Student Material 1 (Hal

Mandatory Experiment 1.2

Investigation of (a) Redox reactions of the halogens (b) Displacement reactions of metals

Student Material

(a) Redox reactions of the halogens

Theory

The halogens, fluorine, chlorine, bromine, iodine and astatine are very reactive elements and are too unstable to exist in nature in an uncombined form. They often react by taking an electron from another element. This means they react as oxidising agents. The smaller the halogen atom, the stronger the oxidising agent it is. So in terms of oxidising power

F > Cl > Br > I > At

Fluorine is extremely poisonous, and astatine is unstable and radioactive, so the investigation here is confined to the other three halogens.

(i) Reactions with halides

Chlorine, being the strongest oxidising agent of the three, is capable of releasing the other two elements from solutions of their salts: - - Cl2 + 2Br (aq)  2Cl (aq) + Br2 - - Cl2 + 2I (aq)  2Cl (aq) + I2 Bromine can release iodine from a solution of its salts - - Br2 + 2I (aq)  2Br (aq) + I2

Your task is to find the evidence to support this theory.

Chemicals and Apparatus

Chlorine solution i Bromine solution Iodine solution Sodium chloride solution Sodium bromide solution Potassium iodide solution

Safety glasses

1 PVC gloves Fume cupboard or well-ventilated room Pasteur pipettes Test tubes Test tube rack Test tube brush

Quantities needed per working group

Name of solution Quantity

Aqueous solutions of chlorine, 2 cm3 aliquots (portions) per test bromine and iodine. Aqueous solutions of chloride, bromide, 2 cm3 aliquots per test and iodide salts.

Procedure

NB: Wear your safety glasses.

1. Copy the following table into your practical report book and fill in your observations. By reference to this table you will be able to draw conclusions from your later observations.

Name of solution Colour of solution Chlorine in water Bromine in water Iodine in water Chloride ions in water Bromide ions in water Iodide ions in water

Table 1

2. Draw a second table into your practical report book with the following headings:

Solutions added to the Observation Conclusion test-tube (a) Chlorine and bromide ions (b) Chlorine and iodide ions (c) Bromine and iodide ions

Table 2

2 3. As there are as many as 16 reagent bottles in use in these experiments it is vital to the success of the experiment not to mix the reagents by putting the wrong stopper on the wrong bottle. Your teacher will demonstrate to you the correct way to hold the stopper while using a reagent bottle. Always replace stoppers immediately after use.

4. For each of the cases (a), (b) and (c) described in Table 2, add 2 cm3 of the solutions mentioned to separate test tubes and mix.

Record your observations and conclusions. Retain the contents of the test tubes for comparison purposes, ensuring that the test tubes are correctly labelled.

(ii) Reactions with iron(II) salts and with sulfites

All three halogen solutions are able to oxidise iron(II) ions to iron(III) ions, and to oxidise sulfite ions to sulfate ions in aqueous solution. For example, chlorine reacts as follows: 2+ - 3+ Cl2 + 2Fe (aq)  2Cl (aq) + 2Fe (aq) 2- - 2- Cl2 + SO3 (aq)  2Cl (aq) + SO4 (aq)

Your task is to find the evidence to support this theory.

Chemicals and Apparatus

Chlorine solution i

Iron(II) sulfate solution n

Iron(III) chloride solution i

3 Sodium sulfite solution Sodium hydroxide solution Silver nitrate solution

Barium chloride solution n Dilute hydrochloric acid

Dilute ammonia solution i

Safety glasses PVC gloves Fume cupboard or well-ventilated room Pasteur pipettes Test tubes Test tube rack Test tube brush

Quantities needed per working group

Name of solution Quantity

Aqueous solutions of chlorine, 2 cm3 aliquots (portions) per test bromine and iodine. Aqueous solutions of iron(II) sulfate, 2 cm3 aliquots per test iron(III) chloride and sodium sulfite. Sodium hydroxide solution. 2 cm3 Aqueous solutions of silver nitrate, 2 cm3 aliquots per test barium chloride, dilute hydrochloric acid and dilute ammonia.

Procedure

NB: Wear your safety glasses.

1. Copy the following table into your practical report book and fill in your observations. By reference to this table you will be able to draw conclusions from your later observations.

Name of solution Colour of solution Chlorine in water Iron(II) sulfate in water Iron(III) chloride in water Iron(II) ions in sodium hydroxide solution

4 Iron(III) ions in sodium hydroxide solution

Table 3

2. Draw another table into your practical report book with the following headings:

Solutions added to the Observation Conclusion test-tube (d) Chlorine and iron(II) sulfate followed by 10 drops of sodium hydroxide (e) Chlorine and sodium sulfite followed by the test for the presence of chloride ions or sulfate ions

Table 4

3. Check that the sulfite solution does not contain any sulfate ions, as follows: Add 2 cm3 of sodium sulfite solution to a clean test tube. Using a dropping pipette add a few drops of barium chloride solution. A white precipitate forms. Now add 2 cm3 of dilute hydrochloric acid. The white precipitate should dissolve. If any of the white precipitate does not dissolve, then sulfate ions are present, and so this solution cannot be used.

4. Check that the iron(II) sulfate solution does not contain any iron(III) ions by adding 10 drops of sodium hydroxide solution and studying the colour of the precipitate formed and comparing it with the table set up in part 1.

5. For each of the cases (d) and (e) described in Table 4, add 2 cm3 of the solutions mentioned to separate test tubes and mix.

6. Record your observations and conclusions. Retain the contents of the test tubes for comparison purposes, ensuring that the test tubes are correctly labelled.

7. Iron(II) ions and iron(III) ions form different coloured floating (flocculent) precipitates with sodium hydroxide. This allows you to determine whether a reaction has or has not taken place in (d) above.

8. In (e) above, the test for sulfate ions is carried out as described in 3 above. The test for chloride ions is carried out as follows: Add the solution to be tested to a clean test tube. Using a dropping pipette add a few drops of silver nitrate solution. If chloride is present, a white precipitate forms. Now add 2 cm3 of dilute ammonia solution. The white precipitate should dissolve.

6 (b) Displacement reactions of metals

Theory

In this experiment, magnesium and zinc respectively are reacted with a solution of copper(II) sulfate.

Metals higher up on the electrochemical series displace metals lower down from aqueous solutions of their salts. The metal higher up in the series is oxidised in the process and forms a soluble positive ion. The metal lower down in the series is reduced, gains electrons and becomes a solid metallic element.

Use of this principle will allow, for example, the use of scrap iron (higher up in the series) to liberate copper (lower down in the series) from an aqueous solution of copper(II) sulfate. Fe(s) + CuSO4(aq)  FeSO4(aq) + Cu(s)

Equations

Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s)

Mg(s) + CuSO4(aq)  MgSO4(aq) + Cu(s)

This experiment works best under acidic conditions; under these conditions, the following reactions take place simultaneously:

Zn(s) + H2SO4(aq)  ZnSO4(aq) + H2 (g)

Mg(s) + H2SO4(aq)  MgSO4(aq) + H2 (g)

Your task is to find the evidence to support this theory.

Chemicals and Apparatus Acidified copper(II) sulfate solution Zinc powder Magnesium ribbon

Safety glasses Boiling tubes Boiling tube rack Test tube brush Pasteur pipettes

7 Quantities needed per working group

Acidified copper(II) sulfate solution 2 cm3 aliquots per test Zinc powder 1.0 g Magnesium ribbon (cleaned with sand- 0.5 g paper)

Procedure

NB: Wear your safety glasses.

1. Copy the following table into your practical report book

Mg Zn (a) Colour of CuSO4(aq) at the beginning (b) Colour of the solution at the end of the reaction (c) Colour of the precipitate formed (d) Any other observations

Table 5

2. Draw another table into your practical report book with the following headings:

Observation Conclusion Mg (a) (b) (c) (d) Zn (a) (b) (c) (d)

Table 6

8 3. Half fill two boiling tubes with the acidified copper(II) sulfate solution.

4. Add the magnesium ribbon to the solution in one boiling tube and the zinc powder to the solution in the other boiling tube. Record your observations and conclusions.

Questions relating to the experiment

1. Write out a full list of equations for all the redox reactions that took place during this experiment. Using oxidation numbers, label the species that have been oxidised and reduced, e.g. - - Cl2 + 2Br (aq)  2Cl (aq) + Br2 0 -1 -1 0 r o

2. The iodine/thiosulfate titration to be dealt with in experiment 4.7 is another redox reaction involving a halogen. The equation for the reaction is 2Na2S2O3 + I2  2NaI + Na2S4O6 Show that the iodine is acting as an oxidising agent there as well.

3. Describe the tests you would use to distinguish between sulfite and sulfate anions in aqueous solution.

4. Explain why it is difficult to make an aqueous solution of iodine. What particular method is used to overcome this problem?

5. An aqueous solution of chlorine is often made by reacting concentrated hydrochloric acid with a diluted commercial bleach

9 solution. The active ingredient in commercial bleach is sodium hypochlorite (NaOCl). The equation for the reaction is NaOCl + 2HCl  Cl2 + NaCl + H2O

Show what species are oxidised and reduced during this reaction.

6. The position of zinc in the Periodic Table would allow you to predict the colour of its sulfate salt solution. Explain.

7. Explain why magnesium is more reactive than zinc.

8. What would you expect to see happen if a piece of copper wire was suspended in a solution of silver nitrate? (Silver nitrate is very expensive but your teacher may be able to demonstrate this experiment.)

9. Carry out some research to find out why commercial photography laboratories might have a special interest in these kinds of reactions.

10 Teacher Material

 This experiment needs quite an amount of preparation time. Additional value can be got from it by making up the stable solutions during previous single periods, thereby familiarising your students with the appearance of the solutions and with the precautions needed when working with them.

 Students generally find this material difficult. It can be useful to repeat some of the reactions as demonstrations at a later stage when reviewing the laboratory notebooks.

 Silver nitrate solution is light-sensitive and should be stored in a dark container, away from sunlight.

 The following solutions must be prepared fresh: aqueous solution of chlorine, aqueous iron(II) sulfate solution and aqueous sodium sulfite solution.

 In the preparation of chlorine solution from sodium hypochlorite solution and hydrochloric acid, the reaction is as follows: - - 2HCl + OCl → Cl2 + H2O + Cl

 The full balanced equation for the reaction of chlorine with sulfite ions is as follows: 2- - 2- + Cl2 + SO3 + H2O  2Cl + SO4 + 2H

 Clean the magnesium ribbon using a small piece of sandpaper before weighing it. Measure the length of the weighed piece and use that approximate length for everyone.

 If the displacement of metals experiments are not carried out using acidic conditions, a black precipitate rather than a copper-coloured precipitate is likely to be formed.

Preparation of reagents

Distilled or deionised water should be used in making up all these solutions.

Aqueous solution of chlorine: This solution must be freshly prepared. Take 100 cm3 of commercial bleach. Add it to 500 cm3 of water. In a fume-cupboard, add concentrated hydrochloric acid drop by drop with constant stirring until a drop of the solution is just acid to litmus. Put 50 cm3 of the solution into each of 12 reagent bottles.

Aqueous solution of bromine: This is best purchased as a solution because liquid bromine itself is a poisonous liquid and difficult to work with. The solution purchased

11 may be diluted before use so as to match the concentrations of the other reagents. Bromine water does deteriorate with time but lasts much longer than chlorine water, provided that it is stored in a brown bottle. In the event that the solution has to be made up directly using bromine, in a fume- cupboard, shake 0.5 cm3 bromine with 100 cm3 water. Store in a tightly stoppered bottle.

Aqueous solution of iodine: Iodine crystals dissolve very poorly in water but dissolve - readily in an aqueous solution of potassium iodide forming KI3(aq). The I3 ion releases I2 molecules in reactions and so is always treated as an aqueous solution of iodine. Dissolve 20 g of potassium iodide in 500cm3 of water. Add about 10 g of iodine crystals, dissolve and make up to 1 litre. This solution is quite stable.

Aqueous solution of sodium chloride (approximately 0.2 mol l-1): Dissolve about 12 g in water and make up to a litre.

Aqueous solution of sodium bromide (approximately 0.2 mol l-1): Dissolve about 20 g in water and make up to a litre.

Aqueous solution of potassium iodide (approximately 0.2 mol l-1): Dissolve about 33 g in water and make up to a litre.

Aqueous sodium hydroxide solution: Dissolve 10 g of pellets in 200cm3 of cold water and make up to 250 cm3 to make an approximately 1 mol l-1 solution. Special care is required as this is a caustic solution.

Aqueous iron(II) sulfate solution: This solution must be freshly prepared. Dissolve 11.2 g of the crystalline salt in 100 cm3 of water containing 2 cm3 of concentrated sulfuric acid, and dilute to 200 cm3 to make an approximately 0.2 M solution.

Aqueous iron(III) chloride solution: Dissolve 11 g of the hydrated salt in 100 cm3 of water containing 4 cm3 of concentrated hydrochloric acid, and dilute to 200 cm3 to make an approximately 0.2 M solution. . Aqueous sodium sulfite solution: This solution must be freshly prepared. Dissolve 5.2 g of the salt in 100 cm3 of water, and dilute to 200 cm3 to make an approximately 0.2 M solution.

Aqueous ammonia solution: In a fume-cupboard, dilute 40 cm3 of concentrated ammonia solution to 250 cm3 to make an approximately 3 M solution.

Silver nitrate solution: Dissolve 4 g of the crystals and make up to 250 cm3. This makes an approximately 0.1 mol dm-3 solution. The solution must be made with deionised water and stored in a brown bottle. Light tends to reduce silver ions, and halide ions in tap water would form a precipitate with the silver ions. This solution is also best freshly prepared and in very small quantities.

12 Aqueous solution of barium chloride (approximately 0.2 mol l-1): Dissolve about 40 g in water and make up to 1 litre.

Dilute aqueous solution of hydrochloric acid (approximately 2 mol l-1): In a fume cupboard, add about 170 cm3 of concentrated hydrochloric acid slowly with stirring to about 500 cm3 of water and make up to 1 litre.

Acidified copper(II) sulfate solution: Dissolve 5 g of copper(II) sulfate pentahydrate 3 3 (CuSO4.5H2O) in about 100 cm water and make up to 200 cm . This makes an approximately 0.1 mol l-1 solution. Carefully add 20 cm3 of concentrated sulfuric acid.

Quantities needed per working group

100 cm3 of each of the solutions should be placed in 125 cm3 reagent bottles.

Name of solution Quantity

Aqueous solutions of chlorine, 2 cm3 aliquots (portions) per test bromine and iodine. Aqueous solutions of chloride, bromide, 2 cm3 aliquots per test and iodide salts. Aqueous solutions of iron(II) sulfate, 2 cm3 aliquots per test iron(III) chloride and sodium sulfite. Aqueous starch solution. A few drops Sodium hydroxide solution. A few drops Aqueous solutions of silver nitrate, 2 cm3 aliquots per test barium chloride, dilute hydrochloric acid and dilute ammonia. Aqueous solution of copper(II) sulfate. 2 cm3 aliquots per test Zinc powder 1.0 g Magnesium ribbon (cleaned with sand- 0.5 g paper)

Safety Considerations

 The chlorine and bromine solutions and their vapours are poisonous. Consequently, a fume cupboard with proper ventilation should ideally be used for some parts of this experiment.

Chemical hazard notes

Aqueous solution of chlorine i: Vapour attacks lungs, eyes and nose.

13 Bromine : The vapour is highly toxic by inhalation. The liquid causes severe burns to eyes and skin. The aqueous solution attacks lungs, eyes and nose.

Iodine n: Harmful by skin contact and by inhalation. Avoid eye contact. Silver nitrate : Solutions are very dangerous to the eyes and blacken skin. Sodium hydroxide : Caustic, harmful to skin and especially to eyes. Always wear eye protection. Ammonia solution : Pungent vapour toxic by inhalation; irritating to eyes and respiratory system; in case of contact with eyes wash immediately with plenty of water and seek medical advice.

Barium chloride n: Harmful by ingestion or inhalation of dust. Concentrated hydrochloric acid : Very corrosive to eyes and skin, and its vapour is very irritating to lungs. Concentrated sulfuric acid : Very corrosive to eyes and skin. Due to its very considerable heat of reaction with water, it is essential that the acid be added to water when it is being diluted.

Iron(II) sulfate n: Harmful if swallowed. Irritating to eyes and skin.

Iron(III) chloride i: Eye and skin irritant. Severe eye burns may result if left unattended.

Copper(II) sulfate n: Skin and eye irritant. Harmful if ingested.

Magnesium : Flammable; burns with an intense flame. Poisonous by ingestion. Inhalation of dust harmful.

Zinc : Zinc dust at the bottom of the container could be flammable.

Disposal of wastes

Solid products of the reactions such as barium sulfate and copper should be filtered, mixed with sand and placed in a refuse bin. Except in the cases that follow, residual liquid waste should be diluted with excess water and flushed to the foul water drain. Waste containing bromine, as a result of the reaction between chlorine and bromide ions, should be treated with 10% sodium carbonate solution before diluting with excess water. Waste containing iodine, as a result of the reaction between bromine and iodide ions, should be treated with 25% sodium thiosulfate solution before diluting with excess water. Residual liquid waste from the replacement reactions of metals experiment should be neutralised with sodium carbonate before diluting with excess water.

14 Specimen results (a)

Name of solution Colour of solution Chlorine in water Pale green Bromine in water Yellow/orange Iodine in water Brown/red Chloride ions in water Colourless Bromide ions in water Colourless Iodide ions in water Colourless Iron(II) sulfate in water Pale green Iron(III) chloride in water Yellow Iron(II) ions in sodium hydroxide solution Green muddy precipitate Iron(III) ions in sodium hydroxide solution Brown precipitate

Solutions added to the Observation Conclusion test-tube

bromine ions oxidised to bromine

(a) Chlorine and bromide Orange/yellow solution Bromide ions oxidised to ions formed bromine iodide ions oxidised to iodine

(b) Chlorine and iodide ions Brown/red solution formed Iodide ions oxidised to iodine

(c) Bromine and iodide ions Darkening of solution Iodide ions oxidised to iodine (d) Chlorine and iron(II) Orange/brown muddy Iron(II) ions oxidised to sulfate followed by a few precipitate iron(III) ions drops of sodium hydroxide (e) Chlorine and sodium Sulfate test will show a Sulfite ions oxidised to sulfite followed by the test permanent white precipitate sulfate ions for the presence of sulfate

15 ion iron(II) ions oxidised to iron(III) ions or the test for the presence a white precipitate which Chlorine has been reduced of chloride ions dissolves with the addition to chloride ions of dilute aqueous ammonia

Specimen results (b)

Mg Zn (a) Colour of Blue Blue CuSO4(aq) at the beginning (b) Colour of the Colourless Colourless solution at the end of the reaction (c) Colour of the Brown Brown precipitate formed (d) Any other Gas Gas observations evolved evolved

Observation Conclusion 2+ Mg Blue solution Cu (aq) present 2+ Colourless solution Mg (aq) present Brown precipitate Copper metal powder Gas evolved Hydrogen produced 2+ Zn Blue solution Cu (aq) present 2+ Colourless solution Zn (aq) present Brown precipitate Copper metal powder Gas evolved Hydrogen produced

Solutions to student questions

16 r o

The following are the equations to be labelled by the student:

- - Cl2 + 2Br (aq)  2Cl (aq) + Br2 0 -1 -1 0 r o

- - Cl2 + 2I (aq)  2Cl (aq) + I2 0 -1 -1 0 r o

- - Br2 + 2I (aq)  2Br (aq) + I2 0 -1 -1 0 r o

2+ - 3+ Cl2 + 2Fe (aq)  2Cl (aq) + 2Fe (aq) 0 +2 -1 +3 r o

2- - 2- Cl2 + SO3 (aq)  2Cl (aq) + SO4 (aq) 0 +4-2 -1 +6-2 r o

Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s) 0 +2 +2 0 o r

Mg(s) + CuSO4(aq)  MgSO4(aq) + Cu(s) 0 +2 +2 0 o r

Mg(s) + H2SO4(aq)  H2(g) + Mg SO4 (aq) 0 +1 0 +2 o r

Zn(s) + H2SO4(aq)  H2(g) + ZnSO4(aq) 0 +1 0 +2 o r

17 2. The iodine/thiosulfate titration to be dealt with in experiment 4.7 is another redox reaction involving a halogen. The equation for the reaction is 2Na2S2O3 + I2  2NaI + Na2S4O6 Show that the iodine is acting as an oxidising agent there as well.

2Na2S2O3 + I2  2NaI + Na2S4O6 +1 +2–2 0 +1-1 +1+2.5-2 o r The iodine is reduced, and is therefore an oxidising agent.

3. Describe the tests you would use to distinguish between sulfite and sulfate anions in aqueous solution.

Add 2 cm3 of the solution to be tested to a clean test tube. Using a dropping pipette add a few drops of barium chloride solution. A white precipitate forms. Now add 2 cm3 of dilute hydrochloric acid. The white precipitate will dissolve if a sulfite is present, and will not dissolve if a sulfate is present.

4. Explain why it is difficult to make an aqueous solution of iodine. What particular method is used to overcome this problem?

Iodine crystals dissolve very poorly in water, because iodine is non-polar. Iodine dissolves readily in an aqueous solution of potassium iodide, - - because it reacts to form I3 ions. A solution of I3 ions is always treated as - an aqueous solution of iodine, as I3 ions release I2 molecules in reactions.

5. An aqueous solution of chlorine is often made by reacting concentrated hydrochloric acid with a diluted commercial bleach solution. The active ingredient in commercial bleach is sodium hypochlorite (NaOCl). The equation for the reaction is NaOCl + 2HCl  Cl2 + NaCl + H2O Show what species are oxidised and reduced during this reaction.

NaOCl + 2HCl  Cl2 + NaCl + H2O +1-2+1 +1-1 0 +1 -1 +1-2 r o

6. The position of zinc in the Periodic Table would allow you to predict the colour of its sulfate salt solution. Explain.

Zinc is a d-block metal but it is not a transition metal. It therefore will not be expected to form coloured compounds.

18 7. Explain why magnesium is more reactive than zinc.

Magnesium atom has a larger atomic radius and a smaller nuclear charge than a zinc atom. Because of this, and despite the fact that there is extra screening of outer electrons by electrons in inner energy levels in a zinc atom compared to a magnesium atom, the outer electrons in a magnesium atom are not as tightly bound as those in an atom of zinc. Magnesium is therefore higher on the electrochemical series and more reactive than zinc.

Crystals of silver should appear on the surface of the copper wire. The +2 solution should gradually take on a blue colour (Cu (aq)).

9. Carry out some research to find out why commercial photography laboratories might have a special interest in these kinds of reactions.

Silver halides are reduced to silver when photographic film or paper is exposed to light. More silver halide is reduced when the film or paper is being developed. Unreacted silver halide is dissolved away near the end of the processing. The silver should be recovered as it is a valuable metal and would damage the environment as a waste chemical. One possible way is to reduce the metal halide by reaction with a metal higher up on the electrochemical series.