
12.17.0 Manganese, Reactions of manganese, Mn
12.8.4 Mercury, Reactions of mercury compounds, Hg
12.14.04 Metals, Reactions of metals with concentrated oxidizing acids
12.9.4 Nickel, Reactions of nickel compounds, Ni
12.13.0 Phosphorus, Reactions of phosphorus, P
12.14.0 Potassium, Reactions of potassium, K
3.73 Sodium, Reactions of sodium with water, Na
12.15.0 Silicon, Reactions of silicon, Si
12.16.0 Silver, Reactions of silver, Ag
12.17.1 Strontium, Reactions of strontium compounds, Sr
12.20.0 Tin, Reactions of tin, Sn
12.21.1 Zinc, Reactions of zinc and zinc compounds

12.1.0 Aluminium, Al
12.01.1 Reactions of aluminium, Al, aluminate, AlO₂, thermit reaction
12.01.2 Reactions of aluminium salts
12.14.1 Prepare potash alum from its constituent salts, K₂SO₄.Al₂(SO₄)₃.24H₂O
12.14.2 Prepare potash alum from aluminium foil

12.5.0 Chromium, Reactions of chromium compounds, Cr
12.5.0 Chromium ions in solution, [Cr(H₂O)₆]³⁺
12.5.7 Chromic acid, ionization reaction, H₂CrO₄
12.5.5 Oxidize chromium compounds to chromates, CrO₄^2-
12.14.4 Prepare chrome alum, K₂SO₄.Cr₂(SO₄)₃.24H₂O
12.5.2 Prepare chromium trioxide, CrO₃12.5.6 Prepare potassium dichromate, K₂Cr₂O₇
12.14.4.1 Properties of chrome alum, K₂SO₄.Cr₂(SO₄)₃.24H₂O
12.5.1 Reactions of chromium, Cr, and chromium compounds
12.5.3 Reactions of dichromates, Cr₂O₇^2-, potassium dichromate
12.5.4 Reactions of chromates, CrO₄^2-

12.9.0 Lead, Pb
12.9.1 Reactions of lead (II) salts, Pb²⁺
12.9.2 Reactions of lead (IV) salts, Pb⁴⁺

12.17.0 Manganese, Mn
12.8.2 Prepare manganates
12.8.3 Prepare potassium permanganate, KMnO4
12.8.1 Reactions of manganese (II) salts

12.13.0 Phosphorus, P
12.13.5 Prepare microcosmic salt, Na.NH₄.H.PO₄.4H₂O
12.13.2 Prepare phosphorus trichloride, PCl₃
12.13.3 Prepare phosphorus pentachloride, PCl₅12.13.3.1 Prepare phosphorus pentoxide
12.13.6 Reactions of phosphites, HPO₃^2-
12.13.1 Reactions of phosphorus, P, and phosphates,PO₄^3-
12.13.4 Water with chlorides of phosphorus, PCl₃, PCl₅

12.14.0 Potassium, K
12.14.01 Potassium safety
12.5.6 Prepare potassium dichromate, K₂Cr₂O₇
12.14.1 Prepare potash alum from its constituent salts, K₂SO₄.Al₂(SO₄)₃.24H₂O,
12.14.2 Prepare potash alum from aluminium foil
12.14.5 Superphosphate production

12.15.0 Silicon, Si
Order online: Silicon crystals, metallurgical grade
12.15.1 Prepare silica, SiO₂, and silicon, Si
7.2.5 Prepare silicate gardens
7.2.4 Prepare silicon glass
7.2.4.2 Prepare silicon glass, coloured glass
7.2.4.1 Prepare silicon glass in a furnace
7.2.3 Silicon compounds
7.2.6 Silly putty, silicone, bouncing putty (Dow Corning 3179 dilatant compound) "Tricky Putty"

12.16.0 Silver, Ag
12.16.1 Reactions of silver compounds
12.19.4.0 Reactions of silver halides, photography12.16.2 Recycle silver
12.19.8.6 Recover silver, Ag, from silver chloride, AgCl₂
12.16.2 Recycle silver
15.1.9 Silvering and desilvering, plating and deplating silver

12.20.0 Tin, Sn
12.20.2 Prepare tin (IV) chloride
12.20.1 Reactions of tin and tin compounds

12.01.1 Reactions of aluminium
1. Dissolve aluminium in heated dilute hydrochloric acid and note that hydrogen gas forms.
2Al + 6H⁺ --> 2A1³⁺ + 3H₂ (g)
Hot concentrated sulfuric acid will attack aluminium with the production of sulfur dioxide. Dilute or concentrated nitric acid acts only very slowly on aluminium.
2. Add sodium hydroxide solution to aluminium powder. Hydrogen gas forms from the very rapid reaction.
2Al + 2OH^- + 2H₂O --> 2AlO₂^-+ 3H₂ (g) (AlO₂= aluminate ion)
3. The thermit reaction is a hazardous experiment. Mix aluminium powder or aluminium turnings with iron oxide and ignite the mixture with a burning magnesium ribbon. Do the experiment in the open with observers at least ten metres away. Do not use > 25 g of the reaction mixture. Prepare the reaction mixture in a cut down aluminium drink-can suspended above a bucket or trough of sand to contain the molten iron formed. The mixture may be difficult to ignite, but burns with white heat, producing molten iron that can be tapped from the bottom of the container. Be careful! Burning magnesium ribbon held close to the eyes may cause eye damage. The mixture may react violently if the aluminium particles are too fine. Any trace of moisture in the reactants or container may cause violent evolution of steam and ejection of the white-hot contents
8Al + 3Fe₃O₄ –> 9Fe + 4Al₂O₃
4. Fine particles of aluminium react violently with iodine, especially after a drop of water has been added. A large amount of unreacted iodine is liberated as purple vapour into the air. This reaction should only be done with < 5 g of materials and in a fume cupboard or outdoors. Mix the ingredients in a small ceramic mortar and pestle. All observers must wear eye protection.

12.01.2 Reactions of aluminium salts
1. Add ammonia solution, NH₃ (aq) ("ammonium hydroxide") to aluminium sulfate solution. Note the white precipitate of aluminium hydroxide that is insoluble in excess ammonia solution.
Al³⁺ + 3OH^- --> Al(OH)₃ (s)
2. Add drops of sodium hydroxide to aluminium sulfate solution. Note the white precipitate that dissolves in excess sodium hydroxide to form sodium aluminate. Aluminium hydroxide is amphoteric.
Al(OH)₃ + OH^- --> AlO^2- + 2H₂O
3. Add blue litmus solution to aluminium sulfate solution. The blue litmus turns red. Add sodium carbonate solution and note the production of carbon dioxide. Aluminium salts in solution can act as acids because of hydrolysis.
Al³⁺ + 3H₂O --> Al(OH)₃ + 3H⁺
4. Mix dry aluminium powder with twice its volume of sulfur powder. Put into a test-tube only enough to cover the bottom of the test-tube. Be careful! Larger quantities may explode! Set up a safety screen. Clamp the test-tube vertically and heat with a Bunsen burner. Note the vigorous action where aluminium sulfide is synthesized. Leave to cool then add drops of water. Hydrogen sulfide forms because of the hydrolysis of the aluminium sulfide.
2l + 3S --> Al₂S₃
Al₂S³⁺ 6H₂O --> 2Al(OH)₃ (s) + 3H₂S (g)
Similarly, pass hydrogen sulfide through aluminium sulfate solution to produce the hydroxide, not the sulfide.
5. Mix aluminium sulfate with twice its volume of anhydrous sodium carbonate and heat it on a charcoal block. Note the white infusible mass. Add cobalt nitrate solution and heat again. A bright blue solid forms.
6. Demonstrate this reaction only to senior students. Place < 5 g of aluminium chloride in a beaker in a fume cupboard and add water drop-by-drop. The material will hiss, crackle and release clouds of hydrogen chloride and fine particles. Anhydrous aluminium chloride, AlCl₃, reacts violently with water to form the hydrated salt, while hydrolysing and liberating large amounts of hydrogen chloride gas. Fine aerosol particles may also be generated. Both the hydrogen chloride gas and the fine particles are extremely
irritant to the lungs. Aluminium chloride should only be used in a fume cupboard and only in small amounts. Do not mix aluminium chloride with alkaline materials, e.g. sodium hydroxide, because a violent reaction may occur. Aluminium chloride is exceedingly hygroscopic so keep it in a tightly-sealed plastic container. Purchase the material only in small amounts, e.g. 100 g. Aluminium bromide, AlBr₃, has dangerous properties similar to anhydrous aluminium chloride.

12.2.1 Reactions of antimony
See: Tartar emetic
1. Prepare antimony sulfide colloidal solution. Put 20 drops of yellow ammonium sulfide into a boiling tube full of water. Put tartar emetic in another boiling tube and fill with water. Mix equal volumes of the two solutions to produce the colloidal solution and test it as follows: .
1.1 Add sodium chloride. Precipitation occurs. .
1.2 Add iron (III) hydroxide solution. Coagulation occurs because the particles in the two solutions have opposite charges. Iron (III) hydroxide sol is positively charged and antimony sulfide is negatively charged.
2. The effect of alteration of concentration, hydrolysis of antimony chloride. Put antimony chloride in a test-tube and add 1 mL of water. Note the white precipitate of antimony oxychloride. Add drops of concentrated hydrochloric acid until the white precipitate disappears. Add drops of water until the reappearance of antimony oxychloride, SbOCl.
SbCl₃ + H₂O <--> SbOCl (s) + 2HCl
3.1 Add 2 mL of starch solution to 2 mL of antimony sulfide solution. Add sodium chloride solution. The sodium chloride solution has no effect where the solution is protected by the starch.
3.2 Dilute 2 mL of the antimony sulfide solution with 2 mL of water to act as a control. Add sodium chloride solution. The sodium chloride solution coagulates the control.

12.2.2 Reactions of arsenic and arsenic compounds
Arsenic and arsenic compounds are not use in school science experiments because these substances are very poisonous. Heated arsenic (III) oxide gives off the garlic smell of arsenic and a black ring of arsenic in the test-tube. Arsenic (III) oxide is amphoteric and is slightly soluble in water.

12.2.2.1 Wood treated with copper chrome arsenate (CCA)
Copper chrome arsenate is highly toxic but the amount of arsenic in treated wood, timber, is not thought to be toxic because a person would have to ingest about 20 cm³ of treated timber to be at risk from arsenic poisoning. However, when CCA-treated wood is burnt it forms arsenic vapour so it should not be burnt but disposed of in a landfill.

12.2.3 Reactions of barium compounds
1. Add calcium sulfate solution to barium chloride solution. Heat the solution and leave to cool. Note the white precipitate of barium sulfate that is insoluble in water.
Ba²⁺ + SO₄^2- --> BaSO₄ (s)
2. Add ammonium carbonate solution to barium chloride solution. Note the white precipitate of barium carbonate.
Ba²⁺ + CO₃^2- --> BaCO₃ (s)
3. Add ammonium oxalate solution to barium chloride solution. Note the white precipitate of barium oxalate that is soluble in dilute hydrochloric acid but insoluble in acetic acid.
Ba. + C₂O₄^2- --> BaC₂O₄ (s)
4. Add potassium chromate solution to barium chloride solution. Note the yellow precipitate of barium chromate.
Ba²⁺ + CrO₄^2---> BaCrO₄ (s)
5. Do the flame test on barium compounds and note the flame has flashes of green.

12.2.4 Reactions of bismuth compounds
1. Mix solid bismuth nitrate with anhydrous sodium carbonate and heat it on a charcoal block with a mouth blowpipe. A pink globule of bismuth forms surrounded by brown bismuth oxide Bi₂O₃. Bismuth oxide is used in medical suppository creams.
2. Pass hydrogen sulfide into bismuth nitrate solution acidified with dilute hydrochloric acid. Note the dark brown precipitate of bismuth sulfide that is insoluble in either yellow ammonium sulfide or in sodium hydroxide. Filter the precipitate then wash it into an evaporating basin with dilute nitric acid. Heat the evaporating basin to dissolve the precipitate.
2Bi³⁺ + 3S^2- --> Bi₂S₃ (s)
3. Dissolve bismuth chloride in dilute hydrochloric acid and then pour it into a boiling tube full of water. A white precipitate of bismuth oxychloride forms. Pour some precipitate into a test-tube and add drops of concentrated hydrochloric acid to dissolve the precipitate.
BiCl₃+ H₂O --> BiOCl (s) + 2HCl

12.3.1 Reactions of cadmium compounds
1. Pass hydrogen sulfide into cadmium sulfate solution. Note the bright yellow precipitate of cadmium sulfide.
Cd²⁺ + S^2---> CdS (s)
2. Add 3 cm of cadmium sulfate solution in a test-tube an equal volume of 5 M concentrated hydrochloric acid. Pass hydrogen sulfide through the solution. No precipitate appears in acid of this concentration.
Repeat the experiment and dilute the solution until the yellow precipitate appears. Cadmium sulfide precipitates incompletely if the solution is too acidic. Filter off some of the yellow cadmium sulfide and show that it is soluble in dilute nitric acid.
CdS + 2H⁺ --> Cd²⁺ + H₂S (g)
3. Add sodium hydroxide solution to cadmium sulfate solution. Note the precipitate of cadmium hydroxide that is insoluble in excess sodium hydroxide.
Cd²⁺ + 2OH^- --> Cd(OH)₂ (s)
4. Add drops of ammonia solution, NH₃ (aq) ("ammonium hydroxide") to cadmium sulfate solution. Note the white precipitate of cadmium hydroxide that dissolves in excess "ammonium hydroxide".

12.4.1 Reactions of calcium and calcium compounds
1. Heat a flake of calcium on wire gauze with a Bunsen burner flame. The calcium burns brilliantly with a red flame and leaves a white residue of calcium oxide. Add drops of water to the calcium oxide in a test-tube and note the vigorous exothermic reaction. Test the solution with red litmus paper that turns blue. Note that calcium oxide is not very soluble in water.
2Ca + O₂ --> 2CaO
CaO + H₂O --> Ca(OH)_{2 (s)}
1.1 Drop a small piece of calcium (not old stock calcium) into a test-tube a quarter full of dilute hydrochloric acid. Press your thumb over the mouth of the test-tube and when you can feel the pressure on your thumb test the gas for hydrogen with a lighted splint. Repeat the experiment with a small piece of magnesium ribbon. The same reaction occurs, but the calcium is obviously more reactive because it slower down group 2 of the periodic table. Usually, the lower an element in the same group of the periodic table, the more reactive it is.
Ca + 2HCl --> CaCl₂ + H₂
Mg + 2HCl --> MgCl₂ + H₂
2. Add ammonium carbonate solution to calcium chloride solution. Note the white precipitate of calcium carbonate.
Ca²⁺ + CO₃^2- --> CaCO₃ (s)
3. Add ammonium oxalate solution to calcium chloride solution. Note the white precipitate of calcium oxalate that is soluble in dilute hydrochloric acid but insoluble in acetic acid.
Ca. + C₂O₄^2- --> CaC₂O₄ (s)
4. Add solutions of calcium salts to potassium chromate solution and to calcium sulfate solution. No precipitate forms with calcium sulfate solution and barium salts with potassium chromate solution.
5. Add sodium phosphate solution to calcium chloride solution. Note the white precipitate of calcium phosphate that is soluble in dilute hydrochloric acid, nitric acid or acetic acid.
3Ca + 2PO₄^3- --> Ca₃(PO₄)₂ (s)
6. Add concentrated hydrochloric acid to dry calcium chloride and do the flame test. Note the brick-red flame and observe the green colour when seen through blue glass.

12.5.0 Chromium ions in solution
The simplest ion is the hexa aqua chromium (III) ion, hexaaquachromium (III) ion, [Cr(H₂O)₆]³⁺, usually shown as Cr³⁺, a complex ion with a violet-blue colour, but when produced in a chemical reaction is often green.
1. The hexaaquachromium (III) ion forms violet-blue pH 3 solutions in water when the water molecule pulls a hydrogen ion off the complex ion.
[Cr(H₂O)₆]³⁺ + H₂O <--> [Cr(H₂O)₅(OH)]²⁺ + H₃O⁺, but usually shown as:
[Cr(H₂O)₆]³⁺ + H₂O <--> [Cr(H₂O)₅(OH)]²⁺ + H⁺ (aq)
2. Heat chromium (III) sulfate solution
The violet-blue chromium (III) sulfate solution turns green.
[Cr(H₂O)₆]^₃₊ + heat --> [Cr(H₂O)₅(SO₄)]⁴⁺
One of the water molecules in the complex ion is replaced by a sulfate ion. Two positive charges are replaced by two negative charges of the sulfate ion.
3. Heat chromium (III) chloride solution
The violet-blue chromium (III) chloride solution turns green.
[Cr(H₂O)₆]^₃₊ + heat --> [Cr(H₂O)₄Cl₂]⁺ green, tetra aqua dichloro chromium (III) ion,
tetraaquadichlorochromium (III) ion
Two of the water molecules in the complex ion are replaced by chloride ions.
4. Chromium ion + sodium hydroxide
The violet-blue chromium ion solution forms a gelatinous light blue precipitate, with excess sodium hydroxide redissolves to form a green solution.
[Cr(H₂O)₆]^₃₊ + 3OH^- sodium hydroxide solution --> 3H₂O + [Cr(H₂O)₃(OH)₃] (s)
A hydrogen is removed from three of the water molecules in the complex ion to form a neutral complex precipitate and water.
[Cr(H₂O)₃(OH)₃] (s) + 3OH^- excess sodium hydroxide solution --> [Cr(OH)₆]^3- + 3H₂O
The precipitate dissolves again to form a solution of green hexa hydroxy chromate (II) ions, hexahydroxychromate (II) ions
[Cr(OH)₆]^3- + H₂O₂ solution + heat --> CrO₄^2-
5. Chromium ion + sodium hydroxide + hydrogen peroxide
The green hexahydroxychromate (II) ions formed by adding excess sodium hydroxide to chromium ion solution are oxidized by heating with hydrogen peroxide solution to form a bright yellow solution of chromate (V) ions, i.e. a change from chromium (III) to chromium (VI).
5. Chromium ion + ammonia solution
The violet-blue hexaaquachromium (III) ion solution forms a light blue precipitate. However, with excess ammonia, most of the precipitate dissolves to form a red-blue solution.
[Cr(H₂O)₆]^₃₊ + 3NH₃ dilute ammonia solution, acting as a base --> [Cr(H₂O)₃(OH)₃] + 3NH₄⁺
A hydrogen removed from three of the water molecules in the complex ion to form a neutral complex precipitate and ammonium ion.
[Cr(H₂O)₆]^₃₊ + 4NH₃ excess concentrated ammonia solution, then left to stand <-->
[Cr(NH₃)₆]³⁺ + 6H₂O
Ammonia replaces water as a ligand in the complex ion to form hexa ammine chromium (III) ions, hexaamminechromium (III) ions. This reaction is a ligand and exchange reaction.
6. Chromium ion + carbonate ion
2[Cr(H₂O)₆]^₃₊ + 3CO₃^2- (aq) --> 2[Cr(H₂O)₃(OH)₃] + 3CO₂ bubbles + 3H₂O
A hydrogen is removed from three of the water molecules in the complex ion to form a neutral complex precipitate, and carbon dioxide and water..

12.5.1 Reactions of chromium and chromium compounds
1. Dry reactions of chromium
Heat a chromium compound on a carbon block and note the green residue of chromium (III) oxide, Cr₂O₃. Heat the residue in a borax bead. Note the emerald green colour in both the oxidizing and reducing flame of the Bunsen burner.
2. Reactions of chromium in solution
Prepare 2 cm of chrome alum solution alkaline with ammonia solution, NH₃ (aq) ("ammonium hydroxide") solution and boil the solution. The green-grey precipitate of chromium hydroxide forms that is soluble in dilute acids.
Cr³⁺ + 3OH^- --> Cr(OH)₃ (s)
3. Reactions of chromium in solution
Add sodium hydroxide solution to 2 cm of chrome alum solution. Note the precipitate of chromium hydroxide that it is soluble in excess of the reagent to give a green solution of sodium chromite.
Cr(OH)₃ + OH^- --> CrO₂^-+ 2H₂O
4. Add sodium carbonate solution or ammonium sulfide solution to 2 cm of chrome alum solution. Note the precipitate of chromium hydroxide. The carbonate and sulfide of chromium are rapidly hydrolysed in
solution.
5. Heat chromium (III) sulfate solution
[Cr(H₂O)₆]³⁺ + heat --> [Cr(H₂O)₅(SO₄)]₄
One of the water molecules in the complex ion is replaced by a sulfate ion.
6. Chromate (VI) - dichromate (VI) equilibrium
CrO₄^2- yellow solution + H⁺ --> Cr₂O₇^2- orange solution --> + OH^- --> CrO₄^2- yellow solution
2 CrO₄^2- + 2H⁺ <--> Cr₂O₇^2- + H₂O (Add hydrogen ions, the equilibrium shifts to the right. Add hydroxide ions and the equilibrium shift to the left as hydroxide ions react with hydrogen ions.) Add dilute sulfuric acid to form orange dichromate ion. | Add sodium hydroxide solution to form yellow chromate ion.

12.5.2 Prepare chromium trioxide, CrO₃
Dissolve 25 g of potassium dichromate in 50 mL of boiling water. Cool the solution to room temperature and very slowly add 35 mL of concentrated sulfuric acid. Leave for two hours, then pour off the liquid from the potassium hydrogen sulfate crystals. Heat the liquid to 85^oC and add 25 mL of dilute sulfuric acid. Evaporate the liquid on a water bath until crystals form on this surface then set aside to crystallize. Filter through glass wool, preferably with suction, and evaporate the filtrate to produce more crystals. To remove traces of sulfuric acid, wash the crystals while still in the filter with concentrated nitric acid. Chromium trioxide is not soluble in nitric acid. Transfer the crystals to a dry evaporating basin and heat in an air oven at 130^oC.
K₂Cr₂O₇ + 2H₂SO₄ --> 2KHSO₄ + 2CrO₃ (s) + H₂O

12.5.3 Reactions of dichromates, potassium dichromate
1. Add one drop of sodium hydroxide solution to 3 cm of potassium dichromate solution. Note the change of colour of the solution from orange to yellow because of the formation of the chromate ion.
Cr₂O₇^2- + 2OH^- --> 2CrO₄^2- + H₂O
2. Add drops of dilute sulfuric acid to 3 cm of potassium dichromate solution. Then pass sulfur dioxide through the solution. The change of colour to green is because of the reduction of potassium dichromate to chromium sulfate. The sulfurous acid is oxidized to sulfuric acid.
Cr₂O₇^2- + 8H⁺ + 3SO₃^2- --> 2Cr³⁺ + 3SO₄^2- + 4H₂O
Hydrogen sulfide and also ethanol can reduce acidified solutions of potassium dichromate.
K₂Cr₂O₇ + 4H₂SO₄+ 3C₂H₅OH --> K₂SO₄ + Cr₂(SO₄)₃+ 7H₂O + 3CH₃.CHO (acetaldehyde)
Cr₂O₇^2- + 8H⁺ + 3X --> 2Cr³⁺ + 3XO + 4H₂O
3. Acidify potassium dichromate solution. Add a 2 cm deep layer of ether above the solution. Be Careful! Add a drop of hydrogen peroxide solution and note the blue colour because of perchromic acid, HCrO_5.
4. Reduce dichromate (VI) ions with zinc and dilute sulfuric acid or hydrochloric acid Add dilute sulfuric acid or hydrochloric acid to zinc and potassium dichromate (VI) solution in a test-tube or flask. Fit cotton wool in the top of the test-tube or flask to allow hydrogen gas to escape but prevent air entering to reoxidize chromium (II) to chromium (III).
Cr₂O₇^2- + 14H⁺ + 3Zn --> 2Cr³⁺ + 7H₂O + 3Zn²⁺ (reduction from +6 to +3 oxidation states, potassium dichromate (VI) solution to chromium (III) ions)
2Cr³⁺ + Zn --> 2Cr²⁺ + Zn²⁺ (reduction from +3 to +2 oxidation states, chromium (III) ions to chromium (II) ions)

12.5.4 Reactions of chromates
1. Add a drop of silver nitrate solution to potassium chromate solution. Note the bricked precipitate of silver chromate.
2. Add potassium chromate solution to the following solutions: 1. lead acetate and 2. barium chloride to form the chromates of the metals as precipitates.
3. Pass hydrogen sulfide into acidified potassium chromate solution. The chromate is reduced to a chromium salt.
2CrO₄^2- + 10H⁺ + 3H₂S --> 2Cr³⁺ + 8H₂O + 3S (s)
4. Pass sulfur dioxide through acidified potassium chromate solution. Sulfurous acid reduces the yellow chromate solution to the green chromium salt.
2CrO₄^2- + 10H⁺ + 3SO₃^2- --> 2Cr³⁺ + 5H₂O + 3SO₄^2-
5. Add 3 drops of a dilute acid to yellow potassium chromate solution. The colour of the solution changes to an orange is because of the formation of the dichromate ion.
2CrO₄^2- + 2H⁺ --> Cr₂O₇^2-+ H₂O

12.5.5 Oxidation of chromium compounds to chromates
Add 1 cc sodium peroxide to a dilute solution of chrome alum, then boil the solution. The yellow colour of the solution shows the presence of sodium chromate. Tests for the chromate ion by acidifying the solution with acetic acid and add lead acetate solution.

12.5.6 Prepare potassium dichromate
1. Dissolve 15 g of potassium chromate in 50 mL of dilute sulfuric acid and evaporate to half the volume. Leave the solution to cool so that potassium dichromate crystals form. Crystallize again from hot water to yield purer crystals.
2K₂CrO₄ + H₂SO₄ --> K₂SO₄ + K₂Cr₂O₇ + H₂O
2. Add potassium hydroxide solution to chromium (III) chloride solution to form a grey-green then dark green precipitate containing [Cr(OH)₆]^3- ions.
[Cr(H₂O)₆]³⁺ hexaaquachromium (III) ion + (NaOH solution) --> [Cr(H₂O)₃(OH)₃] grey-green +
(excess NaOH solution) --> [Cr(OH)₆]^3- dark green hexahydroxochromate (III) ions.
Add hydrogen peroxide solution, then heat the solution to turn yellow as potassium chromate (VI) forms.
[Cr(OH)₆]^3- + (H₂O₂ + heat) --> CrO₄^2-
Add dilute sulfuric acid to the yellow solution to form orange dichromate solution
2CrO₄^2- chromate + 2H⁺ <--> Cr₂O₇^2- dichromate + H₂O (Add H ions equilibrium to right, add OH ions equilibrium to left)
Boil the solution until no more bubbles of oxygen form to decompose any excess hydrogen peroxide. Add concentrated ethanoic acid to acidify the solution. Leave to cool and orange crystals of potassium dichromate form.

12.5.7 Chromic acid
Ionization reactions:
H₂CrO₄ + H₂O <--> H₃O⁺ + HCrO₄^-, K1 = 2 × 10^-1HCrO₄^- + H₂O <--> H₃O⁺ + CrO₄^2-, K2 = 3.2 × 10^_-7

12.8.1 Reactions of manganese (II) salts
1. Add drops of yellow ammonium sulfide solution to manganese (II) chloride solution. Note the pink precipitate.
Mn²⁺ + S^2- --> MnS(s)
This same precipitate occurs if you pass hydrogen sulfide into an alkaline solution of a manganese (II) salt but no precipitate occurs with an acidic solution.
2. Drop sodium hydroxide solution into manganese (II) chloride solution. Note the white precipitate of manganese (II) hydroxide that rapidly turns brown due to atmospheric oxidation. Keep on adding the sodium hydroxide solution and note that the precipitate is not soluble in excess.
Mn²⁺ + 2OH^- --> Mn(OH)₂(s)
3. Repeat (2.) using ammonium hydroxide with same observations.
4. Repeat (2.) after first adding 2 cc of solid ammonium chloride to the manganese (II) chloride solution. No precipitate occurs. The ammonium ion introduced depresses the ionization of the hydroxide.
5. To 1 cc of manganese (II) chloride solution, add 1 mL of sodium hydroxide solution then 2 mL of bromine water or sodium peroxide and heat. The valence 2 oxide or hydroxide is oxidized to the higher valence 4 oxide. manganese dioxide, that forms a dark brown precipitate. The permanganate forms if the manganese (II) salt is heated with excess oxidizing agent. Boil some of the manganese (II), chloride solution with a 2 cc of lead dioxide and 1 mL of concentrated nitric acid. Dilute with water and filter. The solution comes through showing the pink permanganate colour.

12.8.2 Prepare manganates
Heat on a crucible lid a piece of potassium hydroxide, crystals of potassium nitrate and some manganese dioxide until the whole mass has fused. Leave to cool and add some water and filter. A deep green solution of potassium manganate forms. The O₂ comes from the KNO₃.
4KOH + 2MnO₂ + O₂ --> 2K₂MnO₄+ 2H₂O
The solution is unstable and is readily hydrolysed by dilute acids and even by largely diluting the solution into a permanganate.
3K₂MnO₄ + 2H₂O --> 2KMnO₄ + MnO₂ + 4KOH
Dilute the green solution ten times with water and boil. Note the pink colour of the permanganate on allowing the solution to settle.

12.8.3 Prepare potassium permanganate, KMnO₄
Put 10 g of potassium hydroxide on a sand tray and heat gently to melt. Add 1 g of potassium chlorate, mixing with a glass rod. Add 7.5 g of manganese dioxide very gradually and continuing to stir. When all has been added heat to bright red heat for fifteen minutes. Leave to cool, break off the melt and grind to a fine powder in a mortar. Transfer the powder to a 500 mL flask, add 200 mL of water and heat on a gauze. During the boiling pass carbon dioxide into the solution. At first the solution is green and contains potassium manganate:
6KOH + 3MnO₂ + KClO₃ --> 3K₂MnO₄ + KCl + 3H₂O
Later it becomes purple potassium permanganate:
3K₂MnO₄ + 2CO₂ --> 2KMnO₄+ MnO₂ + 2K₂CO₃
After boiling for ten minutes let cool and pour off through glass wool in a filter. Wash the flask and return the filtrate to it. Boil for a further ten minutes while carbon dioxide passes through, withdraw a drop of solution and place on a filter paper. If the drop has a green centre continue the process until complete. Finally cool and filter. Transfer the solution to a large evaporating basin and evaporate until crystals appear, then set aside to crystallize.

12.8.4 Reactions of mercury compounds
Mercury metal and mercury compounds are not use in school science experiments because these substances are poisonous. Mercury forms mercury (I) and mercury (II) salts. Mercury is very low in the electrochemical series and so nearly all metals precipitate it from solution. Mercury hydroxide, carbonate and oxide are unstable and produce mercury when heated. Mercury (I) salts in solution form the unusual ion Hg₂²⁺. However, mercury (II) salts ionize only slightly.

12.9.1 Reactions of lead (II) salts, Pb²⁺
1. Add dilute hydrochloric to lead nitrate solution. Note the white precipitate of lead chloride. Wash the precipitate, add four times its volume of water and heat. The precipitate dissolves and precipitates again cooling.
Pb²⁺ + 2Cl^- --> PbCl₂ (s)
2. Add dilute sulfuric acid to lead nitrate solution. Note the white precipitate of lead sulfate. Wash the precipitate, concentrated ammonium acetate solution and heat. The lead sulfate dissolves.
Pb²⁺ + SO₄^2- --> PbSO₄ (s)
3. Add potassium chromate solution to 3 mL of lead nitrate solution. Note the yellow precipitate of lead chromate.
Pb²⁺ + CrO₄^2- --> PbCrO₄ (s)
4. Add potassium iodide solution to 3 mL of lead nitrate solution. Note the yellow precipitate of lead iodide that is soluble in hot water.
Pb²⁺ + 2I^- --> PbI₂
5. Add drops of sodium hydroxide solution to lead nitrate solution. Note the white precipitate of lead hydroxide that is soluble in excess sodium hydroxide solution.
Pb²⁺ + 2OH^- --> Pb(OH)₂ (s)
2 Pb(OH)₂ + 2OH^- --> PbO₂^2- + 2HO (PbO₂^2- = plumbite ion)
6. Pass hydrogen sulfide through lead nitrate solution. Note the black precipitate of lead sulfide. Wash the precipitate, transfer to an evaporating dish, add dilute nitric acid and heat the solution until it boils. Some lead sulfide dissolves forming lead nitrate solution, and some lead sulfide is oxidized to lead sulfate.
Pb²⁺ + S^2- ---> PbS (s)
7. Add drops of dilute sodium hydroxide solution to lead acetate solution until a precipitate forms then disappears. Add hydrogen peroxide solution and heat the solution. Note the brown precipitate of lead dioxide.
8. Add sodium carbonate solution to lead nitrate solution. Note the white precipitate of basic lead carbonate, Pb(OH)₂.2PbCO₃.
3Pb²⁺ + 3CO₃^2- + H₂O --> Pb(OH)₂.2PbCO₃ (s) + CO₂ (g)
Add sodium hydrogen carbonate solution to lead nitrate solution. Note the white precipitate of lead carbonate.
Pb²⁺ + 2HCO₃^- --> PbCO₃ (s) + CO₂ (g) + H₂O

12.9.2 Reactions of lead (IV) salts, Pb⁴⁺
1. Add 2 cc of red lead to 2 cm with glacial acetic acid. Heat the mixture and the red lead dissolves. If a brown precipitate occurs repeat the experiment using less red lead. Cool under the tap to precipitate white crystals of lead tetra-acetate.
Pb₃O₄ + 8CH₃COOH --> Pb(CH₃COO)₄ + 2Pb(CH₃COO)₂ + 4H₂O
Add three times the volume of water to the mixture and heat it to hydrolyse the lead tetra-acetate. Note the brown precipitate of lead dioxide.
Pb(CH₃COO)₄ + 2H₂O --> PbO₂ (s) + 4CH₃COOH
2. Add 2 cc of lead dioxide to 2 cm of concentrated hydrochloric acid and cool under the tap. Filter the mixture and note the golden yellow solution containing lead (IV) chloride. Divide the solution into 3 parts.
PbO₂ + 4HCl --> PbCl₄ + 2H₂O
Heat part A of the yellow lead (IV) chloride solution and test for chlorine. Cool the remaining solution under the tap and leave to crystallize. Note the white crystals of lead (II) chloride.
PbCl₄ --> PbCl₂ + Cl₂ (g)
Add drops of part B of the yellow lead (IV) chloride solution to 880 ammonia solution, NH₃ (aq) ("ammonium hydroxide") Note the fine yellow crystals of ammonium chloroplumbate.
PbCl₄ + 2NH₃ + 2HCl --> (NH4)₂PbCl₆ (ammonium chloroplumbate)
Add drops of sodium hydroxide solution part B of the yellow lead (IV) chloride solution. Note the red gelatinous precipitate that on heating forms lead dioxide as a brown powder.
PbCl₄ + 2H₂O ---> PbO₂ (s) + 4HCl
3. Prepare lead dioxide and lead nitrate.
Slowly add 20 g of red lead to 50 mL of dilute nitric acid and boil for 1 minute. Be careful! Filter the solution while hot. Leave the filtrate to cool and form lead nitrate crystals. Wash the residue of lead dioxide twice with hot water and dry it by gentle heating in an evaporating basin.
Pb₃O₄ + 4HNO₃ --> PbO₂ + 2Pb(NO₃)₂ + 2H₂O

12.9.3 Reactions of lithium with water
Lithium reacts violently with water to form corrosive lithium hydroxide and hydrogen gas that, if mixed with air, may explode if ignited. Lithium reacts vigorously with water and acids and so is usually stored under oil. Lithium floats on paraffin oil so when returning a piece of lithium to the storage container shake the container to recoat the surface with the oil. Handle lithium in the same way as you would handle sodium metal. However, lithium is harder to cut than sodium so used a single piece strong scalpel, not a scalpel with a disposable blade. Lithium is toxic if ingested and corrosive to the skin.
2Li (s) + 2H₂O (l) --> 2LiOH (aq) + H₂ (g)

12.9.4 Reactions of nickel compounds
1. Heat some nickel carbonate in a hard glass test-tube and note the green brown residue of nickel (II) oxide. Heat the nickel (II) oxide in a crucible and black nickel (III) oxide, Ni₂O₃, forms. Dissolve nickel (III) oxide in dilute sulfuric acid to give green nickel (II) sulfate.
2Ni₂O₃ + 4H₂SO₄ --> 4NiSO₄ + 4H₂O + O₂
2. Add sodium hydroxide solution to nickel sulfate solution. Note the light green precipitate of nickel (II) hydroxide, that is stable in air and is soluble in ammonium hydroxide to give a blue solution of a complex ion.
Ni²⁺ + 2OH^- --> Ni(OH)₂(s)
3. Heat a solution of nickel chloride made by dissolving nickel carbonate in hydrochloric acid. Nickel chloride crystals are stable when heated.

12.10.1 Reactions of magnesium
See 15.3.13.01: Cathodic protection
1. Magnesium reacts easily with oxygen in the air to form a protective coating of magnesium oxide. Magnesium burns in oxygen with an intense white flame that can hurt the eyes. So it has been used in fireworks and photographic flashlights.
2Mg + O₂ --> 2MgO + energy
2. Magnesium together with aluminium is used to make light weight alloys for use in aircraft, racing cars, bicycles. Magnesium pencil sharpeners, with the iron blade removed, are about 95% magnesium so a cheap source of magnesium.
3. Magnesium powder slowly forms hydrogen with cold water. Clean a magnesium pencil sharpener of piece of magnesium ribbon with sandpaper to remove the magnesium oxide and add a drop of water. Tiny bubbles of hydrogen gas form.
Mg + H₂O --> MgO + H₂ + energy
4. Burning or molten magnesium reacts with water in a violent exothermic reaction to produce flammable hydrogen gas. This is a very dangerous experiment that has caused injuries in school laboratories. Magnesium powder should never be heated and is too reactive for most school experiments.
Mg + 2H₂O --> Mg(OH)₂ + H₂ + energy
The reaction may be so hot that the magnesium can react with nitrogen in the air.
3Mg + N₂ --> Mg₃N₂ + energy
Do not try to use water to control burning magnesium because more explosive hydrogen gas may be formed.
Mg + H₂O --> MgO + H₂ + energy
Also, the carbon dioxide liberated from a soda acid fire extinguisher may release even more energy
2Mg + CO₂ --> 2MgO +C + energy
5. Burn 6 cm of magnesium ribbon in the air over a piece of paper. Add water to the remaining white magnesium oxide solid, add water in a beaker, boil and test with red litmus paper. The litmus paper slowly turns blue showing the magnesium oxide solution to be weakly alkaline.

12.10.2 Reactions of magnesium compounds
1. Add ammonium carbonate solution to magnesium sulfate solution. Note the white precipitate of ammonium carbonate.
2. Add ammonia solution, NH₃ (aq) ("ammonium hydroxide") to magnesium sulfate solution. Note the white precipitate of magnesium hydroxide.
3. Add of ammonium chloride to magnesium sulfate solution, then add ammonium carbonate solution or ammonia solution, NH₃ (aq) ("ammonium hydroxide") solution. Note white precipitate of basic carbonate forms because the increased concentration of ammonium ion, from the ammonium chloride, suppresses the ionization of the ammonia solution, NH₃ (aq) ("ammonium hydroxide") to leave
insufficient hydroxyl ions to attain the solubility product of magnesium hydroxide.
NH₄OH <--> NH₄⁺ + OH^-
4. Add ammonium chloride and ammonia to magnesium sulfate solution. Add disodium hydrogen phosphate solution. Note the white crystalline precipitate of magnesium ammonium phosphate.
Mg²⁺ + HPO₄^2- + NH₃ --> MgNH₄PO₄ (s)
5. Heat magnesium sulfate crystals on charcoal and let cool. Moisten the white mass with cobalt nitrate solution, heat again, then leave to cool. Note the pink precipitate.
6. Fit a 250 mL flask fitted with a stopper and delivery tube and connect it to a U-tube. Connect the U-tube to a piece of combustion tube. Mix 5 cc each of ammonium, chloride and sodium nitrite in the flask and add 30 mL of water. Put 2 cm of magnesium ribbon loosely in the combustion tube. Heat the flask slowly until a reaction action begins, then remove the flame, and heat the combustion tube. The reaction produces nitrogen which combines with magnesium to form magnesium nitride, Mg₃N₂. The U-tube allows the steam to condense steam and prevent it passing into the combustion tube. Transfer the white nitride to a test-tube, add water and boil. Test for ammonia with litmus paper.
Mg₃N₂ + 6H₂O --> 2NH₃ + 3Mg(OH)₂

12.13.1 Reactions of phosphorus and phosphates
1. Add three drops of the sodium phosphate solution to 5 cm of ammonium molybdate acidified with concentrated nitric acid. The ammonium molybdate must be much in excess. Heat the solution with the heat of the hand. Note the blue precipitate of ammonium phosphomolybdate, (NH₄)₃PMo₁₂O₄₀. The deeper the blue the greater the amount of phosphate
2. Add drops of sodium phosphate solution to a neutral solution of silver nitrate. Note the yellow precipitate of silver phosphate that is soluble in dilute nitric acid and also in ammonia solution, NH₃ (aq) ("ammonium hydroxide")
3Ag⁺ + PO₄^3- --> Ag₃PO₄ (s)
3. Add drops of sodium phosphate to a solution containing magnesia mixture (magnesium sulfate, ammonia, and ammonium chloride to prevent precipitation of magnesium hydroxide) Note the white crystalline precipitate of magnesium ammonium phosphate.
Mg²⁺ + NH⁴⁺ + PO₄^3- --> Mg.NH₄.PO₄ (s)
4. Add drops of iron (III) chloride solution to sodium phosphate solution. Note the buff coloured precipitate that is soluble in dilute mineral acids and also in excess of iron (III) chloride solution.
HPO₄^2- + Fe³⁺ --> FePO₄ (s) + H⁺
5. To convert an orthophosphate to a pyrophosphate, heat 3 cm of disodium hydrogen phosphate to red heat and dissolve the residual sodium pyrophosphate. Note the residual sodium pyrophosphate solution forms a white precipitate with silver nitrate solution and a yellow precipitate with disodium hydrogen
phosphate solution.
2Na₂HPO₄ --> Na₄P₂O₇+ H₂O
6. Prepare orthophosphoric acid. Use a fume cupboard. Add 2 mL of concentrated nitric acid to red phosphorus in an evaporating basin. Heat the basin gently and note the vigorous production of nitrogen dioxide. Add more nitric acid if any phosphorus remains undissolved and heat again. The remaining liquid is orthophosphoric acid solution. Heat the solution to evaporate and form a thick syrup.
P₄ + 20HNO₃ --> 4H₃PO₄ + 20NO₂ (g) + 4H₂O
7. Prepare sodium salts of orthophosphoric acid.
Titrate a dilute solution of phosphoric acid against N sodium hydroxide solution using litmus as an indicator. Suppose x mL of the acid neutralized 25 mL of the alkali. Repeat the titration without litmus. This solution contains mainly disodium hydrogen phosphate from which forms crystals after evaporation to a small volume and leaving to cool. Filter off the crystals, wash with cold water and dry between filter papers.
2NaOH + H₃PO₄ --> Na₂HPO₄+ 2H₂O
8. To prepare sodium dihydrogen phosphate, add x mL of the same phosphoric add solution to 12.5 mL of the sodium hydroxide solution. To obtain trisodium phosphate, add x mL of the same phosphoric acid solution to 37.5 mL of the sodium hydroxide solution. Proceed in both cases to obtain crystals as above.
NaOH + H₃PO₄ --> NaH₃PO₄+ H₂O
3NaOH + H₃PO₄ --> Na₃PO₄ + 3H₂O

12.13.2 Prepare phosphorus trichloride
See diagram 12.13.2: Prepare phosphorus trichloride.
1. Use a fume cupboard. The apparatus must be dry.
Pass carbon dioxide to displace the air. Remove the delivery tube and put sand then 10 g of pieces of dry phosphorus in the retort. The dry sand protects the retort from cracking. Pass dry chlorine through the delivery tube. Spontaneous ignition occurs as the chlorine and phosphorus react to produce phosphorus trichloride. Further chlorine produces yellow phosphorus pentachloride.
2P + 3Cl₂ --> 2PCl₃
PCl₂ + Cl₂ --> PCl₅
2. To purify the phosphorus pentachloride, transfer it to a distilling flask with a two-holes stopper fitted with a thermometer and delivery tube. Attach the delivery tube to a sloping condenser and use another distilling flask with a calcium chloride guard tube as a receiver. Warm the liquid in the distilling flask on a water bath and collect the product until the temperature is 76^oC.

12.13.3 Prepare phosphorus pentachloride
See diagram 12.13.3: Prepare phosphorus pentachloride.
Dry chlorine by passage through wash bottles containing concentrated sulfuric acid. Pass a stream of dry chlorine into the flask and allow phosphorus trichloride to drop slowly into the atmosphere of chlorine. The funnel prevents blocking of the inlet tube by any solid. Phosphorus pentachloride collects as a yellow crystalline solid on the bottom of the flask. Transfer the phosphorus pentachloride to a storage bottle.
PCl₃ + Cl₂ -->- PCl₅

12.13.3.1 Prepare phosphorus pentoxide
Ignite < 5 g of red phosphorus on a heat resistant mat in a fume cupboard and observe the formation of phosphorus pentoxide, P₄O₁₀.

12.13.4 Water with chlorides of phosphorus, PCl₃, PCl₅
1. Add one drop of phosphorus trichloride to 1 cm of water. Hold a rod moistened with silver nitrate near the mouth of the test-tube. The hydrolysis is vigorous and hydrogen chloride forms.
PCl₃ + 3H₂O --> 3HCl (g) + H₃PO₃
2. Repeat the experiment with a piece of solid phosphorus pentachloride the size of half a small pea. Note the vigorous reaction.
PCl₅ + 4H₂O --> 5HCl (g) + H₃PO₄

12.13.5 Prepare microcosmic salt. Na.NH₄.H.PO₄.4H₂O.
1. Put 14 g of sodium phosphate and 2.2 g of ammonium chloride in separate beakers. Dissolve each substance in 10 mL of hot water. Mix the solutions while hot and leave to crystallize. Crystallize again with a minimum of water.
NaHPO₄ + NH₄Cl --> Na(NH₄)HPO₄ + NaCl
2. Heat the microcosmic salt to decompose it into ammonia, water and sodium metaphosphate.
Na(NH₄)HPO₄ --> NaPO₃ + NH₃ (g) + H₂O
3. Dip a loop of red-hot platinum wire in microcosmic salt. Heat the loop to obtain a glassy bead of sodium metaphosphate. Dust the bead with manganese dioxide and heat. again. Note the amethyst colour because of the formation of manganese orthophosphate.

12.13.6 Reactions of phosphites
Phosphorous acid, H₃PO₃, behaves as a dibasic acid.
Add silver nitrate solution to a neutral solution of sodium phosphite, NaHPO_3.Note the white precipitate of silver phosphite that if heated or allowed to stand darkens due to reduction to metallic silver.
HPO₃^2- + 2Ag⁺ + H₂O --> 2Ag(s) + HPO₄^2- + 2H⁺

12.14.01 Potassium safety
Potassium metal is banned in many school systems. Potassium reacts violently with water to form hydrogen gas that may ignite or explode. Destroy potassium safely with 2-methylpropan-2-ol. Potassium forms a crust of yellow oxide that may explode when cut. Use sodium or lithium instead of potassium.

12.14.1 Prepare potash alum, K₂SO₄.Al₂(SO₄)₃.24H₂O, from its constituent salts
Aluminium sulfate can form a double salt with the sulfate of a monovalent metal. With potassium sulfate it forms potash alum, K₂SO₄.Al₂(SO₄)₃.24H₂O. In other alums where Na or (NH₄) = X, they have the formula: X₂SO₄.Al₂(SO₄)₃.24H₂O. In other alums, the sulfate of other trivalent metals, Y, are substituted for aluminium sulfate, and they have the formula: X₂SO₄.Y₂(SO₄)₃.24H₂O, i.e. X⁺.Y³⁺.2SO₄^2-.12H₂O.
The molecular weight of potassium sulfate = 174. The molecular weight of hydrated aluminium sulfate Al₂SO₄.8H₂O = 666.
So equi-molecular weights of these sulfates = 174: 666 = 1.45 g : 5.6 g.
Dissolve 1.45 g of potassium sulfate in 15 mL of water. Dissolve 5.6 g of aluminium sulfate crystals in 20 mL of hot water. Mix the two solutions and leave until the next day. Select a crystal with a regular shape and let it to grow in the solution.. Dissolve the rest of the crystals in the minimum of water and leave to form new crystals..

12.14.2 Prepare potash alum, K₂SO₄.Al₂(SO₄)₃.24H₂O, from aluminium foil. Dissolve 1 g of potassium hydroxide pellets in 40 mL of water and add upwards of 1 g of aluminium foil. When action ceases, excess magnesium must be left over..
2l + 2KOH + 2H₂O --> 2KAlO₂ + 3H₂
Pour off the liquid and slowly add dilute sulfuric acid to it until litmus paper shows the solution is slightly acidic. Heat the solution to evaporate it to a small volume, then leave to cool. Filter off the crystals, then wash and dry them.
2KAlO₂ + 4H₂SO₄ --> K₂SO₄ + Al₂(SO₄)₃ + 4H₂O

12.14.3 Prepare iron (III) ammonium alum, (NH₄)₂SO₄.Fe₂(SO₄)₃.24H₂O
Dissolve 11. 5 g of iron (II) sulfate in 30 mL of dilute sulfuric acid. Add 5 mL of concentrated nitric acid and evaporate the solution to 15 mL. Dissolve 2.7 g of ammonium sulfate in 10 mL of water. Mix the two solutions and leave to crystallize. Choose a crystal with a regular shape and allow it to grow in the solution. Iron (III) alum crystals have an amethyst colour but break down on standing in air due to the formation of basic iron (III) sulfate.

12.14.4 Prepare chrome alum, KCr(SO₄)₂, K₂SO₄.Cr₂(SO₄)₃.24H₂O, Cr₂(SO₄)₃.K₂SO₄.24H₂O
1. Reducing action of ethanol on potassium dichromate in acid solution.
Heat 7.5 g of potassium dichromate (VI) in 50 mL of water and leave to cool. Add 6 mL of concentrated sulfuric acid, stand in ice and stir with a thermometer until cooled to 35^oC. Slowly add by drops 5 mL of ethanol and keep stirring so temperature remains below 50^oC. Leave the solution until the next day in a refrigerator. Separate crystals from the remaining solution, wash with deionized water and dry with filter paper. Choose a good crystal to grow in the solution of potash alum to form an overgrowth.
K₂Cr₂O₇ + 4H₂SO₄+ 3CH₃CH₂OH --> K₂SO₄ + Cr₂(SO₄)₃+ 7H₂O + 3CH₃.CHO (acetaldehyde)
Cr₂O₇^2- + 8H⁺ + 3CH₃CH₂OH --> 2Cr³⁺ + 7H₂O + 3CH₃.CHO (omitting the spectator ions)
2. Dissolve 60 g of potassium chromium sulfate in 100 ml water. Stir common alum (aluminium potassium sulfate) into warm water until it will no longer dissolve. Mix the two solutions to form deep violet-blue crystals. Use a seed crystal in a saturated solution of chrome alum to form large diamond-shaped crystals, similar to potash alum.
3. Mix same molar concentrations of solutions of potassium sulfate and chromium (III) sulfate then allow to crystallize from a saturated solution on a piece of weighted cotton. If the solution is evaporated instead of leaving for crystallization, a mixture of crystals ofpotassium sulfate and chromium (III) sulfate forms.

12.14.4.1 Properties of chrome alum
Chrome alum,potassium chromium (III) sulphate, chromium (III) potassium sulfate,chromium (III) potassium sulfate-12-water, chromium (III) potassium sulfate dodecahydrate, potassium double sulfate of chromium, is similar to potash alum. It can be used in dyeing and in tanning leather.
See 12.5.0: Chromium ions in solution, [Cr(H₂O)₆]³⁺
1. Chrome alum crystals are soluble in water to form an violet-blue acid solution, about pH 3
2. Boil a dilute solution of chrome alum in a test-tube. The violet-blue colour of the solution turn green but when allowed to cool and stand the violet-blue colour returns.
3. Add washing soda solution, Na₂CO₃.10H₂O, to a violet-blue solution of chrome alum in a test-tube. A light green gelatinous precipitate of the hydrogen carbonate forms or a light blue precipitate forms with bubbles of carbon dioxide. Add drops of hydrogen peroxide to the test-tube and boil the contents. A yellow solution of sodium chromate forms. Filter paper wetted with the yellow solution turns green in sulphur dioxide gas.
4. Add dilute ammonia to a solution of chrome alum. A light green precipitate of chromium hydroxide forms. This precipitate, like aluminium hydroxide, has a great attraction for dyes and is used as a mordant to make dyes stick to cloth.

12.14.5 Superphosphate production
See 6.42: Artificial fertilizers
Superphosphate fertilizer, a mixture of calcium sulfate and calcium dihydrogen phosphate, is produced by concentrated sulfuric acid on phosphate rock.
Ca₃(PO₄)₂ + 2H₂SO₄ → Ca(H₂PO₄)₂ + 2CaSO₄Wear goggles if superphosphate dust is produced. Dispose of superphosphate in the garbage to landfill.

12.15.1 Prepare silica and silicon
1. Add 2 mL of dilute hydrochloric acid to a dilute solution of water glass, then heat the solution. Note the white precipitate of hydrated silica.
SiO₃^2- + 2H⁺ --> SiO₂(s) + H₂O
Add sodium hydroxide solution to the precipitate of hydrated silica. Heat the mixture. The precipitate dissolves forming sodium silicate in solution.
SiO₂ + 2OH^- --> SiO₃^2- + H₂O
2. Mix 3 g of dry silica and 1 g of dry magnesium powder and put in a dry test-tube clamped at an angle. Be careful! Do this experiment behind a safety screen! Heat the test-tube slowly with a Bunsen burner. A violent reaction occurs. Leave the mixture to cool. Note the brown pieces of silicon in the exploded mixture.
SiO₂ + 2Mg --> 2MgO + Si
3. Put two pieces of silicon in a crucible and heat them from above with a Bunsen burner. Silicon oxidizes to form silica
Si + O₂ --> SiO₂
4. Add sodium hydroxide solution to amorphous silicon in a test-tube and heat the mixture. Hydrogen forms and sodium silicate remains in solution.
Si + H₂O + 2OH^- --> SiO₃^2- + 2H₂ (g)

12.16.1 Reactions of silver compounds
1. Grind solid silver nitrate with twice its volume of anhydrous sodium carbonate in a mortar. Heat the mixture on charcoal in the reducing flame of a blowpipe. A white bead of metallic silver forms that will not mark paper but will dissolve in dilute nitric acid.
2. Add drops of concentrated hydrochloric acid to silver nitrate solution. (Expensive!) Note the white precipitate of silver chloride. Shake the mixture to coagulate the silver chloride, wash with water and leave to settle.
Ag⁺ + Cl^- --> AgCl (s)
Pour off the water and divide the solid silver chloride into three parts. Part (i): Expose it to light and it turns violet. Part (ii) Add ammonium hydroxide and it dissolves. Part (iii) Heat with concentrated hydrochloric acid and it dissolves.
3. Add drops of potassium chromate solution to silver nitrate solution. Note the brick red precipitate of silver chromate that is soluble in both dilute nitric acid and sodium hydroxide.
2g⁺ + CrO₄² --> Ag₂CrO₄ (s)
4. Add sodium phosphate solution to silver nitrate solution. Note the yellow precipitate of silver phosphate.
3Ag⁺ + PO₄^3- --> Ag₃PO₄ (s)
5. Dilute bench ammonium hydroxide solution to five times its volume with water and slowly add to silver nitrate solution.. Note the first formed brown precipitate of silver oxide that dissolves in excess of ammonia to form a complex ion [Ag(NH₃)₂]⁺.
2gNO₃ + 2NH₄OH --> Ag₂O (s) + 2NH₄NO₃ + H₂O
Similarly, sodium hydroxide precipitates silver oxide but it is not soluble in excess of the reagent.

12.16.2 Recycle silver
Add solid sodium chloride to silver solutions. Decant the clear solution above the precipitate and wash it down the sink. Store the dried precipitate.

12.17.1 Reactions of strontium compounds
1. Add ammonium carbonate solution to strontium nitrate solution. Note the white precipitate of strontium carbonate.
Sr²⁺ + CO₃^2- --> SrCO₃ (s)
2. Add ammonium oxalate solution to strontium chloride solution. Note the white precipitate of strontium oxalate that is soluble in dilute hydrochloric acid but insoluble in acetic acid.
Sr²⁺ + C₂O₄^2- --> SrC₂O₄(s)
3. Add sodium phosphate solution to strontium chloride solution. Note the white precipitate of strontium phosphate that is soluble in dilute hydrochloric, nitric acid or acetic acid.
3Sr. + 2PO₄^3- --> Sr₃(PO₄)₂ (s)
4. Add calcium sulfate solution to strontium nitrate solution. Heat the solution then leave to cool. Note the white precipitate of strontium sulfate that is much more insoluble than calcium sulfate.
Sr²⁺ + SO₄^2- --> SrSO₄ (s)
5. Perform the flame test with strontium nitrate. Note the crimson colour of the flame and observe no change in colour when viewed through blue glass.

12.20.1 Reactions of tin and tin compounds
1. Pass hydrogen sulfide through tin (II) chloride solution. Note the precipitate that is insoluble in dilute hydrochloric acid.
Sn²⁺ + S₂- --> SnS (s)
Filter off the precipitate and wash with distilled water. Transfer it to an evaporating basin and add yellow ammonium sulfide solution. The precipitate dissolves. Oxidation by the free sulfur in the ammonium sulfide occurs, so the S in the equation come from the ammonium sulfide.
(NH₄)₂S + SnS + S --> (NH₄)₂SnS₃ (ammonium thiostannate)
Add dilute acid to the ammonium thiostannate to precipitate tin (IV) sulfide, SnS₂.
2. Add drops of sodium hydroxide solution to tin (II) chloride solution. Note the white precipitate of tin (II) hydroxide that dissolves in excess sodium hydroxide to form sodium stannite.
Sn²⁺ + 2OH^- --> Sn(OH)₂ (s)
Sn(OH)₂ + 2OH^- --> SnO₂^2- + 2H₂O (stannite ion = SnO₂^2-)
3. Add drops of ammonium hydroxide solution to tin (II) chloride solution. Note the white precipitate of tin (II) hydroxide that is not soluble in excess of ammonium hydroxide.
4. To show that Tin (II) chloride is a powerful reducing agent, add tin (II) chloride solution to solutions of the following reagents. Reduction occurs with every reagent. Iron (III) chloride forms pale green iron (II) ions. Potassium permanganate forms manganese (II) ions. Potassium dichromate forms green chromic ions.

12.20.2 Prepare tin (IV) chloride
1. Use a fume cupboard
Put sand into the retort protect the glass during heating, followed by 5 cc of granulated tin. Insert the delivery tube and connect to a chlorine apparatus. Heat the retort while chlorine passes over. Note the ignition of the tin, the fine white crystals in the upper part of the retort, and the yellow distillate of tin (IV) chloride. The white crystals are SnCl₄.5H₂O, due to traces of moisture in the apparatus.
Sn + 2Cl₂ --> SnCl₄
2. Add drops of water to the tin (IV) chloride mixture and heat the mixture. Test for hydrogen chloride. The white precipitate is hydrated tin (IV) oxide or a tin (IV) acid.
SnCl₄ + 4H₂O --> SnO₂.2H₂O + 4HCl
3. Add ammonium hydroxide solution to the tin (IV) chloride mixture and heat the mixture. Divide the suspension of hydrated tin (IV) oxide into two parts. To show the amphoteric nature of hydrated tin (IV) oxide, dissolve part A in sodium hydroxide solution and dissolve part B in hydrochloric acid.
SnO₂.2H₂O + 4HCl -->- SnCl₄ + 4H₂O
SnO₂2H₂O + 2NaOH -->- Na₂SnO₃ + 3H₂O (Na₂SnO_{3 =} sodium stannate)

12.21.1 Reactions of zinc and zinc compounds
1. Hold a piece of zinc foil in the Bunsen burner flame, using tongs. Note the zinc oxide forms that is yellow when hot and white when cold.
2. Add sodium carbonate solution to zinc sulfate solution. Note the white precipitate of basic zinc carbonate, ZnCO₃.2Zn(OH)₂H₂O.
3. Add sodium hydrogen carbonate to zinc sulfate solution. Note the white precipitate of the normal carbonate, ZnCO₃.
4. Add drops of sodium hydroxide solution to zinc sulfate solution. Note the white precipitate of zinc hydroxide that dissolves in excess of sodium hydroxide solution to form sodium zincate. Pass hydrogen sulfide is passed through the sodium zincate solution. Note the white precipitate of zinc sulfide. Zinc hydroxide is amphoteric.
ZnSO₄ + 2NaOH --> Zn(OH)₂ + Na₂SO₄
Zn(OH)₂ + 2NaOH --> Na₂ZnO₂ + 2H₂O (Note: Na₂ZnO_{2 =}sodium zincate)
5. Add drops of ammonium sulfide solution to zinc sulfate solution. Note the white precipitate of zinc sulfide that may be discoloured.
6. Dip a rolled filter paper into a concentrated solution of zinc sulfate with added cobalt nitrate solution. Burn the filter paper on wire gauze and note the remaining green ash, Rinmann's green.
7. Add drops of ammonium hydroxide to zinc sulfate solution. The precipitate of zinc hydroxide dissolves in excess, due to the formation of a complex ion, [Zn(NH₃)₂]²⁺.