topic12C

School Science Lessons
Topic 12C Halogens, periodic table, sulfur, sulfuric acid
2012-01-24 SP
Please send comments to: J.Elfick@uq.edu.au

Table of contents
12.19.9 Bromine experiments
12.19.8 Chlorine experiments
12.19.7 Fluorine experiments
12.19.0 Halogens
12.19.6 Iodine experiments
12.19.10 Periodic table
12.18.0 Sulfur experiments
12.18.5 Sulfuric acid

12.19.9 Bromine experiments
Bromic acid, hydrobromic acid, HBr, hydrobromic acid solution, hydrogen bromide solution, conc. 9.0 M, Harmful
13.1.2 Bromine, Br, bromide ion Br^-, Toxic by all routes, extremely irritant vapour, skin irritant
1.8 Bromo compounds organic, List of bromo compounds
Bromo compounds, List
12.19.9.1 Reactions of bromine, Br₂12.19.9.1.1 Prepare bromine water
12.19.9.1.2 Open bromine gas ampoule
12.19.9.2 Reactions of bromine water, Br₂
12.19.9.3 Prepare hydrogen bromide, HBr
12.19.9.3a Heat 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^-12.19.9.7 Reaction of bromine water with zinc dust

12.19.8 Chlorine experiments
13.1.6 Chlorine gas, Cl₂
13.1.7 Chlorides, Cl^-

12.19.7 Fluorine experiments
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

12.19.0 Halogens
12.19.0 Halogens, Group 17 of the periodic table
12.19.5.0 CFCs, chlorofluorocarbons
12.19.2.1 Compare halogens, chlorine, bromine, iodine
3.32.1 Composition of the atmosphere and greenhouse gases
12.19.3.0 Displace a less reactive halogen from halogen compounds
12.19.5.1 "Freons"
12.19.5.2 Methyl bromide, bromoethane, CH₃Br
12.19.2.0 Pass halide vapour over hot iron wire to form iron halides
3.42.01 Prepare hydrochloric acid
3.42.0 Prepare hydrogen chloride, HCl
12.19.1.0 Properties of halogens
12.19.4.0 Reactions of silver halides, photography
3.42.1 Tests for hydrogen chloride

12.19.6 Iodine experiments
Iodine
Iodine compounds
12.19.6.4 Compare silver chloride, silver bromide and silver iodide
12.19.6.11 Decolorize iodine solution
12.19.6.0 Heat iodine crystals
12.19.6.6 Iodine extraction, I₂, tincture of iodine
12.9.6.0a Iodine with aluminium powder
12.19.6.7 Iodine with ammonium hydroxide
9.132 Iodine tests for starch
12.19.6.10 Iodine with starch
12.19.6.9 Make fingerprints with iodine
12.19.6.1 Prepare hydrogen iodide, HI
12.19.6.3 Prepare iodic acid, HIO₃, and potassium iodate, KIO₃
12.19.6.5 Prepare iodine solution, tincture of iodine
12.19.6.8 Prepare solid iodine from tincture of iodine
12.19.6.2 Reactions of iodides, I^-
12.19.6.12 Use iodine to write on iron

12.19.10 Periodic table
12.19.10 Periodic table
12.19.10.1 Patterns in the periodic table
12.17.1.1 Oxides and the periodic table

12.18.0 Sulfur experiments
8.2.15 Heat sulfur
12.18.1 Prepare forms of sulfur
12.18.2 Prepare sulfides
12.18.3.1 Prepare sulfur monochloride
12.18.3.2 Prepare thionyl chloride12.18.4 Properties of sulfur dioxide and sulfites
12.18.5 Sulfuric acid
12.18.6.1 Prepare sodium thiosulfate crystals, "hypo", Na₂S₂O₃.5H₂O
12.18.6.2 Reactions of sodium thiosulfate, Na₂S₂O₃.5H₂O
12.18.7 Reactions of sulfamic acid, NH₂.SO₂OH

12.18.5 Sulfuric acid
12.6.0.1 Formation of acid rain, SOx, by burning sulfur or sulfur compounds
12.6.0.4 Hydrogen peroxide oxidizes sulfur dioxide to sulfuric acid
12.18.5.4 Reactions of dilute sulfuric acid as an acid
12.18.5.5 Reactions of dilute sulfuric acid as a sulfate
12.18.5.0 Reactions of sulfuric acid
12.18.5.1 Sulfuric acid dehydrates copper (II) sulfate crystals
12.18.5.2 Sulfuric acid acts as a displacer of acids from their salts, sodium chloride, sodium acetate, sodium formate
12.18.5.3 Sulfuric acid acts as a displacer of acids from their salts, potassium bromide, potassium iodide
12.6.1 Sulfuric acid acts as an oxidizing agent
12.6.2 Sulfuric acid dehydrates copper (II) sulfate crystals
12.6.3 Sulfuric acid dehydrates sucrose (cane sugar)
12.6.0 Sulfuric acid, ionization of hydrogen sulfate ion, sulfur trioxide
14.1.3 Sulfuric acid with water
12.6.0.3.1 Sulfur dioxide to sulfuric acid 1.
12.6.0.3.2 Sulfur dioxide to sulfuric acid 2.
12.6.0.2 The 1 - 3% "sulfur" in coal is usually iron (II) sulfide (pyrite)

12.6.0 Sulfuric acid, ionization of hydrogen sulfate ion, sulfur trioxide
See diagram: 12.6.0
Sulfuric acid is a colourless oily liquid available as: 1. 2.0 M (4.0 N) 1.0 M (2.0 N) and 0.5 M (1.0 N) volumetric solutions 2. Minimum assay 97% solution density 1.83 g cm^-3 3. 98% "ANALAR" solution (d) "Battery acid" solution for lead cell accumulators minimum assay 30% density 1.25 g cm^-3 at 20^oC (battery acid). Sulfuric acid is a strong dibasic acid that forms sulfates and hydrogen sulfates a strong oxidizing agent that dissolves copper and a strong dehydrates agent that can remove water from organic compounds. Sulfuric acid is made by the contact process. Sulfur is burned or the ores zinc sulfide or iron sulfide (pyrites) are heated to form sulfur dioxide. The gases pass over vanadium (V) oxide or platinum catalyst at 450^oC to form sulfur trioxide that combines with water to form sulfuric acid.
Sulfur trioxide is produced by the action of oxygen on sulfur dioxide in the presence of a catalyst, e.g. iron oxide.

12.6.0.1 Formation of acid rain, SOx, by burning sulfur or sulfur compounds
When coal is burnt, the compounds that contain sulfur can form sulfuric acid, as in the equations below, to become components of acid rain (rainwater pH = 5.6, acid rain pH < 5). There may be more than one pathway for the formation of sulfuric acid from sulfur dioxide.

12.6.0.2 The 1 - 3% "sulfur" in coal is usually iron (II) sulfide (pyrite).
4 FeS₂ (s) + 11O₂ (g) --> 2Fe₂O₃ (s) + 8SO₂ (g)
S (s) + O₂ (g) --> SO₂ (g) sulfur dioxide
Also, other sulfide ores may produce sulfur dioxide in the atmosphere during smelting to obtain the pure metal.
Lead (II) sulfide (galena)
3PbS + 3O₂ (g) --> 3Pb (s) + 3SO₂ (g)
Copper (II) sulfide, chalcocite (bornite)
Cu₂S (s) + O₂ (g) --> 2Cu (s) + SO₂ (g)

12.6.0.3.1 Sulfur dioxide to sulfuric acid 1.
Sulfur dioxide is oxidized to sulfur trioxide by oxygen gas.
2SO₂ (g) + O₂ (g) <--> 2SO₃ (g)
Sulfur dioxide is oxidized to sulfur trioxide by nitrogen dioxide.
SO₂ (g) + NO₂ (g) --> SO₃ (g) + NO (g)
Sulfur trioxide dissolves in water to form sulfuric acid.
SO₃ (g) + H₂O (l) --> H₂SO₄ (aq) sulfuric acid

12.6.0.3.2 Sulfur dioxide to sulfuric acid 2.
Sulfur dioxide dissolves in water to form sulfurous acid
SO₃ (g) + H₂O (l) --> H₂SO₃ (aq)
Sulfurous acid is oxidized to to sulfuric acid by ozone
H₂SO₃ (aq) + O₃ (g) --> H₂SO₄ (aq) + O₂ (g)

12.6.0.4 Hydrogen peroxide oxidizes sulfur dioxide to sulfuric acid
SO₂ (g) + H₂O₂ (l) --> H₂SO₄ (aq)

12.6.1 Sulfuric acid acts as an oxidizing agent
See 12.3.13: Concentrated sulfuric acid with copper
BE CAREFUL! YOU ARE USING HOT CONCENTRATED SULFURIC ACID!
Add hot concentrated sulfuric acid to carbon. The reaction forms carbon dioxide and sulfur dioxide.
C (s) + 2H₂SO₄ (l) --> CO₂ (g) + 2SO₂ (g) + 2H₂O (l)
Add hot concentrated sulfuric acid to sulfur. The reaction forms sulfur dioxide and water.
S (s) + 2H₂SO₄ (l) --> 3SO₂ (g) + 2H₂O (l)
Add hot concentrated sulfuric acid to carbohydrates. The reaction forms carbon dioxide or carbon and water.

12.6.2 Sulfuric acid dehydrates copper (II) sulfate crystals
Add drops of sulfuric acid to blue copper (II) sulfate crystals. The crystals turn white as they lose water. Concentrated sulfuric acid combines so readily with water that it can be used as a dehydrates agent, e.g. removing water from hydrated copper (II) sulfate crystals and from other hydrated salts.
CuSO₄.5H₂O (s) <--> CuSO₄ (s) + 5H₂O (l)

12.6.3 Sulfuric acid dehydrates sucrose (cane sugar)
This favourite experiment forms a large quantity of toxic gases, e.g. carbon monoxide and sulfur dioxide, and the voluminous char (remaining carbon and other substances) should be washed thoroughly with water to remove any remaining acid before handling.
1. 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 (s) (H₂SO₄ catalyst) --> 12C (s) + 11H₂O (l)
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".
2. Add 35 mL of sulfuric acid to 50 g of sugar in a 100 mL beaker placed on a heat resistant mat.
3. Add 10 mL of sulfuric acid to a large test-tube almost filled with sugar. The sugar is dehydrated to form carbon and steam which causes the material to expand, forming a black porous column which rises out of the beaker or test-tube. Do the experiment in a fume cupboard because foul-smelling vapours are also released. To remove remaining acid, cool and wash the product thoroughly with water before touching it.
4. Put some sucrose (cane sugar) in a tall beaker. Add drops of concentrated acid to the sugar. BE CAREFUL!
The sugar turns yellow then brown then black and rises in the beaker. It reacts with carbohydrates like sugar and cellulose charring them by removing the elements of water from them and leaving a mass of black carbon behind.
5. Roll paper into a tube and hold it in the middle of a soft plastic container, e.g. ice cream tub. Do not use a glass jar. Fill the container with sugar. Pour just enough water to dampen the sugar down the tube to reach the bottom. Leave to stand for five minutes to allow the water to spread throughout the sugar. Remove the paper tube to leave a hole in the damp sugar. BE CAREFUL! Pour 30 mL of concentrated (98%) sulfuric acid down the hole and onto the top of the sugar. The sugar starts to turn brown, and black in patches. After some minutes bubbles of steam form. The reaction became more vigorous as the material in the container expands. A black cylinder rises out of the jar. Jets of steam spurt out. Heat is given out as the cylinder keeps rising. The black steaming cylinder is spongy carbon. Tap with a spatula to show it is hard, like expanded polystyrene packaging. If the carbon solidifies to make a seal over the top of the jar and the reaction continues deeper in the container, below the seal, pressure may build up to cause an explosion and a shower of black crumbling carbon.

12.18.1 Prepare forms of sulfur
See diagram 12.18.1: Sulfur crystals
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 yellow-green.

12.18.3.1 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 one-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 tests for sulfur dioxide and hydrochloric acid. Note the deposit of sulfur.
2S₂Cl₂ + 2H₂O --> 4HCl + 3S (s) + SO₂ (g)

12.18.3.2 Prepare thionyl chloride, SOCl₂
This experiment may not be allowed in some school systems.
See diagram 12.18.3.2: Prepare thionyl chloride
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 tests 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
See 3.51.1: Tests for sulfur dioxide | See 12.11.5.16: Tests for sulfates
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. Magnesium oxide forms.
2Mg + SO₂ (g) --> 2MgO (s) + S (s)
SO₂ (g) --> S (s) + O₂ (g)
2Mg (s) + O₂ (g) --> 2MgO (s)
4. Pass air (oxygen gas) 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.
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⁺

12.18.5.0 Reactions of sulfuric acid
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 sulfate crystals. After ten minutes, note the colour change from blue copper sulfate crystals to white anhydrous copper sulfate.
CuSO₄.5H₂O + (H₂SO₄) --> CuSO₄ + (H₂SO₄.5H₂O)
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) --> 12C + (H₂SO₄.11H₂O)
Sulfuric acid acts as a displacer of acids from their salts, sulfuric acid being much less volatile than most other acids.
3. 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
4. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium acetate. Note the smell of the displaced the acetic acid.
5. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium formate. Note the displacement of formic acid followed by dehydration.
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.
6. 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)
7. 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
8. Reactions of dilute sulfuric acid as an acid. 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)}
9. Add 2 cm of sodium carbonate to 1 cc of zinc powder. Test for carbon dioxide by passing the gas given off to pass into limewater that turns milky due to the fine precipitate of calcium carbonate.
2H⁺ + CO₃^2- --> H₂O + CO₂ (g)
Ca(OH)₂ + CO₂ (g) --> CaCO₃ (s) + H₂O
10. 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)

12.18.5.1 Sulfuric acid dehydrates copper (II) sulfate crystals
Sulfuric acid removes 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)

12.18.5.2 Sulfuric acid acts as a displacer of acids from their salts, sodium chloride, sodium acetate, sodium formate
Sulfuric acid is 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, potassium bromide, potassium iodide
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 gas 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)

12.18.6.1 Prepare 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-)

12.18.6.2 Reactions of sodium thiosulfate
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
(H₃NSO₃) (tautomer: NH₂.SO₂OH) (amidosulfonic acid, amidosulfuric acid, aminosulfonic acid, sulfamidic acid)
Sulfamic acid is used in preparations to clean stainless steel and copper utensils to remove hard water scale. It is also used to make sweeteners.
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 tests for ammonia with damp red litmus paper. Acidify the remaining solution with hydrochloric acid and add barium chloride solution to tests for sulfate ion.
5. Sulfamic acid solution in water is unstable and forms ammonium bisulfate, however, the colourless crystalline solid is stable, not hygroscopic, and is very soluble in water to form the zwitterion (H3N⁺SO3^-). On heating, the solution produces ammonia.
6. Sulfamic acid reaction with nitrous acid to form nitrogen
HNO₂ + (NH₂)HSO₃ --> H₂SO₄ + N₂ + H₂O
7. Sulfamic acid reacts with nitric acid to form nitrous oxide
HNO₃ + (NH₂)HSO₃ --> H₂SO₄ + N₂O + H₂O
8. Sodium hydroxide solution is standardized by titration with primary standard sulfamic acid solution
NaOH + (NH₂)HSO₃ --> NaNH₂SO₃ + H₂O

12.19.0 Halogens, Group 17 of the periodic table
See 1.6: Iodine solution
The halogens are yellow gas fluorine F₂, green-yellow gas chlorine Cl₂, red-brown liquid bromine Br₂, violet solid iodine I₂, 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, react with hydrogen in the decreasing order F-, then Cl^-, then Br^-, then I^-. Halogens react with most elements, with alkali metals to form salts. Halogens are toxic (poisonous) and are more soluble in hydrocarbon solvents than in water. 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 yellow-green 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, sodium fluoride solution, sodium chloride solution, sodium bromide solution and sodium iodide solution. Add chlorine solution to each solution. Chlorine has no visible effect on sodium fluoride solution or sodium chloride solution. Chlorine turns sodium bromide solution yellow-brown. Chlorine turns sodium iodide solution deep brown. So chlorine displaces bromine from sodium bromide and chlorine displaces iodine from sodium iodide. However, chlorine does not displace fluorine from sodium fluoride. So chlorine is more active than bromine and iodine but chlorine is less active than fluorine.
4. Compare the colours and solubility of silver fluoride, chloride, bromide and iodide by adding silver nitrate solution to sodium fluoride solution, sodium chloride solution, sodium bromide solution and sodium iodide solution. Silver fluoride is soluble. Silver chloride, silver bromide and silver iodide are insoluble. Silver chloride is white. Silver bromide is very pale yellow. Silver iodide is a deep yellow.

.	Fluorine	Chlorine	Bromine	Iodine
Element	yellow gas, irritating smell	yellow green, irritating smell	red liquid, irritating smell	black solid, irritating smell
Prepare	electrolysis of KHF₂	heat mixture of chloride, MnO₂ and conc. H₂SO₄	heat mixture of bromide, MnO₂, and conc. H₂SO₄	heat mixture of iodide, MnO₂, and conc. H₂SO₄
Activity	very reactive, combines with most metals and non-metals	very reactive, combines with most metals and many non-metals	very reactive, combines with most metals and non-metals	reactive, combines with most metals and few non-metals
Replacing action	replaces all halogens	replaces Br₂ and I₂ from bromides and 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 13.4.8: Burn steel wool in chlorine

12.19.2.1 Compare halogens, chlorine, bromine, iodine
1.1 Add drops of 0.02 M chlorine solution to Universal indicator paper on a white tile and note any change.
The chlorine solution bleaches the Universal indicator paper. Chlorine in water forms hydrochloric acid and bleach (chloric (I) acid).
1.2 Add drops of 0.02 M bromine solution, to Universal indicator paper on a white tile and note any change.
The bromine solution bleaches the area on the Universal indicator paper where it touched it.
1.3 Add drops of 0.02 M iodine solution to Universal indicator paper on a white tile and note any change.
The bromine solution faintly bleaches the area on the Universal indicator paper where it touched it.
2.1 Add 0.02 M chlorine solution to 0.2 M potassium bromide solution
Cl₂ (aq) + 2KBr (aq) --> 2KCl (aq) + Br₂ (aq) orange red-brown solution
chlorine + potassium bromide --> potassium chloride + bromine
Cl₂ (aq) + 2Br^- (aq) --> 2Cl^- (aq) + Br₂ (aq)
2.2 Add 0.02 M chlorine solution to 0.2 M potassium iodide solution
Cl₂ (aq) + 2KI (aq) --> 2KCl (aq) + I₂ (aq) black precipitate in a brown solution
chlorine + potassium iodide --> potassium chloride + iodine
Cl₂ (aq) + 2I^-(aq) --> 2Cl^- (aq) + I₂ (aq)
3.1 Add 0.02 M bromine solution to 0.2 M potassium chloride solution - No reaction
3.2 Add 0.02 M bromine solution to 0.2 M potassium iodide solution
Br₂ (aq) + 2KI (aq) --> 2KBr (aq) + I₂ (aq) black precipitate in a brown solution
bromine + potassium iodide --> potassium bromide + iodine
Br₂ (aq) + 2I^- (aq) --> 2Br^- (aq) + I₂ (aq)
4.1 Add 0.02 M iodine solution to 0.2 M potassium chloride solution - No reaction
4.2 Add 0.02 M iodine solution to 0.2 M potassium bromide solution - No reaction
Chlorine is a more powerful oxidizing agent than either bromine or iodine so it can take electrons from both bromide ions and iodide ions. Bromine is a more powerful oxidizing agent than iodine so it can remove electrons from iodide ions to give iodine. Oxidizing ability falls as you go down the Group, chlorine then bromine then iodine. Halogens get less reactive as the group is descended
The halogen molecule is the electron acceptor (the oxidizing agent) and is reduced by electron gain to form a halide ion.
The halide ion is the electron donor (the reducing agent) and is oxidized by electron loss to form a halogen molecule.

12.19.3.0 Displace a less reactive halogen from halogen 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) + Cl_{2 (}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 ammonia solution to the blackened chloride to produce diammine silver ion solution. 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, "Freons"
See 3.50.1: Ozone and photochemical smog
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:.
(1.) Freon 11 (CCl₃F, trichlorofluoromethane)
(2.) Tetrachloromethane (carbon tetrachloride, CCl₄, perchloromethane, dry cleaning fluid)
(3.) 1,1,1-trichloroethane (CH₃CCl₃, methyl chloroform, electrical equipment cleaner)
(4.) 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 with long half lives. 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	10.00	.
Halon-1301	10.00	.

12.19.5.1 "Freons"
Freon is a registered trademark for non-toxic non-flammable gases invented to avoid danger from leaking refrigerator gases. It is a name for compounds of ethane or methane with hydrogen atoms substituted by fluorine or chlorine, i.e. CFCs. The manufacture of Freons is being discontinued because of their ozone-depleting properties. The term "Freon" is not in favour nowadays. Examples of Freons include the following:
Freon 11, CCl₃F,. trichlorofluoromethane, (CFC-11)
Freon 12, CCl₂F₂, dichlorodifluoromethane, (CFC-12), b.p. -30^oC, (most common refrigerant gas, solvent, 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

12.19.5.2 Methyl bromide, bromoethane, CH₃Br, an organobromine compound, is colourless, odourless and not flammable. It is used to sterilize soil, kill rodents and plant pathogens. However, it depletes ozone in the atmosphere more than any CFC and its use is being phased out by most countries. However, seaweed, e.g. Asparagopsis taxiformis, produce great quantities of organohalides and produce the "smell" of the sea.

12.19.6.0 Heat iodine crystals
Use a Bunsen burner flame to heat gently < 0.5 g of iodine crystals in a test-tube with a stopper to show sublimation, then allow the test-tube to cool to show condensation. If heating more than 0.5 g of iodine strongly so only in a fume cupboard to avoid the irritating vapour.

12.9.6.0a Iodine with aluminium powder
Use < 5 g total material of iodine to react with aluminium powder in a fume cupboard. However, be aware that a cloud of unreacted iodine vapour may be released.

12.19.6.1 Prepare hydrogen iodide, HI
See diagram 12.19.6.1: Prepare hydrogen iodide
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, I^-
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, KIO₃
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 silver chloride, silver bromide and silver iodide
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 iodine solution, tincture of iodine
See: 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.

12.19.6.6 Iodine extraction, I₂, tincture of iodine
See 16.3.5.2.8: Povidone, PVP, porphyrins, pyrroles, polypyrroles, tetrapyrroles
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. Leave 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.
A tincture is an alcoholic solution of non-volatile substances with the alcohol acting as a solvent and preservative. Tincture of iodine, used in medicine, is iodine dissolved in ethanol. However, nowadays a povidone-iodine complex is used, e.g. Betadine.

12.19.6.7 Iodine with ammonium hydroxide
This experiment is very dangerous. Wear safety gloves eye and ear protection. Put 2 or 3 very small iodine crystals in a glass beaker. Do not use tincture of iodine. Pour ammonia solution over the crystals to just cover them. Mix the contents of the beaker by swirling it for a few seconds then leave the beaker in a fume cupboard or well-ventilated place for five minutes to allow a brown precipitate to form on in the beaker, fifteen minutes if using cloudy ammonia. Use masking tape to attach two layers of filter paper to a ring clamp attached to a retort stand and leave this apparatus in a safe place which is well-ventilated without wind or direct sunlight. Decant the liquid in the beaker to leave a brown solid. Dispose of decanted liquid down the sink with plenty of water. Wear safety gloves while using a spatula to transfer the brown solid to the filter paper then leave this apparatus in a safe place to dry for two hours leaving the light grey solid ammonium nitrogen triiodide. Be careful! Explosive! When the precipitate is thoroughly dry while wearing ear and eye protection and observers standing well away and protecting their ears, touch the precipitate with a long stick. Some people use a long stick with an attached feather to touch the ammonia nitrogen triiodide. The explosion leaves a purple vapour that should be avoided. Ammonia nitrogen triiodide may be found is some fireworks called Party Poppers but do not try to open these fireworks and remove this chemical

12.19.6.8 Prepare solid iodine from tincture of iodine
Leave 1 cm of iodine solution in an basin cup to evaporate. When only drops of liquid are left in the basin, transfer them to an evaporating dish. Put the dish on a tripod and gauze and place a small watch glass or saucer over the top. Heat the dish over a small flame. The remaining liquid quickly evaporates. Then the iodine vaporizes and becomes deposited as small shiny black crystals on the bottom of the watch glass. Scrape crystals from the bottom of the watch glass into a dry basin with a knife. Transfer the crystals to a dry test-tube and heat the test-tube. The violet vapour of iodine fills the test-tube, and shiny black crystals of solid iodine deposits again in the cooler part of the test-tube. Like ammonium chloride, iodine does not melt when heated but sublimes. To keep the iodine crystals put a stopper in the test-tube before the iodine evaporates.

12.19.6.9 Make fingerprints with iodine
Press the fingers close together on a sheet of plain white paper. While vaporizing the iodine from the evaporating dish in the previous experiment, remove the watch glass for seconds and place the sheet of paper with the finger-prints downward over the dish. When the paper is removed the prints show up clearly on the white paper. This method of developing fingerprints on cigarettes may used by crime investigators.

12.19.6.10 Iodine with starch
1. Use tincture of iodine to detect the presence of iodine starch, when an intensely black-blue complex forms. Some “clock reactions” produce iodine to be detected by the formation of a starch-iodine complex. The complex comprises a linear chain of iodine molecules inside the spiral of glucose molecules that form the starch structure.
Iodine with boiled starch solution produces an intense blue colour. Use the same substances to make an invisible ink. Write on a sheet of paper with a weak solution of boiled starch and when the writing is dry develop it by brushing over the paper a weak solution of iodine. The writing stands out in blue-black lettering. The writing disappear if the paper is warmed, but returns when the paper is cool.
2. Use tincture of iodine to detect the presence of iodine starch, when an intensely black-blue complex forms. Some “clock reactions” produce iodine to be detected by the formation of a starch-iodine complex. The complex comprises a linear chain of iodine molecules inside the spiral of glucose molecules that form the starch structure.

12.19.6.11 Decolorize iodine solution
Alkalis also decolorize iodine. Add drops of sodium hydroxide, sodium carbonate, or ammonia solution to iodine solution.

12.19.6.12 Use iodine to write on iron
Rub the blade of a penknife with sandpaper to clean the surface. Heat the blade and then stroke it on both sides with a small piece of candle or taper. When the blade has cooled a thin coating of wax is found to have adhered to the metal. Use a knitting needle to write initials in the wax. Write deeply enough to reach the iron. Cover the blade with iodine solution in a test-tube and leave it for half an hour. Remove the wax to find initials etched on the blade.

12.19.7.0 Fluorine, F, hydrofluoric acid, HF, Not permitted in schools
Hydrofluoric acid must not be produced as a reaction produce, e.g. by reaction of metal fluorides with mineral acids.
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, Not permitted in schools
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, SiF_4,Not permitted in schools
See diagram 12.19.7.2: Prepare silicon tetrafluoride
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 two-holes 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 hydrated silica as a white precipitate in hydrofluosilicic acid solution (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 may be 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, Cl^-
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

12.19.8.3 Prepare iron (IlI) chloride, FeCl₃
See diagram 12.19.8.3: Prepare iron (IlI) chloride
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 the salt solution 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 with potassium chlorate, KClO₃
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, KClO₄
See diagram 12.19.8.5: Solubility of potassium salts
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, AgCl₂
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, Br₂
Be careful! Liquid bromine can cause sores if in contact with the skin. Also, bromine vapour is painful to the eyes. Store bromine water in a screw-topped bottle in a refrigerator.
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 gas 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.

12.19.9.1.1 Prepare bromine water
Bromine gas, Br₂(g), and bromine liquid, Br₂(l), are too dangerous to prepare in schools so only bromine water may be prepared.
1. Pass chlorine gas through a bromide salt solution (Collect soluble gases in water)
Cl₂ + 2Br --> 2Cl^- + Br₂(aq)
2. Cool a 1 mL ampoule of bromine water, Br₂ (aq) in a refrigerator, break with forceps under 200 mL water. BE CAREFUL! Use safety glasses and nitrile chemical-resistant gloves.
3. Put a 1 mL ampoule of bromine water inside a test-tube of an appropriate size. Fix a short piece of wide bore rubber tubing to the end of the test-tube. Seal the other end of the rubber tubing with a wide bore glass rod. Move the ampoule inside the test-tube so that the top is inside the rubber tubing. Break the ampoule by squeezing the rubber tubing with pliers. Put in water and carefully cut the rubber tubing to release the bromine.
4. Prepare bromine water by adding a small volume of bromine to a large volume of water. The bromine dissolves to a limited extent, colouring the water yellow-orange. Use the relatively harmless bromine water to demonstrate reactions of bromine with various organic and inorganic materials. Reaction is shown by loss of colour from the solution.
5. Mix 5 g of potassium bromide and 5 g of manganese dioxide in a 250 mL conical flask. Add 5 mL of concentrated sulfuric acid and seal with a cork fitted with a distillation tube. Heat with a Bunsen burner to distil the bromine into a 200 mL conical flask containing 50 mL of water. Collect the gas in water solution. Make sure that the end of the distillation tube is just below the surface of the water. Do this reaction in a fume cupboard.
6. Add hydrogen peroxide and a few drops of concentrated sulfuric acid to a dilute solution of sodium or potassium bromide. A yellow colour develops after a few minutes

12.19.9.1.2 Open bromine gas ampoule
Open an ampoule of bromine gas with extreme care in a fume cupboard while wearing safety glasses. Cool the bottom of the ampoule in an ice / water mixture to reduce the vapour pressure of the bromine before opening. Use a glass knife (ceramic impregnated with diamond dust) or the edge of a new file to score around the outside of the neck. Wrap the ampoule in a cloth towel before cracking the ampoule open. Transfer the bromine to a container with a tightly fitting glass stopper or a Teflon-lined screw cap. This container should be placed in a larger container with some packing material to ensure it cannot leak or break. Both containers should be fully labelled. Always store the bromine in a cool secure store area.

12.19.9.2 Reactions of bromine water, Br₂
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 manganese sulfate solution and precipitate lead dioxide from lead nitrate solution.
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, HBr
See diagram 12.19.9.3: Prepare hydrogen bromide
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.

12.19.9.3a Heat hydrogen bromide, HBr
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, HBr
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 gas. 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 potassium carbonate solution, 20 g 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. Observe a drop of solution under a microscope for cubic crystals of the bromide.

12.19.9.6 Reactions of bromides, Br^-
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^-

12.19.9.7 Reaction with zinc dust
Zinc dust decolorizes bromine water.

12.19.10 Periodic table
See: Periodic table | See: Table of the elements
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: group (vertical), 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.

. .	group 1	group 2	group 3	group 4	group 5	group 6	group 17	group 18
1st period	H	.	.	.	.	.	.	He
2nd period	Li	Be	B	C	N	O	F	Ne
3rd period	Na	Mg	Al	Si	P	S	Cl	Ar
4th period	K	Ca	.	.	.	.	.	.

12.19.10.1 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.
Commercial periodic table chart, A4 sheet, three colour, pad of 100 sheets
Commercial Sargent-Welch three colour poster size, laminated chart with metal eyelets, double sided, in protective cardboard tube, 130 cm X 90 cm and 495 mm X 694 mm
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.