SCH4U1
Name: ________________
OXIDATION-REDUCTION REACTIONS
BACKGROUND:
In Part 1 of this experiment, you will observe some possible oxidation-reduction reactions
involving several metals and metallic ions. The results can be used to estimate the ease of
reduction of the metallic ions relative to each other. In Part 2, you will make a similar
comparison of the relative ease of reduction of the three halogen elements C12, Br2 and I2.
Specifically, you will determine which of these halogens is capable of removing electrons from
which of the halide ions, Cl− (aq), Br− (aq) or I− (aq). From this information, you will be able to
arrange the halogen element-halide ion half-reactions in order of decreasing ease of reduction.
For example, magnesium metal, Mg(s), is oxidized when it is placed in a solution of lead (II)
nitrate. In the process, Mg (s) loses electrons to form Mg2+ ions while Pb2+ ions are reduced to
lead metal, Pb (s). Thus, in this reaction, there is a transfer of electrons from Mg(s) to Pb2+ (aq).
The Mg(s) is the reducing agent because it reduced the Pb2+ (aq) ion. The Pb2+ (aq) ion is the
oxidizing agent because it oxidized the Mg(s).
BCE:
Pb(NO3)2 (aq) +
Mg (s)
Pb (s) +
Mg(NO3)2
Oxidation:
Mg (s)
Mg2+ (aq)
+
2e−
Reduction:
Pb2+ (aq)
+
2e−
Pb (s)
Net reaction:
Pb2+ (aq)
+
Mg (s)
Pb (s)
+
Mg2+ (aq)
PURPOSE:
To establish the relative ease of reduction of three metallic ions and three halogens.
HYPOTHESIS:
Based on what you know about the trend in electronegativity in the periodic table, predict the
order of the ease of reduction for the metal ions and non-metallic elements listed below.
MATERIALS:
Part 1: - 2 copper strips
- 2 zinc strips
- 2 lead strips
- 6 large test tubes
- test tube rack
- steel wool
- 0.2 mol/L Zn(NO3)2, Cu(NO3)2 and Pb(NO3)2 solutions
Part 2: - 2 small test tubes
- 2 rubber stoppers
- mineral oil
- chlorine water
- bromine water
- iodine water
- 0.5 mol/L KBr solution
- 0.5 mol/L KI solution
- 0.5 mol/L NaCl solution
SAFETY:
Wear disposable gloves and goggles during the experiment. Lead ions are very toxic. Do not
handle solutions of this substance without permission. Copper (II) nitrate and zinc nitrate are
toxic by ingestion. Halogens solutions can produce very reactive vapours that can damage the
respiratory system. Always use open containers containing these elements in the fume hood.
Wash hands after the lab is completed.
METHOD:
Part 1 - Reduction of Metallic Ions by Metals
1) Observe the prepared reaction combinations of copper, zinc and lead metal in copper (II)
nitrate, lead (II) nitrate and zinc nitrate solutions. Do not touch solutions containing lead ions.
2) Record any observations of a chemical reaction or write “no reaction”.
Part 2 - Reduction of Halogens by Halide Ions
1) Fill one of the small test tubes to a height of about 2 cm with 0.1 mol/L KBr solution.
2) Fill the other small test tube to a height of about 2 cm with 0.5 mol/L KI solution.
3) Add 1 half-filled dropper of mineral oil to each test tube. Then add 2 droppers of a fresh
solution of chlorine in water to each test tube.
4) Stopper and shake both test tubes for a few seconds. Note the colour of the mineral oil (nonpolar) phase.
5) Dispose of the solutions in the “organic waste” container.
6) Repeat Steps 1-5 using NaCl solution, KI solution and bromine water.
7) Repeat Steps 1-5 using NaCl solution, KBr solution and iodine water.
8) Using the color chartin Table 1, interpret which halogens and halide ions were present in each
experiment. For example, a yellow-brown mineral oil phase means bromine was present and
pink-purple means that iodine was present.
Halogen or
C12
Br2
I2
Cl−
Halide Present
Colour of
no
yellowpinkno
Mineral Oil
colour
golden
purple
colour
Phase
colour
colour
Table 1: Colour of various halogens and halide ions in mineral oil.
Br−
I−
no
colour
no
colour
OBSERVATIONS:
Create two data tables (Table 2: Reduction of Metal Ions and Table 3: Reduction of Halogens)
to record your observations for Part 1 and Part 2.
ANALYSIS:
Part 1 [T15; C1]
1) For each of the reactions that occurred in Part 1, write a balanced single displacement
reaction. Include the state symbols for each reactant. [T3; C1]
2) For each reaction that occurred, write the net ionic equation and indicate the direction of
electron movement using an arrow. [T3]
3) Which of the metallic ions was reduced by: a) two of the metals b) one of the metals c) none
of the metals? [T1]
4) For each net ionic equation from question 2,, write the oxidation and reduction half reactions.
[T6]
5) Arrange the metallic ion-metal reduction half-reactions in a column in order of decreasing
ease of reduction (i.e. the metallic ion that was always reduced is at the top). Other
experiments show that the Ag+ (aq) ion is reduced by Zn (s), Pb (s) and Cu(s). Use this
information to add the half-reaction Ag+ (aq) + e− → Ag(s) to your list in the proper place. [T2]
Part 2 [T15; C1]
1) For each of the reactions that occurred in Part 2, write a balanced single displacement
reaction. Include the state symbols for each reactant. [T3; C1]
2) For each reaction that occurred, write the net ionic equation and indicate the direction of
electron movement using an arrow. [T3]
3) Which of the halogens tested was reduced by a) two of the halide ions b) one of the halide
ions c) none of the halide ions? [T1]
4) For each net ionic equation from question 2, write the oxidation and reduction half reactions.
[T6]
5) Arrange the halogen element-halide ion half-reactions in a column in order of decreasing ease
of reduction (i.e. the halogen that was always reduced is at the top). [T2]
CONCLUSION [T2]:
Restate the purpose and evaluate the hypothesis. In the conclusion, construct a series of all
seven half-reactions discussed in this experiment. List them in order of decreasing ease of
reduction, given the following information: Ag+ (aq) ions are less easily reduced than Br2(aq)
and more easily reduced than I2(aq); also I2(aq) is less easily reduced than Ag+(aq) and more
easily reduced than Cu2+ (aq). [T2]
QUESTIONS [A7]:
Use the list you constructed in the conclusion to answer the following:
1) Would it be possible to store a solution of copper(II) sulfate in a container made of metallic
zinc ? Explain why or why not. [A2]
2) Would it be possible to store a solution of copper(II) sulfate in a container made of metallic
silver? Explain. [A2]
3) Would you expect jewelry made from an alloy of silver and copper to tarnish (oxidize) in a
laboratory where fumes of bromine are present? Explain your reasoning. [A2]
4) Alloys of copper and zinc are used to produce brass. Why do these elements not react in this
application? [A1]