School Science Lessons
(topic12B)
Metallic elements, reactions
Contents
12.1.0 Metals displace hydrogen from acids
12.2.1 Antimony, Reactions of antimony
12.3.0 Arsenic, Reactions of arsenic
12.4.1 Barium, Reactions of Ba compounds
12.5.1 Bismuth, Reactions of Bi compounds
12.6.0 Bromo-compounds, Reactions of bromo-compounds
12.7.1 Cadmium, Reactions of cadmium sulfate solution
12.8.0 Chromium, Reactions of chromium compounds
14.8.0 Iron, Reaction of iron, Fe
12.9.0 Phosphorus, Reactions of phosphorus compounds
12.10.0 Silica / Silicon, Reactions of silica, Si / silicon, SiO2 compounds
12.11.0 Silver, Reactions of silver compounds
12.12.0 Strontium, Reactions of strontium compounds
12.19.5.0 CFCs, "Freons"
12.3.0 Arsenic, Reactions of arsenic
12.3.1 Arsenic, As
12.3.2 Arsenic compounds
12.3.3 Arsenic, Reactions of As and compounds
12.3.4 Wood treated with copper chrome arsenate, (CCA)
12.6.0 Bromo-compounds, Reactions of bromo-compounds
12.6.1a Bromo-compounds
12.6.1 Bromoacetanilide
12.6.2 Bromoacetic acid, bromoethanoic acid
12.6.3 Bromobenzene, phenyl bromide
12.6.4 Bromobutane, butyl bromide
12.6.5 Bromochloromethane, methylene chlorobromide
12.6.7 Bromoethane, ethyl bromide
12.6.21 Bromoform, tribromomethane
12.6.8 Bromohexane, hexyl bromide
12.6.9 Bromomethane, methyl bromide
12.6.10 Bromomethylbenzene, benzyl bromide
12.6.11 Bromomethylpropane-1,3-diol
12.6.12 Bromophenol
12.6.13 Bromopropane
12.6.14 Bromopropene, isopropenyl bromide
12.6.15 Bromotoluene
12.6.16 Ethylene dibromide, 1,2-dibromoethane
12.6.17 Dibromomethane, methylene bromide
12.6.18 Dibromopropane
12.6.19 Dinitrobromobenzene
12.6.20 Tetrabromophenolphthalein,
12.6.21 Tribromomethane, bromoform.
12.8.0 Chromium, Reactions of chromium compounds
See: Chromium, (Commercial).
12.8.8 Chrome alum
12.8.9 Chromium ions in solution, hexaaquachromium ion, Cr(H2O)63+
12.8.7 Chromic acid, ionization reaction, H2CrO4
Experiments
12.8.5 Oxidize chromium compounds to chromates, CrO42-
12.8.10 Prepare chrome alum
12.8.2 Prepare chromium trioxide, CrO3
12.8.6 Prepare potassium dichromate, K2Cr2O7
12.8.1 Reactions of chromium, Cr, and chromium compounds
12.8.3 Reactions of dichromates, Cr2O72-, potassium dichromate
12.8.4 Reactions of chromates, CrO42-
12.9.0 Phosphorus, Reactions of phosphorus compounds
12.9.1 Prepare microcosmic salt, Na.NH4.H.PO4.4H2O
12.9.2 Prepare phosphorus trichloride, PCl3
12.9.3 Prepare phosphorus pentachloride, PCl5
12.9.4 Prepare phosphorus pentoxide
12.9.5 Reactions of phosphites, HPO32-
12.9.6 Reactions of phosphorus, P, and phosphates, PO43-
12.9.8 Phosphorus trichloride with water
12.10.0 Silica / Silicon, Reactions of silica, Si / silicon, SiO2 compounds
12.10.1 Prepare silicic acid and pure silica
12.10.2 Prepare silica, SiO2, and silicon, Si
12.10.3 Prepare silicate gardens
12.10.4 Prepare silicon glass
12.10.5 Prepare silicon glass, coloured glass
12.10.6 Prepare silicon glass in a furnace
12.10.7 Silicon compounds, glass
12.10.8 Silicon reverse-resistance temperature effect
12.10.9 Silly putty, silicone, bouncing putty, "Tricky Putty"
12.10.10 Silicon tetrachloride with water
12.10.11 Silicon tetrachloride with water
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 sulfideis 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. 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.
4. 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.
12.3.1 Arsenic, As
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, semimetallic 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.
Arsenic occurs in realgar As4S4, orpiment As2S3, arsenolite As2O3, and arsenopyrite FeAsS.
Arsenic, As, (Greek arsenikon yellow orpiment), arsenic trisulfide, As, metalloid, arsenic mineral, natural arsenic.
Arsenic and arsenic compounds are Not permitted in schools.
Arsenic is Toxic if ingested.
A 65 mg dose or repeated smaller doses may be poisonous.
Also, small quantities of arsenic may be carcinogenic.
High levels of arsenic have been reported in rice grains, particularly unpolished rice.
Maximum intake of arsenic occurs when the cook only just covers uncooked rice in the pot with water, to be absobed by rice during cooking.
So use 5 X the volume of uncooked rice for cooking water and discard the excess water when cooking is finished.
In timber treatments, wood preservatives, pesticides, found free and in combined many minerals, three allotropes are yellow, black and main allotrope grey arsenic sublimes at 613oC, and n-type dopant of silicon semiconductors, hardens lead alloys.
White arsenic, arsenic (III) oxide, As4O6, common in sulfide ore deposits, very toxic, rodenticide.
Salvarsan, Erlich's "compound 606" (arsphenamine), was the first drug to treat syphilis.
The most toxic form of arsenic is As3+, which reacts with enzymes in the body.
Agricultural use of arsenic kills plants before concentration is toxic enough for humans.
The proliferation of shallow tube wells in Bangladesh has caused widespread arsenic poisoning.
Atomic number, 33, Relative atomic mass, 74.9216, r.d. 5.72, m.p. = 814oC, b.p. = 613oC (sublimes)
Specific heat capacity, 326 J kg-1 K-1
Arsenic, As, (Greek arsenikon yellow orpiment, i.e. arsenic trisulfide), As, metalloid, arsenic mineral, natural arsenic
Arsenic and arsenic compounds are Not permitted in schools.
Arsenic is Toxic if ingested.
A 65 mg dose or repeated smaller doses may be poisonous.
Also, small quantities of arsenic may be carcinogenic.
12.3.2 Arsenic compounds
Chemicals Not permitted in schools, Australia: Arsenic compounds
Arsenic (III) acid, HAsO3, (formerly arsenious acid)
Arsenic (III) iodide, arsenic triiodide, Not permitted in schools
Arsenic (III) oxide, As4O6, As2O3, arsenolite (white arsenic), claudetite mineral, Not permitted in schools
Arsenic (V) acid, H3AsO4, (formerly arsenic acid), Not permitted in schools
Arsenic (V) oxide, As2O5, arsenic oxide
Arsenic sulfide, AsS, red arsenic, realgar
Arsenic trioxide, As2O3, arsenic (III) oxide, arsenic (III) trioxide, arsenious oxide, arsenious acid, white arsenic, Extremely toxic
Arsenic trisulfide, As2S3, yellow arsenic, arsenic sulfide, orpiment mineral, yellow orpiment pigment and dye, Toxic
(Arabic, az zarnik gold)
Arsenious acid, H3AsO3, weak acid, (Arsenic (III) oxide dissolved in water)
Arsenolite, As2O3, mineral, colourless cubic Arsenopyrite, mispickel, arsenical pyrites, FeAsS, Erinite, Cacodyl, tetramethyldiarsane CH3)2As-As(CH3)2, poisonous liquid, garlic smell
Copper (II) arsenite, AsCuHO3, cupric arsenite, Scheele's green, very toxic
Copper (II) arsenate, Cu3(AsO4)2.4H2O, Cu5H2(AsO4)4.2H2O
Hardite, case-hardening material containing arsenic
Pigments, copper acetoarsenite (green 21, emerald green, Paris green), copper arsenite (Scheele's green)
Pigments, cobalt arsenite and arsenic trisulfide, (yellow 39), Toxic, carcinogenic, Not permitted in schools
White arsenic, arsenic oxide, colourless, hygroscopic, soluble in water, crystalline oxides As2O3 and As2O5, toxic, used for family murder "inheritance powder", leukaemia medicine
12.3.3 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.
12.3.4 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.
12.4.1 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+ + 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.
12.5.1 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.
12.6.1a 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.
12.6.1 Bromoacetanilide
Bromoacetanilide, C8H8BrNO, CH3CONHC6H4Br, 4-bromoacetanilide, para-bromoacetanilide.
12.6.2 Bromoacetic acid, CH2BrCO2H
Bromoacetic acid, CH2BrCO2H, C2H3BrO2, bromoethanoic acid, colourless crystals, corrodes metals and tissues, citrus harvesting.
fruit drop chemical.
12.6.3 Bromobenzene, C6H5Br
Bromobenzene, C6H5Br, phenyl bromide, monobromobenzene, colourless liquid, pungent odour, insoluble in water, denser than water, skin irritant.
<20%, Not hazardous.
12.6.4 Bromobutane, C4H9Br, CH3(CH2)3Br), butyl bromide
1-Bromobutane, C4H9Br, n-Butyl bromide, butyl bromide, insoluble in water, denser than water
2-Bromobutane, C4H9Br, sec-butyl bromide, colourless, pale yellow, pleasant odour, insoluble in water, denser than water,.
1-Bromo-2-methylpropane, C4H9Br, isobtyl bromide,.
2-Bromo-2-methylpropane, 2-methyl-2-bromopropane (tert-butyl bromide), trimethyl bromomethane.
trimethyl bromododecane (CH3)3CBr.
12.6.5 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.
12.6.7 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.
12.6.8 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.
12.6.9 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, causes neurological damage.
12.6.10 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.
12.6.11 Bromomethylpropane-1,3-diol, C5H10Br2O2
Bromomethylpropane-1,3-diol, C5H10Br2O2, 2-bromo-2-methyl propane, tert-butylbromide, highly flammable, irritant, environmental danger.
12.6.12 Bromophenol, C6H5BrO
Bromophenol, C6H5BrO, BrC6H4OH, 2-bromophenol, Toxic by all routes, occurs in crustaceans.
12.6.13 Bromopropane
1-bromopropane, CH3CH2CH2Br, C3H7Br, propyl bromide, Toxic, Solution < 24%, Not hazardous.
2-bromopropane, CH3CH3CHBr, isopropyl bromide.
12.6.14 Bromopropene, C3H5Br
Bromopropene, C3H5Br, CH2CHCH2Br, 2-bromopropene, isopropenyl bromide, acrid smell, Toxic
12.6.15 Bromotoluene, C7H7Br
Bromotoluene, benzyl bromide, C7H7Br, Organic bromo compound.
12.6.16 Ethylene dibromide
1,2-Dibromoethane, Br(CH2)2Br, C2H4Br2, (CH2BR)2, EDB, ethylene dibromide, ethylene bromide
It is a soil fumigant, insecticide, nematocide, colourless volatile liquid, chloroform-like odour, corrosive and toxic fumes.
It irritates skin, causes collapse, possible carcinogen, Toxic by all routes, avoid inhalation, causes sever burning of skin, irritation of eyes and respiratory tract.
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.
12.6.17 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.
12.6.18 Dibromopropane
1,1-dibromopropane, propylene dibromide, C3H6Br2, CH3CHBrCH2Br, Toxic by all routes, avoid inhalation.
1,3-Dibromopropane, Br(CH2)3Br.
12.6.19 Dinitrobromobenzene, C6H3BrN2O4
Dinitrobromobenzene, C6H3BrN2O4, 2,4-dinitrobromobenzene.
12.6.20 Tetrabromophenolphthalein, C20H10Br4O4
Tetrabromophenolphthalein, C20H10Br4O4, disodium salt dye used for X-ray examinations.
12.6.21 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.
12.7.1 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.
2a. 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".
12.8.1 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)63+ + 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.
12.8.2 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.
12.8.3 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).
12.8.4 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.
12.8.5 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.
12.8.6 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)63- ions.
Cr(H2O)63+ 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)63- + (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.
12.8.7 Chromic acid, Ionization reactions
H2CrO4 + H2O --> H3O+ + HCrO4-, K1 = 2 × 10-1.
HCrO4- + H2O --> H3O+ + CrO42-, K2 = 3.2 × 10-7.
12.8.8 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: 12.8.9: Chromium ions in solution, Cr(H2O)63+.
Experiments.
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.
12.8.9 Chromium ions in solution, hexaaquachromium ion, Cr(H2O)63+
The simplest chromium ion is the hexaaquachromium (III) ion, Cr(H2O)63+
It is 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)63+ + H2O --> Cr(H2O)5(OH)2+ + H3O+, but usually shown simply as:
Cr(H2O)63+ + 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)63+ + 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)63+ + 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)63+ + 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)63- + 3H2O.
The precipitate dissolves again to form a solution of green hexahydroxychromate (II) ions.
Cr(OH)6]sup>3- + H2O2 solution + heat --> CrO42-.

* Cr3+ (aq) + 3NH3 (aq) + 3H2O (l) <=> Cr(OH)3 (s) + 3NH+4 (aq)
chromium(III) ion + ammonia + water <=> chromium(III) hydroxide + ammonium ion
Cr(OH)3 dissolves slightly in excess ammonia, but boiling the solution causes the chromium(III) hydroxide to form a green precipitate again.

* Cr3+ (aq) + 3OH− (aq) <=> Cr(OH)3(s) Sodium hydroxide precipitates chromium(III) hydroxide, but the green precipitate dissolves in excess sodium hydroxide.

* 2Cr(OH)−4 (aq) + 3H2O2 (aq) + 2OH− (aq) --> 2CrO2−4 (aq) + 8H2O (l) Hydrogen peroxide oxidizes Cr(III) to Cr(VI) Addition of Ba2+ the solution precipitates the yellow barium chromate ion, CrO2−4,
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
A bright yellow solution of chromate (V) ions forms, 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)63+ + 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 complexion to form a neutral complex precipitate and ammonium ion.
Cr(H2O)63+ + 4NH3 excess concentrated ammonia solution, then left to stand --> Cr(NH3)63+ + 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.
2Cr(H2O)63+ + 3CO32- (aq) --> 2Cr(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.
12.8.10 Prepare chrome alum
Chrome alum, chromium (III) potassium sulfate, KCr(SO4)2, (K2SO4.Cr2(SO4)3.24H2O, Cr2(SO4)3.(K2SO4.24H2O
Chrome alum (dodecahydrate) KCr(SO4)2·12(H2O).
Chromic potassium sulfate dodecahydrate CrH24KO20S2.
Chromium Potassium Sulfate, Chrome Alum, dark, violet-red crystals, soluble in water, efflorescent, used for fixing in some photographic baths.
It is a mordant used in textile dyeing, in ceramic glazes, growing colourful purple crystals.
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.
 4. 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)
12.9.1 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 again.
Note the amethyst colour, because of the formation of manganese orthophosphate.
12.9.2 Prepare phosphorus trichloride
See diagram 12.9.2: Prepare phosphorus trichloride.
1. Use a fume cupboard.
The apparatus must be dry.
Pass carbon dioxide to displace the air.
Remove the delivery tube and put sand, then 10 g of pieces of dry phosphorus, in the retort.
The dry sand protects the retort from cracking.
Pass dry chlorine through the delivery tube.
Spontaneous ignition occurs as the chlorine and phosphorus react to produce phosphorus trichloride.
Further chlorine produces yellow phosphorus pentachloride.
2P + 3Cl2 --> 2PCl3.
PCl2 + Cl2 --> PCl5.
2. To purify the phosphorus pentachloride, transfer it to a distilling flask with a two-holes stopper fitted with a thermometer and delivery tube.
Attach the delivery tube to a sloping condenser and use another distilling flask with a calcium chloride guard tube as a receiver.
Warm the liquid in the distilling flask on a water bath and collect the product until the temperature is 76oC.
12.9.3 Prepare phosphorus pentachloride
See diagram 12.9.3: Prepare phosphorus pentachloride.
Dry chlorine by passage through wash bottles containing concentrated sulfuric acid.
Pass a stream of dry chlorine into the flask and allow phosphorus trichloride to drop slowly into the atmosphere of chlorine.
The funnel prevents blocking of the inlet tube by any solid.
Phosphorus pentachloride collects as a yellow crystalline solid on the bottom of the flask.
Transfer the phosphorus pentachloride to a storage bottle.
PCl3 + Cl2 -->- PCl5.
12.9.4 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.
12.9.5 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+.
12.9.6 Reactions of phosphorus and phosphates
1. Add three drops of the sodium phosphate solution to 5 cm of ammonium molybdate acidified with concentrated nitric acid.
The ammonium molybdate must be much in excess.
Heat the solution with the heat of the hand.
Note the blue precipitate of ammonium phosphomolybdate (NH4)3PMo12O40.
The deeper the blue the greater the amount of phosphate.
2. Add drops of sodium phosphate solution to a neutral solution of silver nitrate.
Note the yellow precipitate of silver phosphate that is soluble in dilute nitric acid and also in ammonia solution, NH3 (aq) ("ammonium hydroxide").
3Ag+ + 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 solution.
To obtain trisodium phosphate, add x mL of the same phosphoric acid solution to 37.5 mL of the sodium hydroxide solution.
Proceed in both cases to obtain crystals as above.
NaOH + H3PO4 --> NaH3PO4 + H2O.
3NaOH + H3PO4 --> Na3PO4 + 3H2O.
12.9.8 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.
12.10.1 Prepare silicic acid and pure silica
White sand is almost pure silica.
However, white silica can be made from water glass, sodium metasilicate (Na2SiO3).
Prepare strong solutions of sodium silicate and sodium hydrogen sulfate, (sodium bisulfate).
Dissolve 5 mL of water glass (sodium silicate) 50 mL of hot water.
Dissolve 5 mL of sodium hydrogen sulfate (sodium bisulfate), in 2 cm of water in a test-tube.
Mix the two solutions together in a beaker.
A jelly-like precipitate of silicic acid, Si(OH)4, forms.
Filter off the gelatinous precipitate then wash it by running hot water through the filter paper.
Use a spoon to scrape the precipitate out of the filter paper on to a metal lid.
Hold the metal lid in a pair of pliers and heat it over a Bunsen burner.
The fine white powder left is pure silica, SiO2.
12.10.2 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.
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.
Grind it 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.
4. 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.
5. 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).
12.10.3 Prepare silicate gardens
1. Mix one part of sodium silicate (IV) (Na2SiO3) with four parts of water to make waterglass.
Gently add crystals of salts to the solution without mixing to make chemical "flowers" grow:
* Chlorides: Co, Fe, Cu, Ni and Pb,
* Sulfates: Al, Fe, Cu, and Ni,
* Nitrates: Co, Fe, Cu and Ni.
2. Put sand 1 cm deep in a 500 mL jar.
Make a 1: 1 mixture of sodium silicate and water (waterglass), and pour it onto the sand to almost fill the jar.
Leave the jar to stand undisturbed for a day.
Drop in crystals of metal salts, e.g. metal hydroxides, iron sulfate, copper (II) sulfate, alum, Epsom salts.
Observe crystals forming "shoots".
Some shoots are directed up by small bubbles.
The metal hydroxide skin formed around the crystal is permeable only to water and not the salt.
The water diffuses in to balance the concentrations each side of the skin until the skin bursts the skin then forms again further from the crystal.
12.10.4 Prepare silicon glass
Pick up sodium carbonate in a nichrome wire loop.
Dip the loop into powdered silica and heat over a burner to form a transparent bead of glass.
12.10.5 Prepare silicon coloured glass
1. Add a metal oxide to the glass mixture.
2. Heat the end of a glass rod to red heat.
Dip it into a powdered metallic oxide and heat until the oxide fuses into the glass.
Use salts to colour glass:
2.1 amethyst use manganese (IV) oxide,
2.2 green use black copper (II) oxide,
2.3 ruby use red copper (I) oxide,
2.4 white use tin oxide.
3. Mix a little silica with an equal quantity of calcium carbonate and about twice as much anhydrous sodium carbonate.
Grind them to a powder in a mortar with a pestle.
Make a loop in a platinum wire.
Heat the loop and place it in the mixture.
Reheat the wire with the adhering mixture until the mixture fuses.
Cool it and note what the bead looks like.
Hit it with a hammer and note if it is brittle?
Does it dissolve in water?
12.10.6 Prepare silicon glass in a furnace
Prepare glass in a crucible by heating a glass mixture in a furnace or over a Meker burner with a hot wide flame.
Glass mixture A: 17 g clean sand, 4.4 g sodium carbonate, 5.2 g disodium tetraborate (III)-10-water (borax).
Glass mixture B: 6 g clean sand, 2 g sodium carbonate, 1 g calcium carbonate.
12.10.7 Silicon compounds, glass
Silicon is a metalloid, because it has physical properties of metals and chemical properties of non-metals.
Silicon is a semiconductor.
Silicon does not exist free in nature, but occurs mainly as silicon (IV) oxide (SiO2), in silica sand, sandstone, clay, quartz and opal.
Silicates occur in most rocks and glass.
Portland cement is a mixture of calcium and aluminium silicates.
In the silicone oils and greases, the silicon atoms form polymers containing a chain of silicon and oxygen atoms with carbon and hydrogen atoms attached to the chain.
Silicones repel water.
Silica glass, an amorphous solid, is more like a "super cooled liquid" than a crystal, because , unlike crystalline substances, it does not have a sharp melting point.
However, some chemists say that glass is a disordered solid not a supercooled liquid.
Main types of glass:
1. Soda-lime glass
Round glass jars may contain: SiO2, CaO, K2O, MgO, TiO2
Flat window glass may contain: SiO2, CaO, Al2O3, TiO2
2. Borosilicate glass, Pyrex
3. Fused quartz glass may contain: SiO2, used in some camera lenses.
Experiments
Waterglass
1. Sodium carbonate is heated with sand to produce sodium silicate, the water-soluble waterglass used is an inorganic builder in detergents, for preserving eggs and for fireproofing materials.
Na2CO3(s) + SiO2 (s) --> Na2SiO3 + CO2(g).
sodium carbonate + silicon dioxide --> sodium silicate (waterglass), + carbon dioxide
2. When sodium oxide Na2O 15%, silicon dioxide SiO2 70%, and calcium oxide CaO 10% and other oxides are heated together with temperatures up to 1000oC, insoluble silica glass forms in which all the crystalline order of the added minerals has been lost.
In silica glass (soda lime silica glass, crown glass), each silicon atom is surrounded by four oxygen atoms as a tetrahedron and each of these is linked to other tetrahedra.
When ionic oxides are added in the glass making melt they get between the Si-O-Si bridges and weaken it, as shown in the transition glass temperatures: silica glass Tg = about 1200oC, Pyrex Tg = 550oC, window glass Tg = 550oC.
3. The glass in high quality wine glasses (lead crystal), contains lead, which gives the glass a ringing sound, higher refractive index and more brilliance.
Cobalt gives blue glass, chromium gives green glass, and copper gives red or blue-green glass.
Boron oxide, B2O3, gives shockproof borosilicate glass, "Pyrex", that is resistant to all chemicals except hydrofluoric acid, HF.
Flint glass, lead glass, has no colour unlike crown glass that has a slight green to yellow colour due to iron impurity.
4. The 2 distinct constituents of glass are as follows:
4.1 The network former, i.e. the non-metal as an oxide is usually silicon, but it can be boron, aluminium, or phosphorus.
4.2 The network modifiers, e.g. sodium, potassium, calcium, and magnesium.
Glass may crystallize over a period of many years and then become more brittle, but some glass has remained uncrystallized for 4000 years in Egypt.
Sodium sesquicarbonate Na2CO3.NaHCO3.H2O occurs as a mineral.
4.3 A famous urban legend has the opinion that glass in panes of old cathedrals and even early American buildings is thicker at the bottom, because the glass has "flowed" down over along periods.
The most likely explanation is that at a time when the thickness of panes of glass varied at that stage of technology, glaziers always inserted the thicker sides down.
5. Collect drinking glasses, including wine glasses, of roughly the same size.
Strike each glass and listen to the ring to identify the existence of modifiers in the glass.
12.10.8 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.
12.10.9 Dilatant compound, "Silly Putty", silicone, (toy product)
Silly putty, silicone, bouncing putty (Dow Corning 3179 dilatant compound), "Tricky Putty", (toy product)
Super ball, (toy product)
Slime balls, "Silly putty", silicone polymer to amuse children
The silicone polymer in silly putty, polyborosiloxane, have covalent bonds within the molecules, but hydrogen bonds between the molecules.
The hydrogen bonds are easily broken.
A silicone is chains of (OH-Si-O-Si-O-Si-OH), with two methyl groups, CH3 on each of the Si atoms.
However, in "Silly putty", boron atoms that can cross link weakly with oxygen atoms in other chains replace some of the silicon atoms.
Experiments
Prepare silly putty
1. Mix 55% Elmer's glue solution and 16% sodium borate in a 4: 1 ratio.
1a. Squeeze a 118 mL bottle of school glue into a bowl.
* Stir 120 mL of water into the glue.
* Add 1 teaspoon of borax or baking soda and stir until the borax has dissolved.
* Stir the borax water into the glue water until the glue starts to turn into gel.
* Pick up the glob of gel from the bowl and knead it with your fingers for 5 to 10 minutes.
* Store it in a plastic, resalable container, and knead it again when you take it out.
1b. Squeeze a 118 mL bottle of school glue into a bowl.
* Stir 120 mL of water into the glue.
* Slowly stir in liquid starch.
* Keep adding starch and stirring until the glue and starch comes together and forms a putty.
* If you use too much starch the silly putty will turn hard.
* Knead the putty together until the putty will clump together and become difficult to stir.
* Take the lump of putty out of the bowl and knead it until it turns firm.
* Store it in a plastic box with a tight-fitting lid.
1c. Pour 120 mL of dish soap into a bowl.
* Mix in 125 grams of cornstarch.
* Stir everything with a spoon, then use your hands.
* The crumbly mixture will will turn into a gel.
* Knead the mixture until the silly putty comes together.
* Store it in a plastic box with a lid.
2. Apply small amounts of stress are slowly to the silly putty only a few bonds are broken and the putty "flows" and stretches a great distance.
Apply larger amounts of stress quickly, many hydrogen bonds are broken, and the putty breaks or tears.
Roll it into a ball that you can bounce.
Press it onto a pencil drawing so that it lifts off the pencil marks so you can see the drawing on the surface of the silly putty.
12.10.10 Silicon tetrachloride with water
Silicon tetrachloride, a chlorosilane, is an epoxy resin hardener.
Highly toxic by all routes, with irritating fumes.
Stoppers of glass reagent bottles containing it may stick and the bottles may explode, so use very small containers.
It reacts vigorously with sodium and other reactive metals.
Silicon tetrachloride reacts vigorously with water to form insoluble silicon dioxide and an acidic hydrogen chloride solution.
The solution gives off fumes of hydrogen chloride and dissociates to form hydrochloric acid solution.
SiCl4 (s) + 2H2O (l) --> SiO2 (s) + 4HCl (aq)
HCl (aq) --> H+ (aq) + Cl- (aq).
12.10.11 Silicone rubber
Silicone rubbers, e.g. PVMQ. a low-temperature resistant rubber cross-linked with vinyl groups, and used for gaskets, because it is a temperature-resistant compound.
Polysiloxanes, [repeat unit: -Si(RR')O-, where R = organic group, e.g. methyl], are also used for aquarium seals and building seals.
12.11.0 Reactions of silver compounds
1. Grind solid silver nitrate with twice its volume of anhydrous sodium carbonate in a mortar.
Heat the mixture on charcoal in the reducing flame of a blowpipe.
A white bead of metallic silver forms that will not mark paper, but will dissolve in dilute nitric acid.
2. Add drops of concentrated hydrochloric acid to silver nitrate solution.
(Expensive!) Note the white precipitate of silver chloride.
Shake the mixture to coagulate the silver chloride, wash with water and leave to settle.
Ag+ + Cl- --> AgCl (s).
Pour off the water and divide the solid silver chloride into three parts.
Part (i): Expose it to light and it turns violet.
Part (ii) Add ammonium hydroxide and it dissolves.
Part (iii) Heat with concentrated hydrochloric acid and it dissolves.
3. Add drops of potassium chromate solution to silver nitrate solution.
Note the brick red precipitate of silver chromate that is soluble in both dilute nitric acid and sodium hydroxide.
2g+ + CrO42 --> Ag2CrO4 (s).
4. Add sodium phosphate solution to silver nitrate solution.
Note the yellow precipitate of silver phosphate.
3Ag+ + 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.
6. 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.
12.12.0 Reactions of strontium compounds
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.
12.19.5.0 CFCs, "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 following:
* Freon 11 CCl3F, trichlorofluoromethane).
* Tetrachloromethane, (carbon tetrachloride, CCl4, perchloromethane, dry cleaning fluid).
* 1,1,1-trichloroethane, ( C2H3Cl3, methyl chloroform, electrical equipment cleaner).
* 1,1,2,2-tetrachloroethane, (C2H2Cl4, refrigerant and solvent no longer used in USA).
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:.
* Freon 11, CCl3F, .trichlorofluoromethane CFC-11.
* Freon 12, CCl2F2, dichlorodifluoromethane, CFC-12, b.p. -30oC (most common refrigerant gas, solvent, used in fire extinguishers).
* Freon 21, CHCl2F, dichlorofluoromethane.
* Freon 114, CClF2CClF2, dichlorotetrafluoroethane.
* Freon 142, CH2CClF2, 1-dichloro1:1difluoroethane.
HCCl3 + 2HF --> HCF2Cl + 2HCl
chloroform + hydrogen fluoride --> chlorodifluoromethane + hydrogen chloride.
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 12.19.5.0, RODP = the relative ozone depletion potential (RODP).
CFC.
 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 .
12.1.0 Metals displace hydrogen from acids
1. Pour 5 cm of the acids in the table below into test-tubes.
Place a piece of metal foil in each test-tube.
Note the formation of hydrogen and compare the different rates at which the bubbles are formed.
Rate of formation of hydrogen gas with 3M hydrochloric acid and 3M sulfuric acid.
Metal
 3 M HCl
 3M (H2SO4
Magnesium
 Very rapid reaction
 Rapid reaction
Aluminium
 Slight reaction
 No reaction
Zinc
 Moderate reaction
 Slight reaction
Iron
 Very slight reaction
 Very slight reaction
Tin
 No reaction
 No reaction
Lead
 No reaction
 No reaction
Copper
 No reaction
 No reaction

2. Recover the zinc after the reaction has stopped.
Evaporate the solution to leave zinc sulfate crystals.
Dissolve the colourless zinc sulfate crystals in water and put two carbon electrodes (central poles of dry cell batteries) in the solution.
Connect the electrodes to a 6 V or 12 V DC supply.
Zinc forms rapidly on the cathode.