School Science Lessons
Topic 12B
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12B Reactions of metallic elements and compounds
Table of contents

See: Chemicals, (Commercial)

Aluminium, Al


12.2.1 Antimony, Reactions of Sb

12.2.2 Arsenic, Reactions of As and compounds

12.2.3 Barium, Reactions of Ba compounds

12.2.4 Bismuth, Reactions of Bi compounds

12.18.0 Bromo-compounds

12.3.1 Cadmium, Reactions of cadmium sulfate solution

12.4.1 Calcium, Reactions of Ca and compounds

12.5.0 Reactions of Chromium compounds

Lead, Pb

Magnesium, Mg

Manganese, Mn

Nickel, Ni

12.13.0 Phosphorus, Reactions of phosphorus, P

12.16.0 Silver, Reactions of silver, Ag

12.15.0 Silicon, Si

12.17.1 Strontium, Reactions of Sr compounds

12.20.1 Reactions of tin and tin compoundsbr>
12.21.1 Reactions of zinc and zinc compounds CFCs, chlorofluorocarbons, "Freons" "Freons"

12.18.0 Bromo-compounds

12.5.0 Reactions of chromium compounds
See: Chromium, (Commercial) Chrome alum
12.5.0 Chromium ions in solution, hexaaquachromium ion, [Cr(H2O)6]3+
12.5.7 Chromic acid, ionization reaction, H2CrO4 Prepare chrome alum
12.5.5 Oxidize chromium compounds to chromates, CrO42-
12.5.2 Prepare chromium trioxide, CrO3
12.5.6 Prepare potassium dichromate, K2Cr2O7
12.5.1 Reactions of chromium, Cr, and chromium compounds
12.5.3 Reactions of dichromates, Cr2O72-, potassium dichromate
12.5.4 Reactions of chromates, CrO42-

12.13.0 Reactions of phosphorus, P
See: Phosphorus, (Commercial)
12.13.5 Prepare microcosmic salt, Na.NH4.H.PO4.4H2O
12.13.2 Prepare phosphorus trichloride, PCl3
12.13.3 Prepare phosphorus pentachloride, PCl5 Prepare phosphorus pentoxide
12.13.6 Reactions of phosphites, HPO32-
12.13.1 Reactions of phosphorus, P, and phosphates, PO43-
12.13.4 Phosphorus trichloride with water.

12.15.0 Silica, Si
See: Silicon, (Commercial)
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 Prepare silicon glass, coloured glass Prepare silicon glass in a furnace
7.2.3 Silicon compounds, glass
12.15.2 Silicon reverse-resistance temperature effect
7.2.6 Silly putty, silicone, bouncing putty (Dow Corning 3179 dilatant compound) "Tricky Putty"
Soda-lime glass

12.16.0 Reactions of silver, Ag
See: Silver, (Commercial)
12.16.1 Reactions of silver compounds
12.16.2 Recycle silver.

12.18.0 Bromo-compounds
See: Bromine, (Commercial)
12.18.1 Bromoacetanilide
12.18.2 Bromoacetic acid, bromoethanoic acid
12.18.3 Bromobenzene, phenyl bromide
12.18.4 Bromobutane, butyl bromide
12.18.5 Bromochloromethane, methylene chlorobromide
5.0 Bromocresol green (acid-base indicator)
6.0 Bromocresol purple (acid-base indicator) Bromoeosin, eosin
12.18.7 Bromoethane, ethyl bromide
12.18.21 Bromoform, tribromomethane
12.18.8 Bromohexane, hexyl bromide
12.18.9 Bromomethane, methyl bromide
12.18.10 Bromomethylbenzene, benzyl bromide
12.18.11 Bromomethylpropane-1, 3-diol
12.18.12 Bromophenol
7.0 Bromophenol blue (acid-base indicator)
8.0 Bromophenol red (acid-base indicator)
12.18.13 Bromopropane
1-bromopropane, propyl bromide
2-bromopropane, isopropyl bromide
12.18.14 Bromopropene, isopropenyl bromide
8.0a Bromthymol blue (acid-base indicator) 12.18.15 Bromotoluene
12.18.16 Ethylene dibromide, 1, 2-dibromoethane
12.18.17 Dibromomethane, methylene bromide
12.18.18 Dibromopropane
12.18.19 Dinitrobromobenzene
Eosin, C20H6Br4Na2O5, Acid red 87
12.18.20 Tetrabromophenolphthalein,
12.18.21 Tribromomethane, bromoform.

12.2.1 Reactions of antimony
[Tartar emetic, K(SbO)C4H4O6. H2O, crystalline, poisonous, but used as expectorant and to treat schistosomiasis.]
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 is widespread and abundant in the earth.
It is used in dyes, pigments, medicines, lead shot alloy, glass-making, fireworks.
Arsenic and arsenic compounds are not use in school science experiments because these substances are very poisonous.
The two forms are yellow arsenic, S.G. 1.97 and grey arsenic, metallic arsenic S.G. 5.73, respectively.
It has steel-grey colour and is a very brittle, crystalline, semi-metallic solid, a metalloid solid.
It tarnishes in air.
When heated it oxidizes to arsenous oxide, which has a garlic odour.
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.
Occurs in realgar (As4S4), orpiment (As2S3), arsenolite (As2O3), and arsenopyrite (FeAsS). 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
See: Barium, (Commercial)
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+ + SO42- --> BaSO4 (s).

2. Add ammonium carbonate solution to barium chloride solution.
Note the white precipitate of barium carbonate.
Ba2+ + CO32- --> 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 + C2O42- --> BaC2O4 (s).

4. Add potassium chromate solution to barium chloride solution.
Note the yellow precipitate of barium chromate.
Ba2+ + CrO42- --> 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+ + 3S2- --> Bi2S3(s).

3. Dissolve bismuth chloride in dilute hydrochloric acid, 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 sulfate solution
1. Pass hydrogen sulfide into cadmium sulfate solution.
Note the bright yellow precipitate of cadmium sulfide.
Cd2+ + S2---> 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".

12.5.0 Chromium ions in solution
See: Chromium, (Commercial)
Chromite, FeCr2O4
The simplest ion is the 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-grey" pH 3 solutions in water when the water molecule pulls a hydrogen ion off
the complex ion.
So the complex ion is acting as an acid because it gives an hydrogen ion to a water molecule.
[Cr(H2O)6]3+ + H2O <--> [Cr(H2O)5(OH)]2+ + H3O+, but usually shown simply 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, tetraaquadichlorochromium (III) ion,
hexaaquachromium (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 hexahydroxychromate (II) ions
[Cr(OH)6]3- + H2O2 solution + heat --> CrO42-

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).

6. 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 hexaamminechromium (III) ions.
This reaction is a ligand and exchange reaction.

7. Chromium ion + carbonate ion
2[Cr(H2O)6]3+ + 3CO32- (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
CrO42- yellow solution + H+ --> Cr2O72- orange solution --> + OH- --> CrO42- yellow solution
2 CrO42- + 2H+ <--> Cr2O72- + 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 85oC and add 25 mL of dilute sulfuric acid.
Evaporate the liquid on a water bath until crystals form on this surface, then set it 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 130oC.
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.
Cr2O72- + 2OH- --> 2CrO42- + 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.
Cr2O72- + 8H+ + 3SO32- --> 2Cr3+ + 3SO42- + 4H2O
Hydrogen sulfide and also ethanol can reduce acidified solutions of potassium dichromate.
K2Cr2O7 + 4H2SO4 + 3C2H5OH --> K2SO4 + Cr2(SO4)3 + 7H2O + 3CH3.CHO (acetaldehyde)
Cr2O72- + 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, HCrO5.

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).
Cr2O72- + 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.
2CrO42- + 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.
2CrO42- + 10H+ + 3SO32- --> 2Cr3+ + 5H2O + 3SO42-

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.
2CrO42- + 2H+ --> Cr2O72- + H2O

12.5.5 Oxidize chromium compounds to chromates, CrO42-
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, Na2CrO4.
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) --> CrO42-
Add dilute sulfuric acid to the yellow solution to form orange dichromate solution
2CrO42- chromate + 2H+ <--> Cr2O72- 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 × 10-1
HCrO4- + H2O <--> H3O+ + CrO42-, K2 = 3.2 × 10-7. 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 35oC.
Slowly add by drops 5 mL of ethanol and keep stirring so temperature remains below 50oC.
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)
Cr2O72- + 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 of potassium sulfate and chromium (III) sulfate
forms. Chrome alum
Chrome alum, potassium chromium sulfate, K2SO4.Cr2(SO4)3.24H2O, KCr(SO4)2.12H2O, potassium chromium (III) sulfate,
chromium (III) potassium sulfate, chromium alum, chromium potassium, sulfate, chromium (III) potassium sulfate-12-water,
chromium (III) potassium sulfate dodecahydrate, an alum and mordant, Toxic if ingested
It is a potassium double sulfate of chromium, so it is similar to potash alum.
Chrome alum is used in dyeing and in tanning leather.
See 12.5.0: Chromium ions in solution, [Cr(H2O)6]3+.

1. Dissolve chrome alum crystals in water to form a 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 sulfur 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.
Add small amount of ammonia.
Cr(H2O)63+ --> Cr(H2O)3(OH)3
blue chrome alum --> light green neutral complex of chromium hydroxide
Add excess ammonia that replaces water as a ligand.
Cr(H2O)3(OH)3 --> Cr(NH3)63+
light green chromium hydroxide --> violet hexaamminechromium (III) ion
(In a test-tube, forms a violet over green precipitate.)

5. Oxidize chromium (III) ions to chromium (VI) ions
Add excess sodium hydroxide solution to hexaaquachromium (III) ion solution to form green hexahydroxochromate (III) ion solution.
Cr(H2O)63+ --> Cr(H2O)3(OH)3
Heat with hydrogen peroxide solution to form bright yellow chromate (VI) ion solution.

12.13.1 Reactions of phosphorus and phosphates
See: Phosphorus, (Commercial)
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
3Ag+ + PO43- --> 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+ + PO43- --> 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.
HPO42- + 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 acid-base 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
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
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 76oC.

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 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 Phosphorus trichloride with water
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.

12.13.5 Prepare microcosmic salt
Microsmic salt, ammonium sodium hydrogen phosphate (V)-4-water, [Na(NH4)HPO4.4H2O], (from urine)
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.
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, NaHPO3.
Note the white precipitate of silver phosphite that, if heated or allowed to stand, darkens, because of reduction to metallic silver.
HPO32- + 2Ag + + H2O --> 2Ag(s) + HPO42- + 2H+

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.

12.15.1 Prepare silica and silicon
See: Silicon, (Commercial)
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.
SiO32- + 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- --> SiO32- + 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. Dry 50 g of clean sand in an oven for hours, then grind the sane to powder with a mortar and pestle.
Mix 7g of this dry powdered sand with 8 g dry aluminium powder and 10 g of powdered sulfur.
Shake this mixture to mix the contents, but do not grind this mixture with a mortar and pestle.
Block the hole in bottom of a clay flower pot then put the mixture in the pot.
Fill a 2 cm indentation in the top of the mixture with magnesium powder, or with potassium permanganate + drops of glycerine.
Use scissors to fray one end of a 5 cm piece of magnesium ribbon, then push it into the indentation.
Put the flower pot on a heat resistant pad in a fume cupboard or outside away from combustibles.
Wear protective eye wear and gloves, then use a long stemmed fire lighter to light the other end of the magnesium ribbon.
The reaction produces intense heat, bright orange-red light and spattering, so keep well away from it.
When cool, break the clay pot, separate the silicon residue and use a hammer to break it into small pieces.
Add dilute hydrochloric acid to the residue and note the formation of hydrogen sulfide gas.
AlS3 + HCl --> 2AlCl3 + 3H2S
Put the residue in a sieve, running water through it and note the remaining crystalline globules of silicon.

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 + 2NaOH + H2O --> Na2SiO3 + 2H2 (g)

12.15.2 Silicon reverse-resistance temperature effect
Set up a simple circuit to include 2 X 1.5 V batteries (AA) in series, one 3.2 V lamp in a lampholder, one ammeter or multimeter,
2 crocodile clip wires without insulation.
Complete the circuit by holding each end of a piece or pure silicon with the crocodile clips.
Note the brightness of the lamp and the reading on the ammeter.
Remove the piece of silicon, heat it with a Bunsen burner and replace it in the circuit.
Note the increased brightness of the lamp and increased reading of the ammeter.
Silicon is a semi-conductor with a reverse temperature resistance coefficient, so resistance decreases a temperature increases.
Repeat the experiment by replacing the piece of silicon with an iron nail.
Metals increase resistance with increase of temperature

12.16.1 Reactions of silver compounds
See: Silver, (Commercial)
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+ + PO43- --> 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
See: Strontium, (Commercial)
1. Add ammonium carbonate solution to strontium nitrate solution. Note the white precipitate of strontium carbonate.
Sr2+ + CO32- --> 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+ + C2O42- --> 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 + 2PO43- --> 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+ + SO42- --> 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. CFCs, chlorofluorocarbons, "Freons"
Table, RODP = the relative ozone depletion potential (RODP)
See 3.50.1: Ozone and photochemical smog "Freons"
Compounds of fluorine or fluorine and chlorine with ethane or methane are called freons.
Freons were widely used for refrigerating fluids, aerosols and fire extinguishers.
However, scientists believe that chemicals like freons combine with the ozone (O3) that forms a layer of the atmosphere between the
heights of 15 to 30 km.
A depleted ozone layer allows more high energy radiation from the sun to reach the earth and damage living cells.
The Montreal Protocol of November 1992 recommended the stopping of manufacture and consumption of CFCs, including the
1. Freon 11 (CCl3F, trichlorofluoromethane)
2. Tetrachloromethane (carbon tetrachloride, CCl4, perchloromethane, dry cleaning fluid)
3. 1, 1, 1-trichloroethane (C2H3Cl3, methyl chloroform, electrical equipment cleaner)
4. 1, 1, 2, 2-tetrachloroethane (C2H2Cl4, refrigerant and solvent no longer used in USA)

Modern aerosols are labelled: "NO CFC OZONE FRIENDLY".
Modern refrigerators are labelled: "CFC DEPLETED" or, better still, "NO CFC".
Two fluorocarbons used as refrigerants (Freons) were also used as aerosol propellants.
They are non-flammable, odourless, non-toxic at low concentrations, and chemically inert:
(1.) CFC-11, CCl3F, was used for spraying hair and the body
(2.) CFC-12, CCl2F2 was used in high pressure sprays for insecticides and paints.
They were also used to replace pentane in the production of the foam plastics polyurethane and polystyrene.
CFC-13 (CCl2FCClF2) was used in the electronics and dry cleaning industries.
In the upper atmosphere, UV radiation breaks up CFCs to produce chlorine atoms, which can combine with ozone, O3, to form ClO
and an oxygen molecule, O2.
Then ClO and an oxygen atom, O, combine to produce another O2 and a free chlorine atom, Cl, again.
The initial ozone is lost, and the free chlorine atom can repeat the process.
The chlorine atom may react with methane to form hydrogen chloride, and contribute to acid rain.
Cl + O3 --> ClO + O2
ClO + O --> O2 + Cl

CFCs are persistent with long half lives.
Table, RODP = the relative ozone depletion potential (RODP)
Half life
CFC-11 (Freon 11) 1.00 75 years
CFC-12 (Freon 12) 0.86 112 years
CFC-13 .
90 years
CFC-22 0.05 20 years
CFC-113 0.80 .
CFC-114 0.60 .
1, 1, 1-Trichloroethane 0.15 6.5 years
Carbon tetrachloride 1.11 50 years
Halon-1211, bromochlorodifluoromethane (CBrClF2),
10.00 .
Halon-1301, bromotrifluoromethane (CBrF3) 10.00 . "Freons"
Freon is a registered trademark for non-toxic non-flammable gases invented to avoid danger from leaking refrigerator gases.
It is a name for compounds of ethane or methane with hydrogen atoms substituted by fluorine or chlorine, i.e. CFCs.
The manufacture of Freons is being discontinued because of their ozone-depleting properties.
The term "Freon" is not in favour nowadays.
Examples of Freons include the following:
(1.) Freon 11, CCl3F, .trichlorofluoromethane (CFC-11)
(2.) Freon 12, CCl2F2, dichlorodifluoromethane, (CFC-12), b.p. -30oC (most common refrigerant gas, solvent, in fire extinguishers)
(3.) Freon 21, CHCl2F, dichlorofluoromethane
(4.) Freon 114, CClF2CClF2, dichlorotetrafluoroethane
(5.) Freon 142, CH2CClF2, 1-dichloro1:1difluoroethane
HCCl3 + 2HF --> HCF2Cl + 2HCl
chloroform + hydrogen fluoride --> chlorodifluoromethane + hydrogen chloride

12.18.0 Bromo-compounds
Bromo compounds contain a carbon-bromine bond.
Organohalide, Seaweed, e.g. Asparagopsis taxiformis, produce great quantities of organohalides and produce the "smell" of the sea.

12.18.1 Bromoacetic acid, CH2BrCO2H
Bromoacetic acid, CH2BrCO2H, C2H3BrO2, bromoethanoic acid, colourless crystals, corrodes metals and tissues, citrus harvesting
fruit drop chemical

12.18.2 Bromoacetanilide,
Bromoacetanilide, C8H8BrNO, CH3CONHC6H4Br, 4-bromoacetanilide, para-bromoacetanilide.

12.18.3 Bromobenzene, C6H5Br
Bromobenzene, C6H5Br, phenyl bromide, monobromobenzene, colourless liquid, pungent odour, insoluble in water, denser than water,
skin irritant
<20%, Not hazardous

12.18.4 Bromobutane, C4H9Br, (CH3(CH2)3Br), butyl bromide,
(1.) 1-Bromobutane, C4H9Br, n-Butyl bromide, butyl bromide, insoluble in water, denser than water
(2.) 2-Bromobutane, C4H9Br, sec-butyl bromide, colourless, pale yellow, pleasant odour, insoluble in water, denser than water
(3.) 1-Bromo-2-methylpropane, C4H9Br, isobutyl bromide,
(4.) 2-Bromo-2-methylpropane, 2-methyl-2-bromopropane (tert-butyl bromide), trimethyl bromomethane
trimethyl bromododecane (CH3)3CBr

12.18.5 Bromochloromethane, CH2BrCl
Bromochloromethane, CH2BrCl, methylene chlorobromide, chlorobromomethane, Halon 1011, Fluorocarbon 1011, colourless fluid,
chloroform-like odour, insoluble in water, denser than water, dangerous vapour, toxic fumes when heated, not flammable, in fire
extinguisher liquids.

12.18.7 Bromoethane, CH3CH2Br
Bromoethane, CH3CH2Br, C2H5Br, ethyl bromide, colourless volatile liquid, slightly soluble in water, denser than water, toxic by
inhalation, irritates skin and eyes, used as a solvent

12.18.8 Bromohexane, C6H13Br
Bromohexane, C6H13Br, Br-CH2CH2CH2CH2CH2CH2, hexyl bromide, 1-bromohexane, clear liquid, immiscible with water, stable
compound, combustible, incompatible with strong oxidizing agents and strong bases

12.18.9 Bromomethane, CH3Br
Bromomethane, CH3Br, methyl bromide, colourless, odourless, not flammable, insecticidal, nematicidal, used as a fumigant, but now
phased out because ozone-depleting chemical, fire extinguisher liquid, highly toxic, neurological damage.

12.18.10 Bromomethylbenzene, C7H7Br
Bromomethylbenzene, C7H7Br, benzyl bromide, α-bromotoluene, phenyl methyl bromide, colourless liquid, pleasant odour, slightly
soluble in water, denser than water, corrosive to metals and tissue, toxic by inhalation and skin contact, Highly toxic by all routes
Solution < 20% Not hazardous

12.18.11 Bromomethylpropane-1, 3-diol, C5H10Br2O2
Bromomethylpropane-1, 3-diol, C5H10Br2O2, 2-bromo-2-methyl propane, tert-butylbromide, Highly flammable, irritant, environmental

12.18.12 Bromophenol, C6H5BrO
Bromophenol, C6H5BrO, BrC6H4OH, 2-bromophenol, Toxic by all routes, in crustaceans

12.18.13 Bromopropane
1-bromopropane, CH3CH2CH2Br, C3H7Br, propyl bromide, Toxic, Solution < 24%, Not hazardous
2-bromopropane, CH3CH3CHBr, isopropyl bromide

12.18.14 Bromopropene, C3H5Br
Bromopropene, C3H5Br, CH2CHCH2Br, 2-bromopropene, isopropenyl bromide, acrid smell, Toxic

12.18.15 Bromotoluene
, benzyl bromide, C7H7Br, Organic bromo compound
Catalytic cracking of kerosene: 10.6.4

12.18.16 Ethylene dibromide, 1,2-Dibromoethane
1,2-Dibromoethane, Br(CH2)2Br, C2H4Br2, (CH2BR)2, EDB, ethylene dibromide, ethylene bromide, soil fumigant, insecticide
nematocide, colourless volatile liquid, chloroform-like odour, corrosive and toxic fumes, irritates skin, causes collapse, possible
carcinogen, Toxic by all routes, avoid inhalation, causes sever burning of skin, irritation of eyes and respiratory tract, possible carcinogen
Solution < 1%, Not hazardous
1, 2-dibromoethane, scarce element extracted from sea water, 65 ppm, as bromide ion.
Colourless to brown, heavy, volatile liquid, with a mild sweet odour.
Used in leaded petrol, soil fumigant, grains, fruits, and vegetables, log treatment, preparation of dyes, waxes, plastics, latex, and
vinyl bromide CH2=CHBr.

12.18.17 Dibromomethane, CH2Br2
Dibromomethane, CH2Br2, methylene bromide, methylene dibromide, Toxic
Solution < 12.5%, Not hazardous
Colourless liquid, pleasant odour, insoluble in water, used as solvent and in motor fuel

12.18.18 Dibromopropane
1, 1-dibromopropane, propylene dibromide, C3H6Br2, CH3CHBrCH2Br, Toxic by all routes, avoid inhalation
1, 3-Dibromopropane, Br(CH2)3Br

12.18.19 Dinitrobromobenzene, C6H3BrN2O4
Dinitrobromobenzene, C6H3BrN2O4, 2, 4-dinitrobromobenzene
Reactions of benzene, C6H6: 16.8.1
Reactions of the nitrites, NO2-: 12.11.1 (See 6.)
Tests for unsaturated fats, bromine water test:
Tests for unsaturated hydrocarbons, bromine water test:

12.18.20 Tetrabromophenolphthalein, C20H10Br4O4
Tetrabromophenolphthalein, C20H10Br4O4, disodium salt dye for X-ray examinations

12.18.21 Tribromomethane, CHBr3
Tribromomethane, CHBr3, bromoform (use < 50 mL), colourless liquid, chloroform-like odour, slightly soluble in water,
denser than water, not flammable, Toxic by all routes, eye / lung irritant, may cause nervous system disorders, used to separate minerals,
stabilized by adding ethanol