Group 7 AS

Fluorine (F2 ): very pale yellow gas. It is highly reactive Chlorine : (Cl2 ) greenish, reactive gas, poisonous in high concentrations Bromine (Br2 ) : red liquid, that gives off dense brown/orange poisonous fumes Iodine (I2 ) : shiny grey solid sublimes to purple gas. Trend in melting point and boiling point Increase down the group As the molecules become larger they have more electrons and so have larger van der waals forces between the molecules. As the intermolecular forces get larger more energy has to be put into break the forces. This increases the melting and boiling points Trend in electronegativity Electronegativity is the relative tendency of an atom in a molecule to attract electrons in a covalent bond to itself. As one goes down the group the electronegativity of the elements decreases. As one goes down the group the atomic radii increases due to the increasing number of shells. The nucleus is therefore less able to attract the bonding pair of electrons

1. The displacement reactions of halide ions by halogens. A halogen that is a strong oxidising agent will displace a halogen that has a lower oxidising power from one of its compounds The oxidising strength decreases down the group. Oxidising agents are electron acceptors. Chlorine will displace both bromide and iodide ions; bromine will displace iodide ions Chlorine (aq) Bromine (aq) Iodine (aq) potassium chloride (aq) Very pale green solution, no reaction Yellow solution, no reaction Brown solution, no reaction potassium bromide (aq) Yellow solution, Cl has displaced Br Yellow solution, no reaction Brown solution, no reaction potassium iodide (aq) Brown solution, Cl has displaced I Brown Solution, Br has displaced I Brown Solution, no reaction The colour of the solution in the test tube shows which free halogen is present in solution. Chlorine =very pale green solution (often colourless), Bromine = yellow solution Iodine = brown solution (sometimes black solid present) know these observations ! Cl2 (aq) + 2Br – (aq)
2Cl – (aq) + Br2 (aq) Cl2 (aq) + 2I – (aq)
2Cl – (aq) + I2 (aq) Br2 (aq) + 2I – (aq)
2Br – (aq) + I2 (aq) be able to write these reactions as two half equations showing oxidation or reduction e.g. 2Br – (aq)
Br2 (aq)+ 2eCl2 (aq)+ 2e-
2Cl- (aq)

2. The reactions of halide ions with silver nitrate. This reaction is used as a test to identify which halide ion is present. The test solution is made acidic with nitric acid, and then Silver nitrate solution is added dropwise. The role of nitric acid is to react with any carbonates present to prevent formation of the precipitate Ag2CO3 . This would mask the desired observations 2 HNO3 + Na2CO3
2 NaNO3 + H2O + CO2 Fluorides produce no precipitate Chlorides produce a white precipitate Ag+ (aq) + Cl- (aq)
AgCl(s) Bromides produce a cream precipitate Ag+ (aq) + Br- (aq)
AgBr(s) Iodides produce a pale yellow precipitate Ag+ (aq) + I- (aq)
AgI(s) The silver halide precipitates can be treated with ammonia solution to help differentiate between them if the colours look similar: Silver chloride dissolves in dilute ammonia to form a complex ion AgCl(s) + 2NH3 (aq)
[Ag(NH3 )2 ] + (aq) + Cl- (aq) Colourless solution Silver bromide dissolves in concentrated ammonia to form a complex ion AgBr(s) + 2NH3 (aq)
[Ag(NH3 )2 ] + (aq) + Br – (aq) Colourless solution Silver iodide does not react with ammonia – it is too insoluble.

3. The reaction of halide salts with concentrated sulphuric acid. Explanation of differing reducing power of halides A reducing agent donates electrons. The reducing power of the halides increases down group 7 They have a greater tendency to donate electrons. This is because as the ions get bigger it is easier for the outer electrons to be given away as the pull from the nucleus on them becomes smaller. The Halides show increasing power as reducing agents as one goes down the group. This can be clearly demonstrated in the various reactions of the solid halides with concentrated sulphuric acid. Know the equations and observations of these reactions very well. F- and Clions are not strong enough reducing agents to reduce the S in H2SO4 . No redox reactions occur. Only acid-base reactions occur. Fluoride and Chloride NaF(s) + H2SO4 (l)
NaHSO4 (s) + HF(g) Observations: White steamy fumes of HF are evolved. NaCl(s) + H2SO4 (l)
NaHSO4 (s) + HCl(g) Observations: White steamy fumes of HCl are evolved. These are acid –base reactions and not redox reactions. H2SO4 plays the role of an acid (proton donor). Br- ions are stronger reducing agents than Cl- and F- and after the initial acidbase reaction reduce the Sulphur in H2SO4 from +6 to + 4 in SO2 Bromide Acid- base step: NaBr(s) + H2SO4 (l)
NaHSO4 (s) + HBr(g) Redox step: 2HBr + H2SO4
Br2 (g) + SO2 (g) + 2H2O(l) Observations: White steamy fumes of HBr are evolved. Red fumes of Bromine are also evolved and a colourless, acidic gas SO2 Ox ½ equation 2Br –
Br2 + 2eRe ½ equation H2SO4 + 2 H+ + 2 e- → SO2 + 2 H2O Iodide I- ions are the strongest halide reducing agents. They can reduce the Sulphur from +6 in H2SO4 to + 4 in SO2 , to 0 in S and -2 in H2S. NaI(s) + H2SO4 (l)
NaHSO4 (s) + HI(g) 2HI + H2SO4
I2 (s) + SO2 (g) + 2H2O(l) 6HI + H2SO4 → 3 I2 + S (s) + 4 H2O (l) 8HI + H2SO4
4I2 (s) + H2S(g) + 4H2O(l) Observations: White steamy fumes of HI are evolved. Black solid and purple fumes of Iodine are also evolved A colourless, acidic gas SO2 A yellow solid of Sulphur H2S (Hydrogen Sulphide), a gas with a bad egg smell, Ox ½ equation 2I –
I2 + 2eRe ½ equation H2SO4 + 2 H+ + 2 e- → SO2 + 2 H2O Re ½ equation H2SO4 + 6 H+ + 6 e- → S + 4 H2O Re ½ equation H2SO4 + 8 H+ + 8 e- → H2S + 4 H2O Often in exam questions these redox reactions are worked out after first making the half-equations Reduction product = sulphur dioxide Note the H2SO4 plays the role of acid in the first step producing HBr and then acts as an oxidising agent in the second redox step. Note the H2SO4 plays the role of acid in the first step producing HI and then acts as an oxidising agent in the three redox steps Reduction products = sulphur dioxide, sulphur and hydrogen sulphide

4. The disproportionation reactions of chlorine and chlorate(I). Disproportionation is the name for a reaction where an element simultaneously oxidises and reduces. Chlorine with water: Cl2 (aq) + H2O(l)
HClO(aq) + HCl (aq) Chlorine is both simultaneously reducing and oxidising If some universal indicator is added to the solution it will first turn red due to the acidity of both reaction products. It will then turn colourless as the HClO bleaches the colour. Reaction with water in sunlight If the chlorine is bubble through water in the presence of bright sunlight a different reaction occurs 2Cl2 + 2H2O
4H+ + 4Cl- + O2 The same reaction occurs to the equilibrium mixture of chlorine water. The greenish colour of chlorine water fades as the Cl2 reacts and a colourless gas (O2 ) is produced Chlorine is used in water treatment to kill bacteria. It has been used to treat drinking water and the water in swimming pools. The benefits to health of water treatment by chlorine outweigh its toxic effects. Reaction of Chlorine with cold dilute NaOH solution: Cl2 ,(and Br2 , I2 ) in aqueous solutions will react with cold sodium hydroxide. The colour of the halogen solution will fade to colourless Cl2 (aq) + 2NaOH(aq)
NaCl (aq) + NaClO (aq) + H2O(l) The mixture of NaCl and NaClO is used as Bleach and to disinfect/ kill bacteria The greenish colour of these solutions is due to the Cl2 Naming chlorates/sulphates In IUPAC convention the various forms of sulfur and chlorine compounds where oxygen is combined are all called sulfates and chlorates with relevant oxidation number given in roman numerals. If asked to name these compounds remember to add the oxidation number. NaClO: sodium chlorate(I) NaClO3 : sodium chlorate(V) K2SO4 potassium sulfate(VI) K2SO3 potassium sulfate(IV)

A-level Chemistry exemplar for required practical No. 4 To carry out tests for the presence of cations and anions and to make accurate observations Student sheet These tests may be split over several lessons. Requirements You are provided with the following: General  test tubes and stoppers  test-tube racks  plastic graduated dropping pipettes  deionised or distilled water  forceps. Test 1  0.1 mol dm–3 barium chloride solution  0.6 mol dm–3 sodium hydroxide solution  0.1 mol dm–3 calcium bromide solution (or calcium nitrate/potassium bromide)  0.1 mol dm–3 magnesium chloride solution  0.1 mol dm–3 strontium chloride solution. Test 2  0.1 mol dm–3 barium chloride solution  1.0 mol dm–3 sulfuric acid  0.1 mol dm–3 calcium bromide solution (or calcium nitrate/potassium bromide)  0.1 mol dm–3 magnesium chloride solution  0.1 mol dm–3 strontium chloride solution. Test 3  0.1 mol dm–3 ammonium chloride  0.4 mol dm–3 sodium hydroxide solution  red litmus paper  kettle  water bath. Test 4  0.4 mol dm–3 sodium hydroxide solution  red litmus paper (or universal indicator paper)  1.0 mol dm–3 ammonia solution  petri dish with lid. Test 5  0.5 mol dm–3 sodium carbonate solution  0.5 mol dm–3 hydrochloric acid  0.02 mol dm–3 calcium hydroxide solution (limewater). Test 6  0.1 mol dm–3 barium chloride solution  0.1 mol dm–3 magnesium sulfate solution. Test 7  0.1 mol dm–3 potassium chloride solution  0.1 mol dm–3 potassium bromide solution  0.1 mol dm–3 potassium iodide solution  0.1 mol dm–3 nitric acid  0.05 mol dm–3 silver nitrate solution  concentrated ammonia solution  2.0 mol dm –3 ammonia solution Test 8  potassium chloride solid  potassium bromide solid  potassium iodide solid  0.1 mol dm–3 lead nitrate solution (or lead ethanoate solution)  blue litmus paper  filter paper  small spatula.  concentrated sulfuric acid in dropping bottles  0.5 mol dm–3 acidified potassium dichromate(VI) solution Suggested method In every case, you should present all of your observations in a neat table. The presentation of a clearly organised record of your observations is an important skill which you will be expected to demonstrate. Tests 1 and 2: Testing for Group 2 metal cations Test 1 – dilute sodium hydroxide a) Place about 10 drops of 0.1 mol dm–3 barium chloride in a clean test tube. Add about 10 drops of 0.6 mol dm–3 sodium hydroxide solution, mixing well. b) Now continue to add this sodium hydroxide solution, dropwise with gentle shaking, until in excess. The test tube should not be more than half full. Once completed, dispose of the contents by placing the test tube in a bowl of water. c) Repeat this test with the calcium bromide, magnesium chloride and strontium chloride. Test 2 – dilute sulfuric acid a) Place about 10 drops of 0.1 mol dm–3 barium chloride in a clean test tube. Add about 10 drops of 1.0 mol dm–3 sulfuric acid, mixing well. b) Now continue to add this dilute sulfuric acid, dropwise with gentle shaking, until in excess. The test tube should not be more than half full. Once completed, dispose of the contents by placing the test tube in a bowl of water. c) Repeat this test with the calcium bromide, magnesium chloride and strontium chloride. Test 3: Testing for ammonium ions a) Place about 10 drops of 0.1 mol dm–3 ammonium chloride in a clean test tube. Add about 10 drops of 0.4 mol dm–3 sodium hydroxide solution. Shake the mixture. b) Warm the mixture in the test tube gently using a water bath. c) Test the fumes released from the mixture by using forceps to hold a piece of damp red litmus paper in the mouth of the test tube. d) Dispose of the contents by using the previous method. Tests 4, 5, and 7: Tests for anions in aqueous solution Test 4: Test for hydroxide ions in aqueous solution a) Test about 1 cm3 of 0.4 mol dm–3 sodium hydroxide solution in a test tube with red litmus paper or universal indicator paper. b) Record your observations. Dispose of the test tube contents. This approach can also be used to test for the alkaline gas, ammonia, which forms hydroxide ions when it comes into contact with water. c) Take 5 drops of 1 mol dm–3 ammonia solution and place on a filter paper and place inside a petri dish with lid. Dampen a piece of red litmus paper with deionised water and place on the other side of the petri dish. Replace the lid and observe over a few minutes. d) Record your observations. Test 5: Test for carbonate ions in aqueous solution a) Have about 2 cm3 of calcium hydroxide (limewater) ready in a test tube. b) To about 3 cm3 of 0.5 mol dm–3 sodium carbonate solution in a test tube, add an equal volume of 1.0 mol dm-3 dilute hydrochloric acid. c) Immediately put in delivery tube with open end into the limewater test tube. Make sure that the end of the tube is below the level of the liquid. d) Record your observations. Dispose of the test tube contents. Test 6: Test for sulfate ions in aqueous solution a) To about 1 cm3 of 0.1 mol dm–3 magnesium sulfate solution in a test tube, add an equal volume of dilute hydrochloric acid followed by an equal volume of 0.1 mol dm–3 barium chloride solution. b) Record your observations. Dispose of the test tube contents. Test 7: Test for halide ions in aqueous solution Test for chloride, bromide and iodide ions in aqueous solution a) Place about 10 drops of 0.1 mol dm–3 potassium chloride in a clean test tube. Add about 5 drops of dilute nitric acid. Shake well. b) To the solution add another 10 drops of 0.05 mol dm–3 silver nitrate solution. c) Then add an excess of 2 mol dm–3 ammonia solution and shake to mix thoroughly. Dispose of the tube contents d) Repeat steps a) and b), but this time add an excess of concentrated ammonia solution, working in a fume cupboard. Dispose of the tube contents e) Repeat steps a) to d) but using potassium bromide and then potassium iodide instead of potassium chloride. Test 8: Test for halide ions in solid salts using concentrated sulfuric acid Test for chloride, bromide and iodide ions in solid potassium halides Note: Gloves must be worn for this procedure These experiments must be done in a fume hood a) Place a small spatula of solid potassium chloride in a clean dry test tube. b) Slowly add a few (2 to 5) drops of concentrated sulfuric acid. c) Record what happens. d) Test the gas evolved with moist blue litmus paper, taking care that the paper does not touch the sides of the test tube. e) Repeat this experiment with solid potassium bromide, but this time test the gas produced using a narrow strip of filter paper that has been dipped in acidified potassium dichromate solution. f) Repeat this experiment with potassium iodide, but this time test the gas produced using a narrow strip of filter paper that has been dipped in lead nitrate solution. Additional notes In test 3, step (b) will work slowly at room temperature or use water from a recently boiled kettle poured into a beaker. In test 8, step (b), only 2 to 5 drops of concentrated sulfuric acid should be added and this should be done slowly. Sample results Test 1 Barium chloride Calcium bromide Magnesium chloride Strontium chloride Initial colourless solution colourless solution colourless solution colourless solution 10 drops of 0.6 mol dm–3 NaOH colourless solution slight white precipitate slight white precipitate slight white precipitate Excess NaOH colourless solution slight white precipitate white precipitate slight white precipitate Test 2 Barium chloride Calcium bromide Magnesium chloride Strontium chloride 10 drops of 1.0 mol dm–3 H2SO4 white precipitate slight white precipitate slight white precipitate white precipitate Excess H2SO4 white precipitate slight white precipitate colourless solution white precipitate Test 3 Ammonium chloride + sodium hydroxide – Damp red litmus paper = blue Test 4 Sodium hydroxide – Damp red litmus paper = blue Ammonia solution – Damp red litmus paper = blue Test 5 Limewater – colourless solution to cloudy Test 6 Test of sulfate ions – white precipitate forms. Test 7 + HNO3 + 2 M NH3 + concentrated NH3 Potassium chloride white precipitate colourless solution colourless solution Potassium bromide cream precipitate cream precipitate colourless solution Potassium iodide yellow precipitate yellow precipitate yellow precipitate Test 8 Concentrated sulfuric acid Result of paper test Potassium chloride effervescence red Potassium bromide effervescence brown gas produced solution turns deep brown/red Potassium iodide solution goes red/brown immediately brown gas produced turns black/grey