Experiment No by 93l430Nn

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									OUR LADY OF THE ROSARY COLLEGE CHEMISTRY PRACTICAL MANUAL
   Advanced Level, F.6
   Experiment should be carried out under the supervision of a Chemistry teacher
   Safety precaution must be noted.
Division of work: individual

                     A test tube study of redox reactions
Introduction
In this practical you will attempt to place a number of redox pairs in order of oxidizing strength by
carrying out suitable experiments. The redox pairs concerned are
        I      I2 + 2e  2I-
        II SO42- + H+ + e  H2SO32- + H2O
        III ClO- + H2O + e  Cl- + OH-
        IV Br2 + 2e  2Br-
        V     Fe3+ + e  Fe2+
Tabulate your test result and immediate deduction. Write a balanced ionic equation for the reaction
in each experiment. Place equations from I to V in order of oxidizing strength
Pre-laboratory work
1. Define 'reduction reaction', 'oxidation reaction' and 'redox pair'.
2. A redox reaction may be represented by an oxidation half-equation and a reduction
    half-equation. Write half-equations for the redox reaction of iron(II) and silver (I) ions.
3. A textbook states "redox equilibrium is the competition for electrons…". Use your own words
    to elaborate this idea. Appropriate examples should be given.
4. Many reduction half-reactions are arranged in the electrochemical series (E.C.S.). How are
    they arranged?
Chemicals/materials:
   universal indicator and wooden splints.
Bench solution:
    iron(III) chloride, potassium chloride, sulphuric acid, sodium chlorate(I), iron(II) sulphate,
    iodine in KI, starch, hydrogen peroxide, solution of SO2 , bromine water, chlorine water and
    sodium bromide solution.
Procedure
Experiment 1
Add a little solution of iron(III) ions to a solution of sulphate(IV) ions in a test tube. Add also a
little solution of sulphate(VI) to a solution of iron(II) ions in another test tube. Describe any
chemical change and determine the relative strength of the redox pairs.


Experiment 2
Add a little bromine water to a solution containing iron(II) ions. Also add a little solution of iron(III)
into a solution containing bromide ion. Describe any chemical change and determine the relative
strength of the two redox pairs.


Experiment 3
Add a little iodine solution to a solution containing iron(II) ions in a test tube. Add also a little
solution containing iron(III) ions into a solution of iodide solution in another test tube. Describe
whether any chemical change and determine the relative strength of the redox pairs.


Experiment 4
Add a little iodine solution to a solution containing sulphate(IV) ions in a test tube. Add also a little
solution containing sulphate(VI) ions into a solution of iodide ions in another test tube. Describe
whether any chemical change and determine the relative strength of the redox pairs.


Experiment 5
Add a little iodide solution to a solution containing chlorate(I) ions in a test tube. Add also a little
solution containing chloride ions into a solution of iodine in another test tube. Describe whether
any chemical change and determine the relative strength of the redox pairs.


Experiment 6
Add a little bromide solution to a solution containing chlorate(I) ions in a test tube. Add also a little
solution containing chloride ions into a solution of bromine in another test tube. Describe whether
any chemical change and determine the relative strength of the redox pairs.


Experiment 7 and 8
Hydrogen peroxide, H2O2, can behave as both an oxidizer and a reducer depending on the
conditions, Half-equation VI shows hydrogen peroxide behaving as an oxidizing agent.
                    VI      H2O2 + 2H+ + 2e  2H2O
Half-equation VII shows hydrogen peroxide behaving as a reducing agent.
                    VII      O2 + 2H+ + 2e  H2O2
When hydrogen peroxide acts as a reducing agent half-equation VII will, of course, proceed in the
reverse direction.


Experiment 7
Add a little hydrogen peroxide solution to a solution of sodium chlorate(I) and note what happens.
Decide whether the hydrogen peroxide is behaving as an oxidizer or as a reducer. Give your result
and reasons.


Experiment 8
Acidify a little potassium iodide solution with dilute sulphuric(VI) acid. Add a
little hydrogen peroxide solution. Describe what occurs.


Further work
Devise tests to establish more precisely the positions of half-equation VII and VIII in the table.


Questions for discussion
Questions 1-4: Experiment 1
1. Half-equations II and III are not balanced. Answer these questions with reference to each of II
    and III.
    (a) Which element is undergoing redox?
    (b) Work out the oxidation number of this element in both its oxidized and reduced forms
         and
    (c) Balance the half-equation.
2. For experiment 1-6:
    (a) Write a balanced ionic equation for the reaction in each experiment.
    (b) Place equations from I to V in order of oxidizing strength


Answer questions 3-5 by using half-equations I and V in a tabular form.
3. Write a balanced ionic equation for the reaction that has occurred.
4. In this reaction, which of the half-equations moved from right to left instead of left to right?
5. Place equations I and V in order of oxidizing strength.
For experiment 7 and 8:
6. Which half-equation, VI or VII describes the behaviour of hydrogen peroxide in (a) experiment
   7 and (b) experiment 8?
7. Which position should equation VI be placed?
8. Which position should equation VII be placed ?
9. Write balanced ionic equations for the reaction of hydrogen peroxide in experiment 7 and 8.


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