topic12C

Topic 12C Periodic table and the halogens
Updated 2008-08-19 R
Please send comments to: J.Elfick@uq.edu.au
See also: Interesting websites

Table of contents
12.18.0 Periodic table
12.18.1e Patterns in the periodic table
12.17.1.1 Different oxides and the periodic table
Halogens
Fluorine
Bromine
Chlorine
Iodine

Halogens
12.19.0 Group 17 of the periodic table, the halogens
12.19.1.0 Properties of halogens
12.19.2.0 Pass halide vapour over hot iron wire to form iron halides
12.19.3.0 Displace from its compounds less reactive halogens
12.19.4.0 Reactions of silver halides, AgCl_2,photography
12.19.5.0 CFCs, chlorofluorocarbons
3.32.1 Composition of the atmosphere and greenhouse gases

Iodine
12.19.6.0 Iodine extraction, I₂
12.19.6.1 Prepare hydrogen iodide, HI
12.19.6.2 Reactions of iodides, I^-
12.06.2.2 Reactions of potassium iodide solution with copper (II) sulfate solution, KI
12.19.6.3 Prepare iodic acid and potassium iodate, HIO₃, KIO₃
12.19.6.4 Compare the chloride, bromide and iodide of silver, AgCl, AgBr, AgI
12.19.6.5 Prepare tincture of iodine

Fluorine
12.19.7.0 Fluorine, F, hydrofluoric acid, HF
12.19.7.1 Prepare hydrogen fluoride, HF
12.19.7.2 Prepare silicon tetrafluoride, SiF₄

Chlorine3.40 Prepare chlorine, Cl₂
3.40.1 Tests for chlorine
3.40.2 Pass chlorine through water
12.19.8.1 Reactions of chlorides, Cl^-
12.19.8.2 Prepare chlorine, Cl_2, with concentrated hydrochloric acid and potassium manganate (VIII)
12.19.8.3 Prepare anhydrous iron (IlI) chloride, FeCl₃
12.19.8.4 Concentrated sulfuric acid on potassium chlorate, KClO₃
12.19.8.5 Prepare potassium perchlorate by fractional crystallization KClO₄12.19.8.6 Recover silver from silver chloride, AgCl₂

Bromine
12.19.9.1 Reactions of bromine, Br₂
12.19.9.2 Reactions of bromine water, Br₂
12.19.9.3 Prepare hydrogen bromide, HBr
12.19.9.4 Reactions of hydrogen bromide, HBr
12.19.9.5 Prepare potassium bromide, KBr
12.19.9.6 Reactions of bromides, Br^-

12.18.0e Periodic table
The periodic table is an orderly way to arrange the properties of the elements. The periodic table shows each element as a symbol with its atomic number atomic mass (whole number) electron notation and valence. The groups have group notation numbers, 1 to 18, as approved by the IUPAC (International Union of Pure and Applied Chemistry). The atomic number is shown above the symbol for each element. It is equal to the number of protons in the nucleus. The relative atomic mass is shown below the symbol for each element. It is the average of the values for the different isotopes of the element. The relative atomic mass of Carbon, C, is defined as Metallic properties are dominant towards the lower left corner and non-metallic properties are dominant towards the upper right corner.
Periods and groups for the first 20 elements: where G = Group (vertical), P = Period (horizontal)
In the full periodic table, Groups 3 to 12 contain the transition elements including: 4th Period: Mn, Fe, Co, Ni, Cu, Zn | 5th period: Mo, Ag, Cd, Sn | 6th period Pt, Au, Hg, Bi.

-	G 1	G 2	G 13	G 14	G 15	G 16	G 17	G 18
1st P	H	-	-	-	-	-	-	He
2nd P	Li	Be	B	C	N	O	F	Ne
3rd P	Na	Mg	Al	Si	P	S	Cl	Ar
4th P	K	Ca	-	-	-	-	-	-

12.18.1e Patterns in the periodic table
Use A4 size periodic charts for student use or a classroom size 2,000 X 1,500 mm periodic chart to find the following information and fill in a blank A4 size chart of the periodic table.
1. Periods: The elements have electrons in the same outer shell, i.e. the rows.
2. Groups: The elements have the same number of electrons in their outer shells, i.e. the columns. A group of elements has similar chemical properties.
3. Metals: Al and elements below and to left of Al and Sn, Pb (Sb, Bi) Po, but not H.
4. Non-metals: He, C, N, O, F, Ne, P, S, Cl, Ar, Se, Br, Kr, I, Xe, At, and Rn.
5. Metalloids: As, Ge, Si, Te (and Sb, Bi, B).
6. Elements that are gases at room temperature 27^oC: H₂, He, N₂, O₂, Ne, Cl₂, Ar, Kr, Xe and Rn. The elements that are liquids at room temperature are Hg and Br

12.19.0 Group 17 of the Periodic Table, the Halogens
See also 5.4.8: Iodine solution
The halogens are fluorine F, chlorine Cl, bromine Br, iodine I, and the rare element Astatine, At. Halogens form halide ions F^-, Cl^-, Br^-and I^-, are strong oxidizing agents, react with alkali metals to form salts and react with hydrogen in the decreasing order F-, then Cl^-, then Br^-, then I^-. Halogens react with most elements, are toxic and are very soluble in hydrocarbon solvents. The hydrogen halides are gases that dissolve in water to form acidic solutions that conduct electric current.

12.19.1.0 Properties of halogens
BE CAREFUL! CHLORINE AND BROMINE ARE HARMFUL WHEN INHALED OR WHEN THEY CONTACT THE SKIN!
Fluorine, F, is a yellow gas. Chlorine, Cl, is a green gas. Bromine, Br, is a red-brown liquid. Iodine, I, is a grey-black crystal. All halogens are slightly soluble in water to form weak acidic solutions that are bleaches. Chlorine is the most soluble and the strongest bleach. Test each solution with universal indicator. Universal indicator turns red then is bleached.
1. Compare the colours and states of the elements at room temperature. At room temperatures fluorine is a pale yellow gas, chlorine is a green-yellow gas, bromine is a red-brown liquid that gives a brown vapour and iodine is a grey-black solid.
2. Compare the colours, states and solubility in water of sodium fluoride, chloride, bromide and iodide. They are all soluble, white, crystalline solids.
3. Compare the activity of fluorine, chlorine, bromine and iodine by investigating which will displace another from their compounds. Prepare solutions of sodium fluoride, chloride, bromide and iodide in separate test-tubes. Add to each a solution of chlorine in water. Chlorine has no visible effect on solutions of sodium fluoride or sodium chloride but it turns sodium bromide solution yellow-brown and sodium iodide solution deep brown. This suggests that chlorine displaces bromine from sodium bromide and iodine from sodium iodide but that it does not displace fluorine from sodium fluoride, and so is more active than bromine and iodine but less active than fluorine.
4. Compare the colours and solubility of silver fluoride, chloride, bromide and iodide by adding silver nitrate solution to separate solutions of sodium fluoride, chloride, bromide and iodide. Silver fluoride is soluble. Silver chloride, bromide and iodide are insoluble. Silver chloride is white. Silver bromide is very pale yellow and silver iodide is a deeper yellow.

-	Fluorine	Chlorine	Bromine	Iodine
Element	yellow gas, irritating smell	yellow green, irritating smell	red liquid, irritating smell	black solid, irritating smell
Preparation	electrolysis of KHF₂	heat mixture of chloride, MnO₂ & conc. H₂SO₄	heat mixture of bromide, MnO₂, & conc. H₂SO₄	heat mixture of iodide, MnO₂, & conc. H₂SO₄
Activity	very reactive, combines with most metals & non-metals	very reactive, combines with most metals & many non-metals	very reactive, combines with most metals & non-metals	reactive, combines with most metals & few non-metals
Replacing action	replaces all halogens	replaces Br₂ & I₂ from bromides & iodides	replaces iodine from iodides	-
Oxidizing action	very powerful oxidizing agent	very powerful oxidizing agent	powerful oxidizing agent	weak oxidizing agent
Action with alkalis	forms fluoride	forms hypochlorite	forms hypobromite	forms hypoiodite
Halogen acid	fuming gas, weak acid	fuming gas, strong acid	fuming gas, strong acid	fuming gas, strong acid

12.19.2.0 Pass halide vapour over hot iron wire to form iron halides
See also 13.4.8: Burn steel wool in chlorine

12.19.3.0 Displace a less reactive halogen from its compounds
More reactive halogen displaces less reactive halogen from its compound.
1. Add iodine solution to colourless potassium bromide solution. No reaction because iodine is less active than bromine. Pass chlorine gas through colourless potassium bromide solution. The more active chlorine displaces the less active bromine and the solution turns orange.
2KBr(aq) + Cl₂(g)---> 2KCl(aq) + Br₂(aq)
2. Pass chlorine gas through potassium iodide solution. The more active chlorine displaces the less active iodine and the solution turns deep brown.

12.19.4.0 Reactions of silver halides, photography
Add silver nitrate solution to solutions of potassium chloride, potassium bromide and potassium iodide. Silver chloride (white), silver bromide (pale yellow) and silver iodide (deep yellow) are all sensitive to light and are used in photography.
Carl Wilhelm Scheele (1742 - 1786) exposed silver chloride beneath water to light. He added silver nitrate to precipitate new silver chloride then added a solution of ammonia to the blackened chloride to produce the solution of diammine silver ion. He also noticed that violet rays of the spectrum blackened the silver chloride much more than red rays.
Ag⁺ + Cl^- + light energy --> Ag⁺+ Cl + e^-, i.e. one electron lost from chlorine, oxidation of chlorine
Ag⁺ + e- --> Ag (metal), i.e. one electron gained by silver, reduction of silver, to form a dark image on film
AgCl + 2NH₃ --> Ag(NH₃)₂⁺ + Cl^-

12.19.5.0 CFCs, chlorofluorocarbons
See also 3.50: Ozone, O₃
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, e.g. Freon 11 (CCl₃F, trichlorofluoromethane)
and any other ozone depleting chemicals, Tetrachloromethane (carbon tetrachloride, CCl₄, perchloromethane, dry cleaning fluid), 1,1,1-trichloroethane (CH₃CCl₃, methyl chloroform, electrical equipment cleaner), 1,1,2,2-tetrachloroethane. 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. The relative ozone depletion potential (RODP) and half life in the atmosphere is as follows: CFC-11 RODP 1.00, Half life 75 years | CFC-12 RODP 0.86, Half life 112 years | CFC-13, Half life 90 years | CFC-22 RODP 0.05, Half life 20 years | CFC-113 RODP 0.80, - | CFC-114 RODP 0.60, - | 1,1,1-Trichloroethane RODP 0.15, Half life 6.5 years | Carbon tetrachloride RODP 1.11, Half life 50 years | Halon-1211 RODP 10.00, - | Halon-1301 RODP 10.00, - |
Methyl bromide is used to sterilize soil, kill rodents and plant pathogens. However, it depletes ozone in the atmosphere more than any CFC and there is international pressure to phase out its use. However seaweed, e.g. Asparagopsis taxiformis, produce great quantities of organohalides and produce the "smell" of the sea.

12.19.6.0 Iodine extraction
Set up a white background. retort stand and retort ring to fit three separating flasks. Dilute tincture of iodine to solution produce 200 mL of a deep golden solution. Mix 200 mL of the dilute iodine solution with 200 mL 0.2 M copper (II) sulfate to produce a green solution. Pour one third of the mixture into each of two 250 mL separating flasks. Keep a third flask for comparison. Add 100 mL of ether to one flask, insert a stopper and shake while venting the vapours to release any build up of pressure. Allow to stand. Add about 00 mL Freon or chloroform to the second separating flask. Insert a stopper, shake the flask and leave it to stand. The ether extraction produces two layers, a blue, lower, aqueous layer, and a yellow, upper, ether layer. With care in the dilution of the original iodine solution, and the selection of volumes of solution to be used, the final ether layer will appear the same colour as the original iodine solution. The Freon or chloroform extraction initially produces a purple colour, during the shaking, which should separate on standing to give an upper (aqueous) blue layer and a lower (Freon) pink layer. Iodine is a non-polar substance so it is more soluble in the non-polar Freon and the slightly polar ether, than in water, that is highly polar. Copper (II) ions are soluble in water but not in non-polar solvents.

12.19.6.1 Prepare hydrogen iodide
See diagram 12.19.6.1
1. Grind a 2 cc each of dry red phosphorus and iodine in a mortar and introduce into a boiling tube. Add four drops of water and fit the boiling tube.
P₄ + 6I₂ --> 4PI₃
PI₃ + 3H₂O --> H₃PO₃ + 3HI(g)
Heat the boiling tube to produce further quantities of hydrogen iodide.
2. Pass the gas into silver nitrate solution in a test-tube. Note the yellow precipitate of silver iodide.
3. Pass the gas into a test-tube containing drops of 880 ammonia. Note the white fumes of ammonium, iodide.
HI + NH₃ --> NH₄I
4. Pass the gas into a test-tube containing drops of concentrated nitric acid. The hydrogen iodide is easily oxidized to iodine and the nitric acid reduced to nitrogen dioxide.
2HNO₃ + 2HI --> 2H₂O + 2NO₂ + I₂
5. Pass the gas for some time into a test-tube containing concentrated sulfuric acid. The acid is reduced to sulfur dioxide, hydrogen sulfide or sulfur, showing that hydrogen iodide is a powerful reducing agent.
6. Heat the delivery tube with a Bunsen burner. Note the violet vapours of iodine.
2HI <--> H₂ + I₂

12.19.6.2 Reactions of iodides
Use one crystal the size of a match head or 1 mL of a 10% solution for each reaction.
1. Grind the iodide with a small quantity of manganese dioxide and add 1 mL of concentrated sulfuric acid to the mixture in a test-tube. Heat gently and observe the violet vapours of iodine.
MnO₂ + 2KI + 2H₂SO₄ --> MnSO₄ + K₂SO₄ + 2H₂O + I₂
With concentrated sulfuric acid alone iodine is also obtained on heating because hydrogen iodide is a powerful reducing agent.
2. Add some silver nitrate solution to potassium iodide solution. Note the yellow precipitate of silver iodide that is insoluble in both dilute nitric acid and ammonium hydroxide solution.
Ag⁺ + I- --> AgI(s)
3. Add drops of chlorine water to potassium iodide solution. Note that iodine is given off. The iodine turns the solution brown and some black crystals may be seen at the surface.
Cl₂ + 2I- --> I₂ + 2Cl^-
4. Add drops of lead acetate solution to potassium iodide solution. Note the yellow precipitate of lead iodide.
Pb²⁺ + 2I- --> PbI₂(s)

12.19.6.3 Prepare iodic acid, HIO₃, and potassium iodate
1. To prepare iodic acid, use 5 g of iodine in a retort and add a measured 40 mL of fuming nitric acid. Heat the retort on a sand tray, keeping the temperature high enough to promote action. Collect any nitric acid which distils over and return it to the retort. When the iodine has all been oxidized to white crystals of iodic acid, pour the contents of the retort into an evaporating basin and heat almost to dryness on a water bath. Collect the crystals and dry between filter paper.
I₂ + 10HNO₃ --> 2HIO₃ + 10NO₂ + 4H₂O
Heat some of the crystals in a dry test-tube, gently then strongly. Note the formation of moisture to leave iodine pentoxide.
2HIO₃ --> I₂O₅ + H₂O
This is followed by decomposition to iodine and oxygen. (Test with glowing splint.)
2I₂O₅ --> 2I₂ + 5O₂2. To prepare potassium iodate, put 2 g of potassium hydroxide into a test-tube and add 4 cm of water. When dissolved, slowly add 4.5 g of iodine to the heat solution. Pour the solution into a watch glass and let cool
6OH^- + 3I₂ --> IO₃^- + 5I^- + 3H₂O
Pour off the solution from the crystals, wash the latter with some water and dry them on a filter paper. Heat crystals in a dry test-tube and show oxygen forms.
3. Prepare potassium bromate by a similar experiment using 30 drops (1 mL) of bromine in place of the iodine.

12.19.6.4 Compare the chloride, bromide and iodide of silver
1. Add silver nitrate solution to 2 cm of potassium chloride solution, potassium bromide solution and potassium iodide.
Ag⁺ + Cl^- --> AgCl(s) white precipitate
Ag⁺ + Br^- --> AgBr(s) slightly yellow precipitate
Ag⁺ + I^- --> AgI(s) yellow precipitate
Divide each solution into two parts, Part 1 and Part 2:
Part 1. Add dilute nitric acid.
Part 2. Add drops of dilute ammonium hydroxide.
All three precipitates are insoluble in nitric acid. Silver chloride is soluble in ammonia solution forming the soluble complex ion [Ag(NH₃)₂]⁺. Silver bromide is slightly soluble in ammonia solution. Silver iodide is insoluble in ammonia solution.
AgCl + 2NH₃ --> [Ag(NH₃)₂]⁺ + Cl^-
Silver fluoride is soluble in water, so potassium fluoride solution gives no precipitate with silver nitrate.
2. Repeat the experiment or divide the original precipitate into three parts instead of two and show that all three precipitates are soluble in sodium thiosulfate solution. Silver chloride dissolves sodium thiosulfate to produce the sodium silver thiosulfate ion, NaAgS₂O₃.
AgCl + Na₂SO₃ --> NaAgS₂O₃ + NaCl

12.19.6.5 Prepare tincture of iodine
Add 70 g of iodine, I2 and 50 g of potassium iodide, KI, water, dilute to 1 litre with addition of ethanol.

12.19.7.0 Fluorine, hydrofluoric acid
Fluorine, F, is a poisonous green-yellow gas. Fluorine is the most reactive corrosive and electronegative of all elements. It never occurs as a free gas and is not used in school laboratories. It occurs in the minerals fluorite, CaF2, Fluorine combines with carbon to form inert polymers, e.g. Teflon coated frying pans. In some countries, sodium fluoride is added to drinking water to improve the hardness of tooth enamel apatite of children's teeth. However, in some countries, this is not allowed because some people believe that sodium fluoride is too reactive to be put into drinking water and it may cause discoloration of teeth.
Dissociation of hydrofluoric acid
HF + H₂O <--> H₃O⁺ + F-

12.19.7.1 Prepare hydrogen fluoride, HF
Fluorine combines with nearly all known elements. Use a fume cupboard.
Coat a microscope slide with wax and remove part of the wax by writing on the slide with a pin. Pour concentrated sulfuric acid over powdered calcium fluoride in the bottom of a lead basin. Put the microscope slide face downwards over the basin and heat gently. Blowing across an ammonia bottle in the directions of the gas. Put a piece of damp blue litmus paper in the gas. The action of silver nitrate on fluorides in solution is not typical as silver fluoride is soluble in water. Note the steam-like fumes hydrogen fluoride. After three minutes, remove the wax from the slide and note the writing etched on the glass because of the formation of silicon fluoride.
CaF₂ + H₂SO₄ --> CaSO₄ + 2HF
SiO₂ + 4HF --> SiF4 + 2H₂O (The SiO₂ comes from the glass in the microscope slide.)

12.19.7.2 Prepare silicon tetrafluoride
See diagram 12.19.7.2
1. All the apparatus must be dry. Mix 5 cc of fine dry sand and 5 cc of powdered calcium fluoride and put the mixture into a 250 mL flask with a 2-hole stopper fitted with a funnel and delivery tube. Put other end of the delivery tube into a gas jar, to keep the delivery tube dry. Pour concentrated sulfuric acid on to the mixture in the flask, and shake to moisten the whole mass. Pour water to a depth of 6 cm and then heat the contents of the flask. The silicon tetrafluoride, which is a gas, passes into the water and hydrolyses to form a white precipitate of hydrated silica and a solution of hydrofluosilicic acid, H₂SiF₆.
3SiF₄ + 2H₂O --> SiO₂(s) + 2H₂SiF₆
The precipitate may be regarded as silicic acid (with a formula H₂SiO₃ or H₄SiO₄ or as hydrated silica (SiO₂.xH₂O). Some of the water may be combined and some occluded.
2. To obtain a specimen of pure silica from sand, filter off the solution with suspended silica obtained above, and wash the hydrated silica well with four washings of hot distilled water. Transfer the silica to a crucible and heat to redness. Leave to cool. The product is pure silica.

12.19.8.1 Reactions of chlorides
1. Add 1 mL of concentrated sulfuric acid to 2 cc of sodium chloride solution. Hydrogen chloride is given off.
NaCl + H₂SO₄ --> NaHSO₄ + HCl(g)
2. Grind 2 cc of sodium chloride with twice its volume of manganese dioxide and transfer the mixture to a boiling tube. Add 3 mL of concentrated sulfuric acid and heat. Chlorine forms and some hydrogen chloride forms. Chlorine bleaches wet litmus paper.
2NaCl + 2H₂SO₄ + MnO₂ --> Na₂SO₄ + MnSO₄ + 2H₂O + Cl₂(g)
3. Prepare chromyl chloride, CrO₂Cl₂
Heat a dry test-tube in the Bunsen flame to soften the glass a third of the distance from the open end. Draw out the glass to reduce the diameter of the test-tube to a 0.5 cm at the heated part and at the same time bend the open end slightly downwards. The apparatus will then serve as a small retort. When cold introduce into the test-tube a mixture of not more than 2 cc of finely ground potassium dichromate and half that amount of sodium chloride. Add just enough concentrated sulfuric acid to cover the mixture. Grasp the test-tube in one holder and in another hold a dry test-tube to act as a receiver. Heat the mixture gently and collect drops of the red-brown liquid, chromyl chloride.
K₂Cr₂O₇ + 4NaCl + 3H₂SO₄ --> K₂SO₄ + 2Na₂SO₄ + 2CrO₂Cl₂ + 3H₂O
Add drops of water to the compound. Test the hydrogen chloride given off with litmus paper and with silver nitrate solution on a glass rod. The yellow solution contains chromic acid. Add sodium hydroxide solution until neutral, then acidify with acetic acid and add lead acetate solution. Note the yellow precipitate that shows the presence of chromate ion.
CrO₂Cl₂ + 2H₂O --> H₂CrO₄ + 2HCl

12.19.8.2 Prepare chlorine with concentrated hydrochloric acid and potassium permanganate
This experiment may not be allowed in some school systems. Use a fume cupboard.
See diagram 12.19.8.2
1. Put 5 cc of potassium permanganate into a flask. Fill the dropping funnel with concentrated hydrochloric acid and allow the acid to run on to the permanganate to produce chlorine.
2MnO₄^- + 16H⁺ + 10Cl^- --> 2Mn²⁺ + 8H₂O + 5Cl₂(g)
2. Pass the gas through water to remove hydrogen chloride and pass the gas through concentrated sulfuric acid to dry it. Collect the gas by downward displacement of air.

12.19.8.3 Prepare anhydrous iron (IlI) chloride
See diagram 12.19.8.3
Use a fume cupboard. This experiment may not be allowed in some school systems.
1. Anhydrous iron (IlI) chloride and some other chloride, e.g. iron (II) chloride, magnesium chloride, zinc chloride, aluminium chloride, and antimony chloride, cannot be prepared by evaporating a solution of the salt to dryness because hydrolysis occurs and the final product is iron (IlI) oxide or any of the other oxides.
Wind 50 cm of thin iron wire around a pencil and put the wire in a combustion tube. Pass dry chlorine through the combustion tube over for a minute to displace the air and then heat the tube with a Bunsen burner until the iron wire commences to burn. After removal of the Bunsen burner flame the wire will continue to burn if the supply of chlorine is sufficient. Most of the iron (IlI) chloride condenses as a mass of black crystals in a cooler part of the combustion tube.
2Fe + 3Cl₂(g) --> 2FeCl₃(s)
2. Repeat the experiment using dry hydrogen chloride is used in place of chlorine. The small colourless scales of iron (II) chloride produced are much less volatile and often stick to the iron.
Fe + 2HCl --> FeCl₂ + H₂(g)

12.19.8.4 Concentrated sulfuric acid on potassium chlorate
This experiment may not be allowed in some education systems.
Drop a crystal of potassium chlorate the size of half a small split pea into a clean dry test-tube and clamp in a nearly horizontal position. Make sure that the mouth of the test-tube points away from you in a safe direction. Drop two drops of concentrated sulfuric acid into the mouth of the test-tube. Adjust the slant of the test-tube so that the acid runs slowly down onto the potassium chlorate. The yellow gas given off is chlorine dioxide, ClO₂.
Install a safety screen between you and the equipment and slowly heat the test-tube while holding it at arm's length. A violent reaction occurs as the chlorine dioxide decomposes.
3KClO₃ + 2H₂SO₄ --> KClO₄ + 2KHSO₄+ 2ClO₂ + H₂O

12.19.8.5 Prepare potassium perchlorate by fractional crystallization
See diagram 12.19.8.5
Half fill a crucible with potassium chlorate. Fit the crucible firmly in a pipe- clay triangle. Heat gently until the potassium chlorate melts then stir the liquid it becomes pasty while supplying heat to keep the mass molten. Leave to cool, add an equal volume of water and heat gently until all the potassium chlorate has dissolved. Pour the solution on to a watch glass and let cool. The crystals which appear are almost pure potassium perchlorate, KClO₄, that can be purified further by dissolving in hot water and crystallizing again.
4KClO₃ --> 3KClO₄ + KCl

12.19.8.6 Recover silver from silver chloride
1. Wash the residues with water, dry them and mix with twice the volume of a mixture of anhydrous sodium and potassium carbonates. Transfer the mixture to a crucible and heat strongly in a furnace, then leave to cool. Note a button of silver remaining in the bottom of the crucible.
2. Transfer the residues after washing to an evaporating basin and add sodium hydroxide solution and glucose and heat the mixture. When a portion of the solid dissolves completely in dilute nitric acid, pour off the liquid from the grey silver which remains in a finely divided condition.

12.19.9.1 Reactions of bromine
Be careful! Liquid bromine can cause sores if in contact with the skin. Also, bromine vapour is painful to the eyes.
1. Let drops of bromine fall into a test-tube and cover the mouth by 2 / 3 with a stopper. The bromine evaporates and fills the test-tube. Fix the stopper firmly in the mouth of the test-tube. Invert a test-tube of hydrogen over a test-tube of bromine and let them mix. Apply a flame and note the weak explosion
H₂ + Br₂ --> 2HBr
2. Invert a test-tube of hydrogen sulfide over a test-tube of bromine. Note the sulfur precipitate. Misty fumes of hydrogen bromide replace the colour of bromine.
Br₂ + H₂S --> 2HBr + S(s)
3. Dip a filter paper in an alcoholic solution of fluorescein and let dry. Put it in a gas jar of bromine vapour, when the paper turns red because of the formation of eosin.

12.19.9.2 Reactions of bromine water
Bromine dissolves to a slightly in water forming a 4% red solution. A red vapour remains above the saturated solution.
1. Add iron filings to 3 cm of bromine water and shake the mixture. Note the pale green solution of iron (II) bromide if iron is in excess, or the yellow solution of iron (IlI) bromide if bromine is in excess. Tests for the presence of iron (II) or iron (IlI) iron by adding sodium hydroxide solution to give a black precipitate of Fe₃O₄.
Fe + Br₂ --> FeBr₂
Fe + Br₂ --> Fe²⁺ + 2Br^-
2Fe + 3Br₂ --> 2FeBr₃
2Fe + 3Br₂ --> 2Fe³⁺ + 6Br^-
2. Hold a piece of blue litmus paper in the vapour above bromine water. The litmus paper turns red and becomes bleached.
3. Add drops of bromine water to 3 cm of sulfurous acid. Test the solution for sulfate by adding dilute hydrochloric acid followed by barium chloride.
SO₃^2- + Br₂ + H₂O --> SO₄^2- + 2Br^- + 2H⁺
4. Add drops of sodium hydroxide solution to 3 cm of bromine water until the colour disappears. The remaining solution contains the hypobromite ion, BrO^-.
2OH^- + Br₂ --> BrO^- + Br^- + H₂O
The hypobromite solution can precipitate manganese dioxide from a solution of manganese sulfate and precipitate lead dioxide from a solution of lead nitrate.
5. Add 2 cc of red phosphorus to 3 cm of bromine water. Shake the mixture and leave to stand. Note that the colour of the bromine water disappears. The bromine and phosphorus combine and the resulting bromide of phosphorus decomposes to give phosphorous or phosphoric acids.
P₄ + 6Br₂ --> 4PBr₂
P4 + 10Br₂ --> 4PBr₅
2PBr₃ + 6H₂O --> 2H₃PO₃ + 6HBr
PBr₂ + 4H₂O --> H₃PO₄ + 5HBr
Add bromine water to potassium iodide solution. Iodine is displaced.
2I^- + Br₂ --> I₂(s) + 2Br^-

12.19.9.3 Prepare hydrogen bromide
See diagram 12.19.9.3
1. Add drops totalling 2 mL of acid to 1 mL of water in a boiling tube. Add 2 cc of potassium bromide and heat gently. Be careful!
KBr + H₂SO₄ --> KHSO₄ + HBr(g)
2. Put a paste of 5 g of red phosphorus with water and sand into the flask. The sand is to moderate the action. Slowly let drops of bromine fall from the tap funnel. The first few drops react with a flash of light. Pass the gases through a U-tube containing beads smeared with damp red phosphorus to remove the bromine volatilized by the heat of the reaction. Collect the hydrogen bromide by displacement of air or by passing it through an inverted funnel over water.
3. Show the action of heat on hydrogen bromide by filling a boiling tube with hydrogen bromide, inserting a loose stopper, and heating strongly with a Bunsen burner. Hold a piece of white paper behind the boiling tube as soon as decomposition starts.
2HBr <-> H₂ + Br₂

12.19.9.4 Reactions of hydrogen bromide
Test the misty fumes of hydrogen bromide by slowly lowering a drop of the following reagents on the end of a glass rod into the gas.
1. Silver nitrate solution: A pale yellow precipitate of silver bromide forms.
Ag⁺ + Br^- --> AgBr(s)
2. 880 ammonia: Ammonium bromide fumes form.
NH₃ + HBr --> NH₄Br
3. Litmus solution: Litmus turns red.
4. Chlorine water: Yellow coloration because of bromine.
2Br^- + Cl₂ --> Br₂ + 2Cl^-
5. A drop of water: This may be removed and tested by dipping the rod into drops of silver nitrate solution in a test-tube. The positive reaction shows the high solubility of the gas in water.
6. Concentrated nitric acid: The hydrogen bromide is rapidly oxidized to bromine.
2HNO₃ + 2HBr --> 2H₂O + 2NO₂ + Br₂
7. Attach a stopper with a delivery tube bent at right angles and heat strongly. Hold a piece of white paper behind the test-tube. Strong heat decomposes the hydrogen bromide into bromine and hydrogen. If the bromine is not easily visible, put in the bromine vapour a filter paper dipped in an alcoholic solution of fluorescein and let dry. The filter paper turns red because of the formation of eosin.

12.19.9.5 Prepare potassium bromide, KBr
1. Use 100 mL of distilled water in the flask. Measure 5 mL of bromine in a measuring cylinder which already contains 3 mL of water. Pour the bromine and water into the flask (and wash out the cylinder at once). Weigh 8 g of iron filings and add to the solution in portions of 0-5 g, shaking well on each addition. If this operation of adding the iron is hurried, much heat is generated and some iron (IlI) bromide forms and persists throughout the preparation. Heat the flask on a water bath for ten minutes and filter quickly.
Fe + Br₂ --> FeBr₂
2. Prepare a solution of 20 g of potassium carbonate in 50 mL of water. Add this solution to the green iron (II) bromide solution, mix and heat on a water bath for ten minutes. The white later green precipitate is iron (II) carbonate.
FeK₂ + K₂CO₃ --> 2KBr + FeCO₃Filter quickly and evaporate the colourless solution to crystallization. Examine a drop of solution under a microscope for cubic crystals of the bromide.

12.19.9.6 Reactions of bromides
Use one crystal the size of a match head or 1 mL of 10% potassium bromide for each experiment, chlorine water, carbon tetrachloride.
1. Grind the bromide with a small quantity of manganese dioxide, add 1 mL of concentrated sulfuric acid to the mixture in a test-tube and heat gently. The red vapour of bromine may condense to small drops of liquid bromine on the sides of the test-tube.
MnO₂ + 2KBr + 2H₂SO₄ --> MnSO₄ + K₂SO₄ + 2H₂O + Br₂
2) Add drops of silver nitrate solution to potassium bromide solution. Note the pale yellow precipitate of silver bromide that is insoluble in dilute nitric acid but dissolves in excess ammonium hydroxide, i.e. it is sparingly soluble.
Ag⁺ + Br^- --> AgBr(s)
3. Add drops of chlorine water to a potassium bromide solution. Bromine is liberated which turns the solution light brown or red.
Cl₂ + 2Br^- --> Br₂ + 2Cl^-

12.18.1 Preparation of forms of sulfur
Almost fill a dry test-tube with sulfur powder and heat slowly to boiling, using a safety holder. Pour the boiling sulfur into a beaker of water. Immerse any floating sulfur with a stirring rod. Remove and examine the plastic sulfur. Note the gradual loss of elasticity as the plastic sulfur changes to rhombic sulfur.

12.18.2 Prepare sulfides
1. Use an ignition tube with 3 cm of powdered sulfur and heat until melted. Hook a strip of copper over the rim of the ignition tube so that its lower edge is just above the surface of the sulfur. Heat to boil the sulfur and note the glow as copper sulfide forms on the copper.
2. Mix equal parts. of iron filings and sulfur. Heat the mixture until a reaction starts. Note the glow of the mixture as iron (II) sulfide forms.
3. Pass hydrogen sulfide into copper (II) sulfate solution. Filter off the precipitated copper (II) sulfide.Cu²⁺ + S^2- --> CuS(s)4. Prepare sulfides of iron, cobalt and nickel. Prepare solutions of iron (II) cobalt and nickel salts. Pass hydrogen sulfide into each solution Note a slight precipitate of dark iron (II) sulfide but no precipitate with the cobalt or nickel salts. Add ammonia solution, NH3(aq) ("ammonium hydroxide") and pass more hydrogen sulfide through the three solutions. All three solutions form a black precipitate of the metallic sulfide, however only iron (II) sulfide dissolves in dilute hydrochloric acid. The sulfides of cobalt and nickel dissolve in concentrated hydrochloric acid in the presence of potassium chlorate or in aqua regia. Transfer the sulfides to evaporating basins, add concentrated hydrochloric acid and a crystal of potassium chlorate. Heat until the crystals dissolve. The cobalt salt becomes is pink in solution. The nickel salt becomes green-yellow.

12.18.3a Prepare sulfur monochloride, S₂Cl₂
This experiment may not be allowed in some school systems.
Put 10 cc of sulfur in a distilling flask. Use a 1-hole stopper fitted with a delivery tube to reach the level of the sulfur. Connect the delivery tube to a supply of dry chlorine. Heat the sulfur on a gauze and pass in chlorine. Stand the flask in a water bath of cold water. Collect the liquid product of sulfur monochloride, S₂Cl₂, in a dry test-tube. Heat drops of the product in water and test for sulfur dioxide and hydrochloric acid. Note the deposit of sulfur.
2S₂Cl₂ + 2H₂O -->- 4HCl + 3S(s) + SO₂(g)

12.18.3b Prepare thionyl chloride, SOCl₂
This experiment may not be allowed in some school systems.
See diagram 12.18.3b
1. Use 10 cc of phosphorus pentachloride in a dry distilling flask attached to a sloping condenser. Fit the flask with a stopper and delivery tube that reaches deep into the flask. Stand the flask in a cold water bath. Pass sulfur dioxide into the phosphorus pentachloride until it has completely liquefied. Heat the water bath and collect the distillate of thionyl chloride, SOCl₂. The liquid remaining in the flask is phosphorus oxychloride.
PCl₅ + SO₂ --> POC1₃+ SOCl₂
Add drops of the thionyl chloride to water and test for sulfurous acid and hydrochloric acid.
SOCl₂ + 2H₂O --> H₂SO₃ + 2HCl
2. Cut a cube of camphor into pieces and place in the dry distilling flask. Pass sulfur dioxide through the flask until the camphor liquefies. Disconnect the source of sulfur dioxide and pass dry chlorine through the flask until it is no longer absorbed. Heat the. water bath and collect the distillate of sulfuryl chloride, SO₂C1_2.
SO₂ + C1₂ --> SO₂C1₂
Add drops of the distillate to water and show that the resulting solution contains sulfuric acid and hydrochloric acids.
SO₂Cl₂ + 2H₂O --> H₂SO₄+ 2HCl

12.18.4 Properties of sulfur dioxide and sulfites
1. Heat sulfur in an evaporating basin and test the sulfur dioxide formed by drops reagents on a glass rod. Alternatively, pass sulfur dioxide into 6 cm of water. Show that the solution is acid, potassium permanganate loses its colour, reduces potassium dichromate, and forms a deposit of sulfur with hydrogen sulfide.
2. Prepare sodium sulfite and sodium bisulfite in solution. Saturate 10 mL of sodium hydroxide solution with sulfur dioxide to form sodium bisulfite, NaHSO₃. Add 10 mL of sodium hydroxide solution to form sodium sulfite, Na₂SO₃.
NaOH + SO₂(g) --> NaHSO₃
NaHSO₃ + NaOH --> Na₂SO₃ + H₂O
3. Hold 2 cm of magnesium ribbon in a pair tongs and heat until it ignites, then hold the burning magnesium in sulfur dioxide. Sulfur dioxide decomposes into sulfur and oxygen and magnesium oxide forms.
2Mg + SO₂(g) --> 2MgO(s) + S(s)
See also 13.13.8.
4. Pass air through a hot solution of sodium sulfite, Na₂SO₃. Test the solution for sulfate.
2SO₃^2- + O₂(g) --> 2SO₄^2-
5. Dissolve 5 cc of sodium sulfite crystals in 50 mL of water. Add 2 cc of crushed roll sulfur and boil for an hour. Transfer the mixture to an evaporating basin and heat to a small volume. Test the concentrated solution for sodium thiosulfate, Na₂S₂O_3:1. By addition of iodine solution 2. by addition of an acid.
Na₂SO₃ + (O) --> Na₂SO₄
Na₂SO₃ + (S) --> Na₂S₂O₃
6. Add dilute hydrochloric acid to crystals of sodium sulfite and heat. Sulfur dioxide forms. Tests for sulfur dioxide: 1. notice its smell 2. dilute solution of potassium permanganate loses its colour.
SO₃^2- + 2H⁺ --> SO₃(g) + H₂O
7. Add barium chloride solution to a freshly made solution of sodium sulfite. A white precipitate of barium sulfite forms. Unlike barium sulfate, barium sulfite is soluble in dilute hydrochloric acid.
SO₃^2- + Ba²⁺ --> BaSO₃(s)
8. Add drops of iodine in potassium iodide solution (tincture of iodine) to sodium sulfite solution. The iodine loses its colour. Test the final solution for sulfate ion.
SO₃^2- + H₂O + I₂ --> SO₄^2- + 2I^- + 2H⁺

12.18.5.1 Sulfuric acid acts as a dehydrating agent, removing water, or the elements of water, from another substance.
1. Add 2 cm of concentrated sulfuric acid to 1 cm of copper (II) sulfate crystals. After ten minutes, note the colour change from blue copper (II) sulfate crystals to white anhydrous copper (II) sulfate.
CuSO₄.5H₂O + (H₂SO₄) --> CuSO₄ + (H₂SO₄.5H₂O)
12.18.5.1a Dehydration of sugar by sulfuric acid
2. Heat a mixture of 0.5 cm of sucrose and 1.0 cm of concentrated sulfuric acid gently for 2 seconds and then leave to stand. Note the vigorous reaction and the colour change from white sugar to black carbon.
C₁₂H₂₂O₁₁ + (H₂SO₄ catalyst) --> 12C + 11H₂O
This experiment causes great amusement to children if the sugar is in the shape of a volcano in a deep beaker and the sulfuric acid is poured into the "crater " of the "volcano".

12.18.5.2 Sulfuric acid acts as a displacer of acids from their salts, sulfuric acid being much less volatile than most other acids.
1. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium chloride. Test the fuming gas with silver nitrate solution on a glass rod to form white silver chloride with hydrogen chloride. The less volatile sulfuric acid displaces the hydrogen chloride.
NaCl + H₂SO₄ --> HCl + NaHSO₄
2. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium acetate. Note the smell of the displaced the acetic acid.
3. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium formate. Note the displacement of formic acid followed by dehydration.

12.18.5.3 Sulfuric acid acts as a displacer of acids from their salts, sulfuric acid being much less volatile than most other acids. Also, sulfuric acid acts as an oxidizing agent.
1. Add an equal volume of concentrated sulfuric acid to 0.5 cm of potassium bromide. A fuming gas first forms then a brown gas. Hydrogen bromide is displaced then partially oxidized to bromine. Hydrogen bromide turns silver nitrate on a glass rod to pale yellow silver bromide. Potassium permanganate solution on a glass rod decolorizes the sulfur dioxide, formed by reduction of sulfuric acid.
KBr + H₂SO₄ --> HBr + KHSO₄
2HBr + H₂SO₄ --> Br₂(g) + 2H₂O + SO₂
4H⁺ + 2Br- + SO₄^2- --> Br₂(g) + 2H₂O + SO₂(g)
2. Add an equal volume of concentrated sulfuric acid to 0.5 cm of potassium iodide. A fuming gas first forms then a brown gas. Hydrogen iodide is displaced then oxidized to iodine. Note the greater extent of oxidation compared with the previous experiment. Much of the hydrogen iodide is oxidized to iodine. Heat the test-tube and note the violet vapour of iodine.
4H⁺ + 2I^- + SO₄^2- --> I₂+ 2H₂O + SO₂

12.18.5.4 Reactions of dilute sulfuric acid as an acid
1. Add 2 cm of dilute sulfuric acid to 1 cc of zinc powder. Close the test-tube with the thumb until enough hydrogen forms to give a mild explosion when the mouth of the test-tube is held in a flame.
2H⁺ + Zn(s) --> Zn²⁺ + H_2(g)
2. Add 2 cm of sodium carbonate to 1 cc of zinc powder. Tests for carbon dioxide by passing the gas given off to pass into limewater that turns milky because of the fine precipitate of calcium carbonate.
2H⁺ + CO₃^2- --> H₂O + CO₂(g)
Ca(OH)₂ + CO₂(g) --> CaCO₃(s) + H₂O

12.18.5.5 Reactions of dilute sulfuric acid as a sulfate
Add an equal volume of barium chloride solution to 3 cm of dilute sulfuric acid. Note the white precipitate of barium sulfate. Allow the precipitate to settle, filter, wash and leave to dry.
SO₄^2- + Ba²⁺ --> BaSO₄(s)

12.18.6 Preparation of sodium thiosulfate crystals, Na₂S₂O₃.5H₂O, "hypo"
Put 150 mL of water, 30 g of sodium sulfite and 15 g of crushed sulfur in a 250 mL round bottom flask and fit it with a reflux condenser. Heat the flask on a gauze for three hours. Filter the solution and evaporate to 30 mL. Leave to cool and crystallize.
SO₃^2- + S(s) --> S₂O₃^2- (thiosulfite ion = S₂O₃^2-)

12.18.6a Reactions of sodium thiosulfate, Na₂S₂O₃.5H₂O, "hypo"
1. Heat crystals of sodium thiosulfate in a dry test-tube until the test-tube begins to melt. Note water and sulfur as products of the reaction. Leave the mixture to cool. Add dilute hydrochloric acid to the residue and note that hydrogen sulfide is given off.
4Na₂S₂O₃ --> 3Na₂SO₄+ Na₂S₅
Na₂S₅ -->Na₂S + 4S(s)
2. Add iodine solution to sodium thiosulfate solution. The iodine loses its colour and thiosulfate ion is converted to tetrathionate ion.
2S₂O₃^2- + I₂ --> S₄O₆^2- + 2I^- (tetrathionate ion = S₄O₆^2-)
3. Add chlorine water or bromine water in excess to sodium thiosulfate solution and test with barium chloride solution. The products are sodium sulfate and sulfur that may be further oxidized to sulfuric acid.
S₂O₃^2- + Cl₂ + H₂O --> SO₄^2- + S(s) + 2H⁺ + 2Cl^-
4. Add concentrated hydrochloric acid to sodium thiosulfate solution. Sulfur precipitates and sulfur dioxide forms.
S₂O₃^2- + 2H⁺ --> SO₂(g) + H₂O + S(s)

12.18.7 Reactions of sulfamic acid, NH₂.SO₂OH.
1. Dissolve 1 cc of sulfamic acid in 2 cm of water. Note the high solubility of the acid. Tests for sulfate ion by adding dilute hydrochloric acid and drops of barium chloride solution. At first there is little action but leave to stand and white suspension of barium sulfate forms. Boil the mixture and the barium sulfate becomes more apparent as the sulfamic acid hydrolyses.
NH₂.SO₂.OH + H₂O --> NH₄HSO₄
2. Dissolve 1 cc of sulfamic acid in 2 cm of water. Dissolve 1 cc of sodium nitrite in 2 cm of water. Mix the solutions. Note the vigorous effervescence as nitric oxide, nitrogen dioxide and nitrogen are given off. Sulfamic acid is a strong fully ionized acid that reacts with the nitrites to give oxides of nitrogen and its -NH₂ group. sulfamic acid also reacts with the nitrite to give nitrogen.
2H⁺ + 2HNO₂ + 2e^- --> 2H₂O + 2NO
2NO + O₂ --> 2NO₂
NH₂.SO₂.O^- + H⁺ + NO₂^- --> N₂(g) + HSO₄^- + H₂O
Tests for sulfate ion by adding dilute hydrochloric acid and drops of barium chloride solution. A white suspension of barium sulfate forms.
3. Add sodium hydroxide solution to 1 mL of sulfamic acid to a depth of 2 cm for an excess of sodium hydroxide. Heat the solution and test the gas formed for ammonia with damp red litmus paper.
NH₂.SO₂.O^- + 2OH^- --> NH₃(g) + SO₄^2- + H₂O
4. Heat 1 cc of sulfamic acid in a dry test-tube. Tests for sulfur dioxide with a spot of potassium permanganate on a filter paper. Test for sulfur trioxide by allowing the white fumes to flow into a test-tube containing barium chloride solution acidified with hydrochloric acid. Note the crystalline sublimate and dissolve the crystals in 2 cm of sodium hydroxide solution. Heat the solution then test for ammonia with damp red litmus paper. Acidify the remaining solution with hydrochloric acid and add barium chloride solution to test for sulfate ion.

12.20.1 Reactions of tin and its compounds
1. Pass hydrogen sulfide through tin (II) chloride solution. Note the precipitate that is insoluble in dilute hydrochloric acid.
Sn²⁺ + S₂- --> SnS(s)
Filter off the precipitate and wash with distilled water. Transfer it to an evaporating basin and add yellow ammonium sulfide solution. The precipitate dissolves. Oxidation by the free sulfur in the ammonium sulfide occurs, so the S in the equation come from the ammonium sulfide.
(NH₄)₂S + SnS + S --> (NH₄)₂SnS₃ (ammonium thiostannate)
Add dilute acid to the ammonium thiostannate to precipitate tin (IV) sulfide, SnS₂.
2. Add drops of sodium hydroxide solution to tin (II) chloride solution. Note the white precipitate of tin (II) hydroxide that dissolves in excess sodium hydroxide to form sodium stannite.
Sn²⁺ + 2OH^- --> Sn(OH)₂(s)
Sn(OH)₂ + 2OH^- --> SnO₂^2- + 2H₂O (stannite ion = SnO₂^2-)
3. Add drops of ammonia solution, NH₃(aq) ("ammonium hydroxide") solution to tin (II) chloride solution. Note the white precipitate of tin (II) hydroxide that is not soluble in excess of ammonia solution, NH₃(aq) ("ammonium hydroxide").
4. To show that Tin (II) chloride is a powerful reducing agent, add tin (II) chloride solution to solutions of the following reagents. Reduction occurs with every reagent. Iron (IlI) chloride forms pale green iron (II) ions. Potassium permanganate forms manganese (II) ions. Potassium dichromate forms green chromic ions.

12.20.2 Prepare tin (IV) chloride
See diagram 12.13.2
Do this experiment in a fume cupboard. Put sand into the retort protect the glass during heating, followed by 5 cc of granulated tin. Insert the delivery tube and connect to a chlorine apparatus. Heat the retort while chlorine passes over. Note the ignition of the tin, the fine white crystals in the upper part of the retort, and the yellow distillate of tin (IV) chloride. The white crystals are SnCl₄.5H₂O, because of traces of moisture in the apparatus.
Sn + 2Cl₂ --> SnCl₄
1. Add drops of water to the tin (IV) chloride mixture and heat the mixture. Tests for hydrogen chloride. The white precipitate is hydrated tin (IV) oxide or a tin (IV) acid.
SnCl₄ + 4H₂O --> SnO₂.2H₂O + 4HCl
2. Add ammonia solution, NH₃(aq) ("ammonium hydroxide") solution to the tin (IV) chloride mixture and heat the mixture. Divide the suspension of hydrated tin (IV) oxide into two parts. To show the amphoteric nature of hydrated tin (IV) oxide dissolve part A in sodium hydroxide solution and dissolve part B in hydrochloric acid.
SnO₂.2H₂O + 4HCl --> SnCl₄ + 4H₂O
SnO₂2H₂O + 2NaOH --> Na₂SnO₃+ 3H₂O (Na₂SnO_{3 =} sodium stannate)

12.21.1 Reactions of zinc and its compounds
1. Hold a piece of zinc foil in the Bunsen flame, using tongs. Note the zinc oxide forms that is yellow when hot and white when cold.
2. Add sodium carbonate solution to zinc sulfate solution. Note the white precipitate of basic zinc carbonate, ZnCO₃.2Zn(OH)₂H₂O.
3. Add sodium hydrogen carbonate to zinc sulfate solution. Note the white precipitate of the normal carbonate, ZNCO₃.
4. Add drops of sodium hydroxide solution to zinc sulfate solution. Note the white precipitate of zinc hydroxide that dissolves in excess of sodium hydroxide solution to form sodium zincate. Pass hydrogen sulfide is passed through the sodium zincate solution. Note the white precipitate of zinc sulfide. Zinc hydroxide is amphoteric.
ZnSO₄ + 2NaOH --> Zn(OH)₂ + Na₂SO₄
Zn(OH)₂ + 2NaOH --> Na₂ZnO₂ + 2H₂O (Na₂ZnO_{2 =}sodium zincate)
5. Add drops of ammonium sulfide solution to zinc sulfate solution. Note the white precipitate of zinc sulfide that may be discoloured.
6. Dip a rolled filter paper into a concentrated solution of zinc sulfate with added cobalt nitrate solution. Burn the filter paper on wire gauze and note the remaining green ash, Rinmann's green.
8. Add drops of ammonia solution, NH₃(aq) ("ammonium hydroxide") to zinc sulfate solution. The precipitate of zinc hydroxide dissolves in excess, because of the formation of a complex ion [Zn(NH₃)₂]²⁺.