School Science Lessons
Topic 12b Reactions of metallic elements
2009-09-17
Please send comments to: J.Elfick@uq.edu.au
See: Interesting websites

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Table of contents
12.1.0 Aluminium
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 salts, Cd
12.4.1 Calcium, Reactions of calcium and its compounds, Ca
12.5.0 Chromium
12.6.0 Cobalt
12.7.0 Copper
12.8.0 Iron
12.9.0 Lead
12.9A. Lithium, Reaction of lithium with water
12.10.1 Magnesium, Reactions of magnesium and its compounds, Mg
12.17 Manganese
12.17a.1 Mercury, Reactions of mercury compounds, Hg
12.17b.1 Nickel, Properties of nickel salts
12.13.0 Phosphorus
12.14.0 Potassium
3.73 Sodium, Reactions of sodium with water
12.15.0 Silicon
12.16.0 Silver
12.16a.1 Strontium, Reactions of strontium compounds, Sr
12.20.0 Tin
12.21.0 Zinc, Reactions of zinc and zinc compounds

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

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12.5.0 Cr
12.5.1 Reactions of chromium and chromium compounds, Cr
12.5.2 Prepare chromium trioxide, CrO₃
12.5.3 Reactions of dichromates, Cr₂O₇^2-
12.5.4 Reactions of chromates, CrO₄^2-
12.5.5 Oxidize chromium compounds to chromates, CrO₄^2-
12.5.6 Prepare potassium dichromate, K₂Cr₂O₇
12.5.7 Chromic acid, ionization reaction
12.14.4 Prepare chrome alum, K₂SO₄.Cr₂(SO₄)₃.24H₂O

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12.6.0 Co
12.6.1 Properties of cobalt salts, Co
12.6.2 Make heat-sensitive paper with cobalt (II) chloride, CuCl₂

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12.7.0 Cu
12.1.3 Conservation of mass in a cycle of copper reactions, the copper cycle experiment
12.7.1 Reactions of copper oxide, CuO
12.7.2 Reactions of copper (II) ions, Cu²⁺
12.7.3 Reactions of copper (I) compounds, Cu⁺
12.7.4 Prepare copper (II) ammonium sulfate (NH₄)₂SO₄.CuSO₄.6H₂O
12.7.5 Prepare cuprammonium sulfate, Cu(NH₃)₄SO₄.H₂O
12.7.6 Prepare copper (I) oxide, CuO
12.7.7 Prepare copper (I) chloride, CuCl

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12.8.0 Fe
12.8.1 Reactions of iron (II) salts and iron (IlI) salts, Prussian blue
12.8.2 Rusting
12.8.3 Oxidation of iron (II) salts
12.8.4 Burn steel wool.
12.8.5 Reduction of iron (IlI) salts
12.8.6 Heat iron filings with powdered sulfur
12.8.7 Prepare iron (II) oxide, FeO
12.8.8 Heat iron (II) sulfide, (FeS₂, pyrite, fool's gold)
12.8.9 Prepare iron (IlI) oxide, Fe₂O₃
12.8.10 Show that black iron oxide is a mixed base.
12.8.11 Iron displaces hydrogen from sulfuric acid to form iron (II) sulfate.
12.8.12 Iron displaces hydrogen from hydrochloric acid to form pale green iron (II) chloride.
12.8.13 Heat hydrated iron chlorides.
12.8.14 Prepare iron (II) ammonium sulfate (NH₄)₂SO₄.FeSO₄.6H₂O.
12.8.15 Detect iron in fruit juice using back tea.
12.14.3 Prepare iron (IlI) ammonium alum (NH₄)₂SO₄.Fe₂(SO₄)₃.24H₂O.

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

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

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12.17.0 Mn
12.17.1 Reactions of manganese (II) salts, Mn
12.17.2 Prepare manganates, MnO₄^2-
12.17.3 Prepare potassium permanganate, KMnO₄
12.17.4 Reactions of permanganate ion, MnO₄^-

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

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

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12.16.0 Ag
12.16.1 Reactions of silver compounds, Ag
12.16.2 Silvering solution
12.19.4.0 Reactions of silver halides, AgCl₂
12.19.8.6, Recover silver from silver chloride, AgCl_2,

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

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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)

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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.
2Al + 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.

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12.2.1 Reactions of antimony
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 (IlI) hydroxide solution. Coagulation occurs because the particles in the two solutions are oppositely charged. Iron (IlI) 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 chloridet solution coagulates the control.

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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.

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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.

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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

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12.3.1 Reactions of cadmium salts
1. Pass hydrogen sulfide into cadmium sulfate solution. Note the bright yellow precipitate of cadmium sulfide.
Cd²⁺ + S^2---> CdS (s) P
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".

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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)}
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, unlike strontium salts 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.

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12.5.1 Reactions of chromium and chromium compounds 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.

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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

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12.5.3 Reactions of dichromates
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
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
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.

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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-

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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.

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12.5.6 Prepare potassium dichromate
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

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12.5.7 Chromic acid
Chromic acid, Chromic (VI) acid is formed when chromic oxide dissolves in water. It may be applied as sodium dichromate solution for cleaning glass and etching plastics.
Ionization reactions:
H₂CrO₄ + H₂O <--> H₃O⁺ + HCrO₄^-, K1 = 2 X 10^-1
HCrO₄^- + H₂O <--> H₃O⁺ + CrO₄^2-, K2 = 3.2 X 10^_-7

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12.6.1 Properties of cobalt salts
1. Heat cobalt carbonate in a small hard-glass tube. The brown residue is cobalt (II) oxide.
CoCO₃ --> CoO + CO₂
Transfer the oxide to a crucible and heat to redness. The black residue is cobalt (II) cobalt (III) oxide, Co₂O₄.
2. Add excess of hot sodium hydroxide solution to 3 cm of a cobalt salt solution. Note the pink precipitate of cobalt (II) hydroxide. Note the change to brown cobalt (III) oxide, Co₂O₃ on exposure to air. Cobalt (II) hydroxide is soluble in ammonia solution, NH₃ (aq) ("ammonium hydroxide") because of the formation of the complex ion [Co(NH)₆]²⁺.
Co²⁺ + 2OH^- --> Co(OH)₂ (s)
3. Add a suspension of bleaching powder to a test-tube containing a cobalt (II) salt solution. Note the black precipitate of cobalt (II) hydroxide. The bleaching powder acts as an alkaline hydroxide and an oxidizer. Divide the precipitate into two parts. Add excess hydrochloric acid to one part to obtain a brown solution of the unstable cobalt (II) chloride. Heat the second part to obtain oxygen gas and a residue containing cobalt (II) oxide. Cobalt (II) hydroxide behaves as a weak base.
4. Evaporate cobalt (II) chloride solution to dryness and note the blue colour of the anhydrous salt. Cobalt chloride is stable compare to hydrated iron (II) chloride.
Write on a piece of paper with dilute cobalt chloride solution. Leave to dry and the writing becomes invisible. Heat the paper from below and writing appears bright blue because of dehydration of the salt. So cobalt chlorine solution is a kind of "invisible ink".

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12.6.2 Make heat-sensitive paper with cobalt (II) chloride
Add cobalt (II) chloride solution to ammonium chloride solution. Dilute this solution until it is pale pink. Soak absorbent paper in the solution and leave to dry and become almost colourless. Heat produces a bright green colour.

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12.7.1 Reactions of copper oxide, CuO
1. Mix copper oxide with fusion mixture and heat it on a charcoal block in the reducing flame of a blowpipe. Brown scales of copper forms.
2. Add concentrated hydrochloric acid to copper oxide on a watch glass. Dip a platinum wire in the mixture for a flame test. Note the intense blue-green flame.
3. Put a borax bead on a platinum wire. Sprinkle copper oxide on the bead and heat in the oxidizing flame of a Bunsen burner. The bead is blue then green when hot

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12.7.2 Reactions of copper (II) ions, Cu²⁺
1. Pass hydrogen sulfide into copper (II) sulfate solution. Note a dark brown precipitate of copper (II) sulfide.
Cu²⁺ + S^2- --> CuS (s)
Wash the precipitate and pour off excess water. Add excess of dilute nitric acid and boil in an evaporating basin. The copper (II) sulfide dissolves.
CuS + 2H⁺ --> Cu²⁺ + H₂S (g)
2. Add potassium iodide solution to copper (II) sulfate solution. A precipitate of white copper (I) iodide and iodine forms. Add sodium thiosulfate solution to dissolve the iodine and note the white precipitate of copper (I) iodide.
2Cu²⁺ + 4I^- --> CuI₂ (s) + I₂ (s)
3. Add sodium hydroxide solution to copper (II) sulfate solution. Note the blue jelly-like precipitate of copper (II) hydroxide.
Cu²⁺ + 2OH^- --> Cu(OH)₂ (s)
Pour the jelly-like precipitate into a test-tube, add ammonia solution, NH₃ (aq) ("ammonium hydroxide") and note the blue precipitate dissolving to form a deep blue solution. This solution contains the cuprammonium ion [Cu(NH₃)4]²⁺. Ammonia has a similar reaction with silver, copper (I) and copper (II) compounds.
Boil the remaining solution and note the black precipitate of copper (II) oxide.
Cu(OH)₂ --> CuO (s) + H₂O
4. Add potassium ferrocyanide solution to copper (II) sulfate solution. Note the brown precipitate of copper ferrocyanide.
2Cu²⁺ + [Fe(CN)₆]^4- --> Cu₂Fe(CN)₆ (s)

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12.7.3 Reactions of copper (I) compounds, Cu⁺
1. Add drops of potassium iodide solution to copper (I) chloride solution in concentrated hydrochloric acid. Note the white precipitate of copper (I) iodide.
Cu₂Cl₂ + 2KI --> Cu₂I₂ (s) + 2KCl
2. Pour some of the solution of copper (I) chloride in concentrated hydrochloric acid into water. Note the white precipitate of copper (I) chloride that is soluble in a high concentration of chloride ions but is insoluble in water.
Cu₂Cl₂ + 4Cl^- <--> 2[CuCl₃]^2-
3. Add dilute hydrochloric acid to 2 cc of copper (I) oxide, then heat. Note the white precipitate of copper (I) chloride.
Cu₂O + 2HCl --> Cu₂Cl₂ + H₂O
4. Add dilute sulfuric acid to 2 cc of copper (I) oxide. Note the red precipitate of metallic copper in a blue solution.
Cu₂O + H₂SO₄ --> CuSO₄ + Cu (s) + H₂O
5. Add dilute nitric acid to 2 cc of copper (I) oxide. The copper produced in the reaction reacts with any excess dilute nitric acid to form blue-green copper nitrate solution, nitric oxide that turns brown on exposure to air, and water.
Cu₂O + 2HNO₃ --> Cu(NO₃)₂ + H₂O + Cu (s)
3Cu + 8HNO₃ --> 3Cu(NO₃)₂ + 2NO + 4H₂O

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12.7.4 Prepare copper (II) ammonium sulfate (NH₄)₂SO₄.CuSO₄.6H₂O
Dissolve 5 g of copper (II) sulfate in 50 mL of boiling water. Dissolve 2.6 g of ammonium sulfate in 10 mL of water. Mix the solutions and evaporate until crystallization begins, then set aside to cool. The crystals are a double salt.

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12.7.5 Prepare cuprammonium sulfate, Cu(NH₃)₄SO₄.H₂O
Dissolve 10 g of copper (II) sulfate by boiling in 50 mL of water in a 200 mL flask and leave to cool. Slowly add concentrated ammonia solution, NH₃ (aq) ("ammonium hydroxide") until any precipitate redissolves, then leave to cool. Be careful!. Add 20 mL of ethanol to form a layer on top of the blue solution. Stopper the flask loosely and leave undisturbed for a week. Filter off the crystals of cuprammonium sulfate and transfer to a container with a stopper. Cuprammonium sulfate is a complex salt in which the copper ion and ammonia form a single divalent ion [Cu(NH)]²⁺] unlike the double salt copper (II) ammonium sulfate, which behaves in solution as would the two sulfates in its molecule.

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12.7.6 Prepare copper (I) oxide, CuO
Fill a boiling tube to the depth of 1 in. with copper (II) sulfate solution and add 2 cc of Rochelle salt, sodium potassium tartrate. When the salt has dissolved, add sodium hydroxide solution. The solution is now Fehling's solution. Add 2 cc of glucose and boil. An orange-red precipitate of copper (I) oxide forms by the reducing action of glucose on the copper (II) copper in solution.
2Cu(OH)₂ - O --> Cu₂O + 2H₂O
Reduction by glucose.
The precipitate is soluble in concentrated hydrochloric acid, but dilute sulfuric acid or nitric acid gives free copper and the copper (II) salt. Excess nitric acid. acts on the copper.
Cu₂O + 2HCl --> CuCl₂ + H₂O
Cu₂O + H₂SO₄ --> Cu + CuSO₄ + H₂O

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12.7.7 Prepare copper (I) chloride, CuCl
1. Use enough copper (II) oxide to cover the bottom of a test-tube. Add five times that volume of concentrated hydrochloric acid then heat the solution. Note the green copper (II) chloride solution.
CuO + 2HCl -->- CuCl₂ + H₂O
Add copper filings of equal volume to the copper (II) oxide used, and boil for 2 minutes. Filter the mixture through glass wool into a beaker of water. Note the white precipitate of copper (I) chloride.
Cu + CuCl₂ --> Cu₂Cl₂ (s)
Pour off the supernatant liquid into two parts. Use part A to show that the copper (I) chloride is soluble in ammonia solution, NH₃ (aq) ("ammonium hydroxide") because of the formation of a complex ion.
Cu₂Cl₂ + 4NH₃ --> 2[Cu(NH₃)]²⁺ + 2Cl^-
Use part B to show that the copper (I) chloride is soluble in concentrated hydrochloric acid, because of the formation of another complex ion with the chloride ion. This complex ion is unstable and decomposes on dilution with water. Cu₂Cl₂ + 4Cl^- --> 2[CuCl₃]^2-
2. Prepare dry copper (I) chloride with a filter pump and Buchner funnel to get the white solid. Wash it with sulfurous acid then glacial acetic acid, then dry it by heating on a water bath. Keep the dry solid in a sealed container.

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12.8.1 Reactions of iron (II) salts and iron (IlI) salts, Prussian blue
1. Add sodium hydroxide or ammonia solution, NH₃ (aq) ("ammonium hydroxide").
Iron (II) salt: Green precipitate of iron (II) hydroxide Fe(OH)₂
Iron (IlI) salt: Red precipitate of iron (IlI) hydroxide, hydroxide Fe(OH)₃.
2. Add acidified potassium permanganate.
Iron (II) salt: Permanganate manganate loses its colour. Iron (II) salts are reducers.
Iron (IlI) salt: Does not reduce.
3. Add potassium ferrocyanide, K₄[Fe(CN)₆].
Iron (II) salt: Light blue precipitate
Iron (IlI) salt: Deep blue precipitate, Prussian blue
Prussian blue as a dye is made by adding iron (II) sulfate to potassium ferrocyanide, with the later addition of iron (III) chloride. Prussian blue can be distilled to yield prussic acid, hydrocyanic acid, HCN, which is very poisonous.
4. Add potassium ferricyanide K₃[FeCN)₆].
Iron (II) salt: Deep blue precipitate "Turnbull's blue
Iron (IlI) salt: Brown colour
5. Add potassium or ammonium thiocyanate solution to a freshly made iron (II) sulfate solution. Iron (II) ammonium sulfate should give a negative result.
Iron (II) salt: No action.
Iron (IlI) salt: Blood red coloration of iron (IlI) thiocyanate Fe(CNS)₃
6. Add iron (III) ions to thiocyanate ion solutions to form bright red complexes, e.g. (Fe(SCN)₃, Fe(SCN)₆^3-. So a thiocyanate solution can be used as a test for iron (III) ions because iron (II) ions do not cause a color change.
7. Iron (II) thiocyanate oxidizes pale green Fe(SCN)₂·3H₂O crystals to red iron(III) thiocyanate and so can be used as a test for the presence of oxygen gas and peroxides.
8. Iron (II) ions and iron (III) ions react with ferrocyanide ion, (Fe(CN)₆⁴⁺) and ferricyanide ion (Fe(CN)₆³⁺) to form the coloured pigment Prussian blue.
K₄Fe(CN)₆ (aq) + Fe³⁺ (aq) --> KFe[Fe(CN)₆] (s) + 3 K⁺ (aq)
9. Iron (II) react with with ferricyanide ions to form the same coloured pigment
K₃Fe(CN)₆ (aq) + Fe²⁺ (aq) --> KFe[Fe(CN)₆] (s) + 2 K⁺ (aq)
10. In blueprinting, the undeveloped paper is covered with iron (III) ferricyanide ion, and citrate. In the light, the citrate reduces the iron (III) to iron (II). With the adsdition of water the deep blue pigment forms.

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12.8.2 Rusting
Rusting forms red hydrated, Fe₂O₃. Wrap a thermometer bulb in wet steel wool results and note the temperature rise. Fe₂O₃ is a red pigment.
4 Fe (s) + 3 O₂ (g) --> 2 Fe₂O₃ (s)

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12.8.3 Oxidation of iron (II) salts
1. Use 2 cm of iron (II) sulfate solution in a test-tube. Add just more than an equal volume of dilute sulfuric acid and three drops of concentrated nitric acid. Heat until the solution boils. Leave to cool and add sodium hydroxide solution until a red precipitate of iron (IlI) hydroxide forms.
6FeSO₄ + 3H₂SO₄ + 2HNO₃ --> Fe₂(SO₄)₃ + 4H₂O + 2NO
Iron (II) ions are oxidized to iron (IlI) ions by electron loss.
Fe²⁺ - e^- --> Fe³⁺
2. Repeat the experiment by substituting other oxidizing materials, e.g. chlorine, bromine, potassium permanganate or hydrogen peroxide, for nitric acid in the above experiment.

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12.8.4 Burn steel wool
Burn steel wool in air to form black magnetite, FeO·Fe₂O₃, that is weakly magnetic.

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12.8.5 Reduction of iron (IlI) salts
1. Put 2 cm of iron (IlI) chloride solution in a test-tube. Pass hydrogen sulfide through the solution until there is no further precipitate of sulfur occurs. Filter the solution and note the pale green solution. Test the filtrate with potassium ferricyanide for proof of iron (II) salt.
2FeCl₃ + H₂S --> 2FeCl₂ + 2HCl + S (s)
Iron (IlI) ions are reduced to iron (II) ions by electron gain.
Fe³⁺ + e^- --> Fe²⁺
Sulfide ions are oxidized by electron loss.
H₂S <--> 2H ⁺ + S^2-
S^2- - 2e^- --> S
2. Add an equal volume of concentrated hydrochloric acid and pieces of granulated zinc to 3 cm of iron (IlI) salt solution. Leave for half an hour then filter. Test the filtrate with excess of sodium hydroxide solution to show that reduction to iron (II) is complete. In the presence of acid, zinc atoms ionize and the electrons are accepted by iron (IlI) ions that are reduced to iron (II) ions.
Zn --> Zn³⁺ + 2e^-
2Fe³⁺ + 2e^- --> 2Fe²⁺

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12.8.6 Heat iron filings with powdered sulfur
Grey iron (II) sulfide forms, FeS. It is ferrimagnetic.
8Fe + S8 --> 8FeS

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12.8.7 Prepare iron (II) oxide, FeO
Close it with a plug of wool a dry test-tube containing 3 cm of iron (II) oxalate. Heat gently then strongly to convert all the yellow oxalate to black iron (II) oxide. Remove the plug of wool and sprinkles the iron (II) oxide into an evaporating basin. The iron (II) oxide spontaneously ignites as it oxidizes to red iron (IlI) oxide.
FeC₂O₄ --> FeO + CO + CO₂
Dissolve the particles left in the test-tube in hydrochloric acid. Test the solution for iron (II) ions. Iron (II) oxide is a base, but iron (II) salts are prepared with metallic iron and acid.

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12.8.8 Heat iron (II) sulfide, FeS₂ (pyrite) fool's gold
Iron (III) oxide and sulfur dioxide forms. (FeS₂ is not iron (IV) sulfide.)
4 FeS₂ (s) + 11 O₂ --> 2 Fe₂O₃ (s) + 8 SO₂ (g)

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12.8.9 Prepare iron (IlI) oxide, Fe₂O₃
Add excess ammonia solution, NH₃ (aq) ("ammonium hydroxide") to an iron (IlI) salt and filter off the iron (IlI) hydroxide. Heat the filter paper and contents in a crucible to leave red iron (IlI) oxide. Boil some oxide in concentrated hydrochloric acid and show that it is a base.

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12.8.10 Show that black iron oxide is a mixed base
Cover the bottom of a test-tube with black iron oxide and add 3 cm of concentrated acid. Heat the solution slowly then filter it. Divide the filtrate into two parts. Test one part for iron (IlI) ions. Test the other part for iron (II) ions. Both ions are present.
Fe₃O₄ + 8HCl -->- 2FeCl₃ + FeCl₂ + 4H₂O

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12.8.11 Iron displace hydrogen from sulfuric acid to form iron (II) sulfate
Fe (s) + H₂SO₄ (aq) --> FeSO₄ (aq) + H₂ (g)
Evaporate the solution to form blue-green crystals of FeSO₄·7 H₂O, green vitriol. In the air, iron(II) salts are oxidized to iron(III) salts, so brown iron (III) hydroxide and iron (III) sulfate may form on the blue-green crystals.

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12.8.12 Iron displaces hydrogen from hydrochloric acid to form pale green iron (II) chloride
Fe (s) + 2 HCl (aq) --> FeCl₂ (aq) + H₂ (g)
Evaporate the solution to form crystals of FeCl₂·4 H₂O. In the air, the iron (II) is oxidized to FeCl₃ and Fe₂O₃

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12.8.13 Heat hydrated iron chlorides
Prepare iron (II) chloride solution by dissolving iron filings in concentrated hydrochloric acid. Evaporate in a test-tube until crystals appear. Heat strongly and test the vapour for hydrogen chloride with silver nitrate solution on a glass rod. Note the residue of iron (IlI) oxide formed when the iron (II) oxide is oxidized in the air.
FeCl₂. + H₂O --> FeO + 2HCl
2FeO + O (air) --> Fe₂O₃
2. Heat iron (IlI) chloride in a test-tube. Test the gas for hydrogen chloride and note the residue of iron (IlI) oxide. Hydrolysis has occurred.
2FeCl₃ + 3H₂O --> Fe₂O₃ + 6HCl

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12.8.14 Prepare iron (II) ammonium sulfate (NH₄)₂SO₄.FeSO₄.6H₂O
Add 4 mL of concentrated sulfuric acid to 30 mL of deionized water in a conical flask. Slowly add 5 g of iron then heat to boiling. Add 10 g of ammonium sulfate and evaporate to two thirds of the original volume. Add a loose stopper loosely and leave the double salt to crystallize. This salt is not an alum.

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12.8.15 Detect iron in fruit juice using back tea
Add strong black tea to samples of fruit juice, e.g. apple, pineapple, cranberry.  Note the time for a cloudy precipitate of  iron compounds to form. The precipitate may not appear for hours or days and the time for precipitation may depend on the temperature and concentrations of the tea and fruit juice. Pineapple juice should give the shortest time for precipitation . The precipitate is formed by a reaction between the ferric, Fe³⁺, non-haeme iron from the fruit juice with the tannins in the black tea. The non-haeme iron is an important component in our diet but black tea may make this iron indigestible so that we cannot absorb it. Perhaps we should drink black tea only between meals and not with meals. The ferrous, Fe²⁺, haeme iron comes mainly from haemoglobin and myoglobin in red meat.

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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

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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

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12.9A.1 Reaction of lithium with water
Lithium reacts vigorously with water and acids and so is usually stored under oil
2Li (s) + 2H₂O (l) --> 2LiOH (aq) + H₂ (g)

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12.10.1 Reactions of magnesium and its compounds, Mg
1. 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.
2. Add ammonium carbonate solution to magnesium sulfate solution. Note the white precipitate of ammonium carbonate.
3. Add ammonia solution, NH₃ (aq) ("ammonium hydroxide") to magnesium sulfate solution. Note the white precipitate of magnesium hydroxide.
4. 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^-
5. 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)
6. 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.
7. 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. Tests for ammonia with litmus paper.
Mg₃N₂ + 6H₂O --> 2NH₃ + 3Mg(OH)₂

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12.13.1 Reactions of phosphorus and phosphates
Phosphorus has 2 main allotropes: 1. white or yellow phosphorus and 2. red phosphorus. White phosphorus, P₄, is very reactive and reacts spontaneously with oxygen in the air to form P₂O_5, so it is stored under water. White phosphorus is too reactive to be used in school science teaching. Red phosphorus has the phosphorus atoms bound in a covalent network and so is less reactive than white phosphorus. Red phosphorus can be stored in air. However, some school systems do not allow red phosphorus to be used in school science experiments.
Phosphoric acid, H₃PO₄, behaves as a tribasic acid although the normal salts are much hydrolysed in solution.
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 yellow precipitate of ammonium phosphomolybdate, NH₃PO₄.12MoO₃.
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 (IlI) chloride solution to sodium phosphate solution. Note the buff coloured precipitate that is soluble in dilute mineral acids and also in excess of iron (IlI) 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
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

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12.13.2 Prepare phosphorus trichloride
See diagram 12.13.2: Prepare phosphorus trichloride
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
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.

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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

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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

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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.

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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 because of reduction to metallic silver.
HPO₃^2- + 2Ag ⁺ + H₂O --> 2Ag (s) + HPO₄^2- + 2H⁺

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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 = 1.45:5.6 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.

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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 stops, excess magnesium must be left over.
2Al + 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

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12.14.3 Prepare iron (IlI) 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 (IlI) alum crystals have an amethyst colour but breakdown on standing in air because of the formation of basic iron (IlI) sulfate.

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12.14.4 Prepare chrome alum, K₂SO₄.Cr₂(SO₄)₃.24H₂O, by the reducing action of ethanol on potassium dichromate in acid solution.

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12.14.5 Superphosphate production
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 + 2CaSO₄

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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 gas forms and sodium silicate remains in solution.
Si + H₂O + 2OH^- --> SiO₃^2- + 2H₂ (g)

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12.16a.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. Do the flame test with strontium nitrate. Note the crimson colour of the flame and observe no change in colour when viewed through blue glass.

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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 1. Expose it to light and it turns violet.
Part 2. Add ammonia solution, NH₃ (aq) ("ammonium hydroxide") and it dissolves.
Part 3. 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.
2Ag⁺ + 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 ammonia solution, NH₃ (aq) ("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]⁺.
2AgNO₃ + 2NH₄OH --> Ag₂O (s) + 2NH₄NO₃ + H₂O
Similarly, sodium hydroxide precipitates silver oxide but it is not soluble in excess of the reagent

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12.16.2 Silvering solution
Deposit a bright silver mirror on glass with two solutions, solution A and solution B:
Solution 1. Dissolve separately 12.5 g of silver nitrate in 100 mL of water and 32.5 g of sodium potassium tartrate in 100 mL of water. Mix the two solutions, heat to 55^oC for 5 minutes. Leave to cool, then pour off the clear liquid from the precipitate, then dilute to 200 mL.
Solution 2. Dissolve 1.5 g of silver nitrate in 12 mL of water. Add dilute ammonia solution until a precipitate forms then redissolve, then dilute to 200 mL.
Mix Solutions 1 and Solution 2. Clean all grease from the glass to be silvered, then suspend it in the solution.

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12.17.1 Reactions of manganese (II) salts, Mn
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 because of 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 ammonia solution, NH₃ (aq) ("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.

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12.17.2 Prepare manganates, MnO₄2-
Heat it 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 is obtained. 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.

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12.17.3 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 it 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.

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12.17.4 Reactions of permanganate ion, MnO₄^-
Permanganates in acidic solution. are very powerful oxidizing agents, especially When the permanganate ion oxidizes in acid solution, the negatively charged permanganate ion becomes the positively charged manganese (II) ion by accepting electrons as is typical of oxidizing agents.
In acid solution:
MnO₄^- + 8H ⁺ + 5e^- --> Mn²⁺ + 4H₂O
In neutral or alkaline solution
MnO₄^- + 2H₂O + 3e^- --> MnO₂ + 4OH^-
1. Add 1 mL of dilute sulfuric acid, followed by 2 mL of hydrogen peroxide to potassium permanganate solution. Tests for oxygen gas with a glowing splint. The solution loses its colour if the potassium permanganate is not in excess.
2MnO₄^- + 6H ⁺ + 5H₂O₂ --> 2Mn²⁺ + 8H₂O + 5O₂ (g)
MnO₄^- + 8H ⁺ + 5e^- --> Mn²⁺ + 4H₂O (permanganate ion reduced)
H₂O₂ - 2e^- --> 2H ⁺ + O₂ (hydrogen peroxide oxidized)
2. Add 1 mL of concentrated hydrochloric acid to potassium permanganate solution. Chlorine gas forms that can be recognized by its smell or by its bleaching action on damp litmus.
2KMnO₄ + 16HCl --> 2KCl + 2MnCl₂ + 8H₂O + 5Cl₂
2MnO₄^- + 16H⁺ + 10Cl^- --> 2Mn²⁺ + 8H₂O + 5Cl₂ (g)
The chloride ion is oxidized to chlorine.
2Cl^- - 2e^- --> Cl₂
3. Add potassium permanganate solution on acidic solutions of the following: Iron (II) sulfate, tin (II) chloride, sodium nitrite, sodium thiosulfate. Oxidation occurs in all cases.
4. Oxidation of a ferrous salt by potassium permanganate, ferrous ammonium sulfate and ferrous sulfate
Ferrous ammonium sulfate, FeSO₄(NH₄)SO₄.6H₂O is used for the quantitative standardization of potassium permanganate solution because ferrous sulfate crystals, FeSO₄.7H₂O form a brown basic sulfate because of efflorescence and oxidation
12FeSO₄ + 3O₂ + 6H₂O --> 4[Fe(OH)₃.Fe₂(SO₄)₃]
Ferrous ammonium sulfate can be found in a pure form and in solution breaks up into ferrous ions, sulfate ions and ammonium ions. However, only the ferrous ions react with the permanganate ion.
5Fe²⁺ + MnO₄^- + 8H⁺ --> Mn²⁺ + 5Fe³⁺ + 4H₂O
5. Oxidation of oxalic acid, H₂C₂O₄, by potassium permanganate
In practice, sodium oxalate solution (Na₂C₂O₄) is used for the quantitative standardization of potassium permanganate solution because it can be found in a more pure form. A temperature of 60^oC is needed for the reaction.
Na₂C₂O₄ --> 2Na⁺ + C₂O₄^2-
C₂O₄^2- + 2H⁺ (from acid) <--> H₂C₂O₄
2MnO4- + 5H₂C₂O₄ + 6H⁺ --> 10CO₂ + 8H₂O + 2Mn²⁺
Oxidation of oxalic acid
H₂C₂O₄ + [O] (from KMnO₄)--> 2CO₂ + H₂O
6. Reaction of potassium permanganate solution with acidified hydrogen peroxide solution
5H₂O₂ + 2MnO₄^- + 6H⁺ --> 2Mn²⁺ + 8H₂O + 5O₂

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12.17a.1 Reactions of mercury compounds, Hg
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.

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12.17b.1 Properties of nickel salts
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, which is stable in air and is soluble in ammonia solution, NH₃(aq) ("ammonium hydroxide") to give a blue solution of a complex ion.
Ni²⁺ + 2OH^- --> Ni(OH)₂(s)
3. Heat nickel chloride solution made by dissolving nickel carbonate in hydrochloric acid. Nickel chloride crystals are stable when heated.

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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 ammonia solution, NH₃ (aq) ("ammonium hydroxide") solution to tin (II) chloride solution. Note the white precipitate of tin (II) hydroxide that is not soluble in excess of ammonia solution, NH₃ (aq) ("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 (IlI) chloride forms pale green iron (II) ions. Potassium permanganate forms manganese (II) ions. Potassium dichromate forms green chromic ions.

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12.20.2 Prepare tin (IV) chloride
See diagram 12.13.2
Do this experiment in 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, because of traces of moisture in the apparatus.
Sn + 2Cl₂ --> SnCl₄
1. Add drops of water to the tin (IV) chloride mixture and heat the mixture. Tests for hydrogen chloride. The white precipitate is hydrated tin (IV) oxide or a tin (IV) acid.
SnCl₄ + 4H₂O --> SnO₂.2H₂O + 4HCl
2. Add ammonia solution, NH₃ (aq) ("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)

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12.21.0 Reactions of zinc and zinc compounds
1. Hold a piece of zinc foil in the Bunsen 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 (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 ammonia solution, NH₃ (aq) ("ammonium hydroxide") to zinc sulfate solution. The precipitate of zinc hydroxide dissolves in excess, because of the formation of a complex ion [Zn(NH₃)₂]²⁺.

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