Redox Reactions: A Comprehensive Guide for CXC Chemistry

Introduction to Redox Reactions

Redox (reduction-oxidation) reactions are fundamental chemical processes that involve the transfer of electrons between chemical species. These reactions are ubiquitous in nature and are responsible for many biological and industrial processes, from cellular respiration to the extraction of metals from their ores.

A redox reaction involves both reduction (gain of electrons) and oxidation (loss of electrons) occurring simultaneously.

Basic Concepts of Redox Reactions

Oxidation and Reduction

Mnemonic: "OIL RIG" - Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons)

Oxidation Numbers (Oxidation States)

An oxidation number is a number assigned to an atom in a compound that represents the number of electrons gained, lost, or shared by that atom.

Rules for Assigning Oxidation Numbers:

Example: Calculate the oxidation number of Mn in KMnO4

K: +1 (Group 1 metal)
O: -2 (four oxygen atoms = -8)
For a neutral compound: +1 + Mn + (-8) = 0
Therefore, Mn = +7

Identifying Redox Reactions

A chemical reaction is a redox reaction if there is a change in oxidation numbers of one or more elements involved.

Example: Zn + CuSO4 → ZnSO4 + Cu

Zn: 0 → +2 (oxidation)
Cu: +2 → 0 (reduction)

Since both oxidation and reduction occur, this is a redox reaction.

Types of Redox Reactions

1. Combination (Synthesis) Reactions

Two or more substances combine to form a single product.

2Mg + O2 → 2MgO

Mg: 0 → +2 (oxidation)
O: 0 → -2 (reduction)

2. Decomposition Reactions

A single compound breaks down into two or more simpler substances.

2H2O → 2H2 + O2

O: -2 → 0 (oxidation)
H: +1 → 0 (reduction)

3. Displacement (Single Replacement) Reactions

An element replaces another element in a compound.

Zn + 2HCl → ZnCl2 + H2

Zn: 0 → +2 (oxidation)
H: +1 → 0 (reduction)

4. Disproportionation Reactions

A single element is both oxidized and reduced in the same reaction.

3Cl2 + 6NaOH → 5NaCl + NaClO3 + 3H2O

Cl: 0 → -1 (in NaCl, reduction)
Cl: 0 → +5 (in NaClO3, oxidation)

Redox Reactions in Electrochemical Cells

Galvanic (Voltaic) Cells

Galvanic cells convert chemical energy into electrical energy through spontaneous redox reactions.

Zn|Zn²⁺ Cu²⁺|Cu Salt bridge Zn Cu V Zn²⁺ Cu²⁺ Zn → Zn²⁺ + 2e⁻ Cu²⁺ + 2e⁻ → Cu

Figure 1: A Daniell Cell (Zn-Cu Galvanic Cell)

Electrolytic Cells

Electrolytic cells use electrical energy to drive non-spontaneous redox reactions.

Electrolyte Solution Anode (+) Cathode (-)

Figure 2: An Electrolytic Cell

Note the terminal polarity difference between galvanic and electrolytic cells:

Balancing Redox Reactions

Half-Reaction Method

This method separates the reaction into oxidation and reduction half-reactions, balances them separately, and then combines them.

Steps to Balance Redox Reactions in Acidic Solution:

  1. Write the unbalanced equation.
  2. Separate into half-reactions (oxidation and reduction).
  3. Balance all elements except H and O.
  4. Balance O by adding H2O.
  5. Balance H by adding H+.
  6. Balance charges by adding electrons (e-).
  7. Multiply half-reactions to equalize electrons transferred.
  8. Add half-reactions and cancel common terms.
  9. Verify that the equation is balanced.

Example: Balance the following equation in acidic solution.

MnO4- + Fe2+ → Mn2+ + Fe3+

Step 1: Separate into half-reactions

Reduction: MnO4- → Mn2+

Oxidation: Fe2+ → Fe3+

Step 2: Balance elements except H and O

Reduction: MnO4- → Mn2+ (Mn is balanced)

Oxidation: Fe2+ → Fe3+ (Fe is balanced)

Step 3: Balance O by adding H2O

Reduction: MnO4- → Mn2+ + 4H2O

Oxidation: Fe2+ → Fe3+ (no O to balance)

Step 4: Balance H by adding H+

Reduction: 8H+ + MnO4- → Mn2+ + 4H2O

Oxidation: Fe2+ → Fe3+ (no H to balance)

Step 5: Balance charges by adding electrons

Reduction: 5e- + 8H+ + MnO4- → Mn2+ + 4H2O

Oxidation: Fe2+ → Fe3+ + e-

Step 6: Multiply to equalize electrons

Reduction: 5e- + 8H+ + MnO4- → Mn2+ + 4H2O (× 1)

Oxidation: (Fe2+ → Fe3+ + e-) × 5

Oxidation: 5Fe2+ → 5Fe3+ + 5e-

Step 7: Add half-reactions

5e- + 8H+ + MnO4- + 5Fe2+ → Mn2+ + 4H2O + 5Fe3+ + 5e-

Step 8: Cancel common terms (5e-)

8H+ + MnO4- + 5Fe2+ → Mn2+ + 4H2O + 5Fe3+

Step 9: Verify balance

Atoms: Balanced

Charge: Left side: +8 - 1 + 10 = +17, Right side: +2 + 0 + 15 = +17

The equation is balanced.

Balancing in Basic Solution

For basic solutions, follow these additional steps after balancing in acidic solution:

  1. Add OH- to both sides of the equation to neutralize H+.
  2. Combine H+ and OH- to form H2O.
  3. Cancel common H2O molecules.

Standard Electrode Potentials

Standard Hydrogen Electrode (SHE)

The SHE is the reference electrode with a standard electrode potential of 0.00 V.

Convention: The more positive the standard electrode potential (E°), the greater the tendency of the substance to be reduced (gain electrons).

Electrochemical Series

Half-Reaction E° (V)
F2 + 2e- → 2F- +2.87
Cl2 + 2e- → 2Cl- +1.36
O2 + 4H+ + 4e- → 2H2O +1.23
Br2 + 2e- → 2Br- +1.07
Ag+ + e- → Ag +0.80
Fe3+ + e- → Fe2+ +0.77
I2 + 2e- → 2I- +0.54
Cu2+ + 2e- → Cu +0.34
2H+ + 2e- → H2 0.00
Pb2+ + 2e- → Pb -0.13
Ni2+ + 2e- → Ni -0.25
Fe2+ + 2e- → Fe -0.44
Zn2+ + 2e- → Zn -0.76
Al3+ + 3e- → Al -1.66
Mg2+ + 2e- → Mg -2.37
Na+ + e- → Na -2.71
K+ + e- → K -2.93
Li+ + e- → Li -3.05

Using the Electrochemical Series

Cell Potential (Electromotive Force, EMF)

The cell potential (Ecell) is the difference between the reduction potential at the cathode and the reduction potential at the anode.

Ecell = Ecathode - Eanode

For a Zn-Cu galvanic cell:

Anode (oxidation): Zn → Zn2+ + 2e- (E° = -0.76 V)

Cathode (reduction): Cu2+ + 2e- → Cu (E° = +0.34 V)

Ecell = Ecathode - Eanode = 0.34 V - (-0.76 V) = 1.10 V

Applications of Redox Reactions

Batteries and Fuel Cells

Batteries and fuel cells convert chemical energy into electrical energy through redox reactions.

Corrosion

Corrosion is the oxidation of metals in the presence of oxygen and moisture, resulting in their deterioration.

Rusting of Iron:

Anode reaction: Fe → Fe2+ + 2e-

Cathode reaction: O2 + 4H+ + 4e- → 2H2O

Fe2+ ions are further oxidized to Fe3+ and form iron(III) oxide (Fe2O3·nH2O), commonly known as rust.

Methods to Prevent Corrosion:

Iron (Fe) Zinc (Zn) coating H₂O + O₂ e⁻ Zn → Zn²⁺ + 2e⁻ O₂ + 2H₂O + 4e⁻ → 4OH⁻

Figure 3: Galvanic Protection of Iron with Zinc

Electroplating

Electroplating is a process that uses electricity to deposit a thin layer of metal onto a conducting surface.

Electroplating Silver onto Copper:

Anode (oxidation): Ag → Ag+ + e- (silver electrode dissolves)

Cathode (reduction): Ag+ + e- → Ag (silver deposits on copper)

Extraction of Metals

Redox reactions are used to extract metals from their ores through processes like smelting and electrolysis.

Extraction of Aluminum by Electrolysis:

Cathode: Al3+ + 3e- → Al

Anode: 2O2- → O2 + 4e-

Biological Redox Processes

Redox reactions are essential for many biological processes, including cellular respiration and photosynthesis.

Practical Experiments

Experiment 1: Metal Reactivity Series

Aim:

To investigate the reactivity of various metals with metal salt solutions and establish a reactivity series.

Materials:

Procedure:

  1. Clean each metal strip with sandpaper to expose a fresh surface.
  2. Place each metal strip in separate test tubes containing each metal salt solution.
  3. Leave for 10-15 minutes and observe any changes.
  4. Record your observations in a table.

Expected Results:

A more reactive metal will displace the metal ion of a less reactive metal from its salt solution. For example, zinc will displace copper from copper(II) sulfate solution, producing a copper coating on the zinc strip and zinc sulfate in solution.

Conclusion:

Based on the displacement reactions observed, you can arrange the metals in order of decreasing reactivity (most reactive to least reactive).

Experiment 2: Building a Simple Galvanic Cell

Aim:

To construct a simple galvanic cell and measure its voltage.

Materials:

Procedure:

  1. Fill one beaker with zinc sulfate solution and place the zinc strip in it.
  2. Fill the other beaker with copper(II) sulfate solution and place the copper strip in it.
  3. Connect the beakers with the salt bridge (filter paper soaked in sodium chloride solution).
  4. Connect the zinc strip to the negative terminal of the voltmeter and the copper strip to the positive terminal.
  5. Record the voltage reading.

Expected Results:

The voltmeter should show a reading of approximately 1.10 V, indicating the flow of electrons from the zinc electrode (anode) to the copper electrode (cathode).

Conclusion:

This experiment demonstrates how a galvanic cell converts chemical energy into electrical energy through a spontaneous redox reaction.

Experiment 3: Determining the Oxidizing Power of Halogens

Aim:

To investigate the relative oxidizing power of halogens by observing displacement reactions.

Materials:

Procedure:

  1. Add a few drops of cyclohexane to each test tube.
  2. Prepare various combinations of halogen solutions and halide solutions.
  3. Shake each test tube gently and allow the layers to separate.
  4. Observe the color in the organic layer.

Expected Results:

Chlorine will displace bromine and iodine from their salts. Bromine will displace iodine but not chlorine. Iodine will not displace either chlorine or bromine.

The colors in the organic layer will indicate which halogen has been displaced:

Conclusion:

The experiment demonstrates the decreasing oxidizing power in the order: Cl2 > Br2 > I2, which corresponds to their positions in the electrochemical series.

Glossary of Terms

Redox reaction: A chemical reaction in which electrons are transferred between substances, resulting in changes in oxidation states.
Oxidation: The loss of electrons or an increase in oxidation state by a molecule, atom, or ion.
Reduction: The gain of electrons or a decrease in oxidation state by a molecule, atom, or ion.
Oxidation number/state: A number assigned to an atom in a compound that represents the number of electrons gained, lost, or shared by that atom.
Oxidizing agent: A substance that causes another substance to be oxidized (lose electrons) while being reduced itself.
Reducing agent: A substance that causes another substance to be reduced (gain electrons) while being oxidized itself.
Half-reaction: Either the oxidation or reduction reaction component of a redox reaction.
Electrochemical cell: A device that generates electrical energy from chemical reactions or uses electrical energy to drive chemical reactions.
Galvanic (Voltaic) cell: An electrochemical cell that produces electricity from spontaneous redox reactions.
Electrolytic cell: An electrochemical cell that uses electrical energy to drive non-spontaneous redox reactions.
Anode: The electrode where oxidation occurs.
Cathode: The electrode where reduction occurs.
Salt bridge: A device that connects the two half-cells of a galvanic cell, allowing ion flow while preventing the direct mixing of the electrolyte solutions.
Standard electrode potential (E°): The potential of a half-cell measured against the standard hydrogen electrode under standard conditions.
Cell potential (Ecell): The potential difference between two half-cells of an electrochemical cell.
Electrochemical series: A ranking of chemical species according to their standard electrode potentials.
Corrosion: The deterioration of metals due to oxidation reactions with their environment.
Electroplating: A process that uses electrical current to reduce dissolved metal cations to form a coating on an electrode.
Disproportionation: A redox reaction where a single substance is both oxidized and reduced.
Standard Hydrogen Electrode (SHE): The reference electrode with a standard electrode potential of 0.00 V.

Self-Assessment Questions

Multiple Choice Questions

1. In the reaction: MnO4- + C2O42- → Mn2+ + CO2, what is the change in oxidation number of manganese?

  1. +7 to +2
  2. +7 to +4
  3. +4 to +2
  4. +2 to +7

Answer: a) +7 to +2

Explanation: In MnO4-, the oxidation number of Mn is +7. In Mn2+, the oxidation number is +2. Therefore, the change is from +7 to +2, which means Mn is reduced by gaining 5 electrons.

2. Which of the following is NOT correctly matched?

  1. Oxidation - Loss of electrons
  2. Reduction - Decrease in oxidation number
  3. Oxidizing agent - Undergoes reduction
  4. Reducing agent - Loses electrons

Answer: d) Reducing agent - Loses electrons

Explanation: A reducing agent causes reduction in another substance and is itself oxidized. During oxidation, a substance loses electrons. Therefore, it is correct to say that a reducing agent loses electrons. However, this is not the definition of a reducing agent but rather what happens to it during the reaction.

3. In an electrolytic cell:

  1. The anode is positive and the cathode is negative
  2. The anode is negative and the cathode is positive
  3. Both electrodes are positively charged
  4. Both electrodes are negatively charged

Answer: a) The anode is positive and the cathode is negative

Explanation: In an electrolytic cell, the anode is connected to the positive terminal of the external power supply, making it positive. The cathode is connected to the negative terminal, making it negative. This is opposite to the polarity in a galvanic cell.

4. The oxidation number of sulfur in H2SO4 is:

  1. +2
  2. +4
  3. +6
  4. +8

Answer: c) +6

Explanation: In H2SO4, the oxidation numbers are: H = +1 (×2), O = -2 (×4), and S = x.

For a neutral compound: 2(+1) + x + 4(-2) = 0

2 + x - 8 = 0

x = 6

Therefore, the oxidation number of sulfur is +6.

5. Which of the following is a redox reaction?

  1. NaOH + HCl → NaCl + H2O
  2. BaCl2 + Na2SO4 → BaSO4 + 2NaCl
  3. Zn + CuSO4 → ZnSO4 + Cu
  4. CaCO3 → CaO + CO2

Answer: c) Zn + CuSO4 → ZnSO4 + Cu

Explanation: In this reaction, Zn is oxidized from an oxidation state of 0 to +2, while Cu2+ is reduced from +2 to 0. Since there are changes in oxidation states, this is a redox reaction. The other reactions are either acid-base, precipitation, or decomposition reactions without changes in oxidation states.

Short Answer Questions

6. Balance the following redox equation in acidic solution:

Cr2O72- + Fe2+ → Cr3+ + Fe3+

Answer: Cr2O72- + 14H+ + 6Fe2+ → 2Cr3+ + 6Fe3+ + 7H2O

Step-by-step solution:

1. Split into half-reactions:

Reduction: Cr2O72- → 2Cr3+

Oxidation: Fe2+ → Fe3+

2. Balance Cr:

Cr2O72- → 2Cr3+ (already balanced)

3. Balance O by adding H2O:

Cr2O72- → 2Cr3+ + 7H2O

4. Balance H by adding H+:

Cr2O72- + 14H+ → 2Cr3+ + 7H2O

5. Balance charges by adding electrons:

Cr2O72- + 14H+ + 6e- → 2Cr3+ + 7H2O

Fe2+ → Fe3+ + e-

6. Multiply to equalize electrons:

Cr2O72- + 14H+ + 6e- → 2Cr3+ + 7H2O (×1)

(Fe2+ → Fe3+ + e-) × 6

6Fe2+ → 6Fe3+ + 6e-

7. Combine half-reactions and cancel common terms:

Cr2O72- + 14H+ + 6Fe2+ → 2Cr3+ + 6Fe3+ + 7H2O

7. Calculate the cell potential for a galvanic cell with a zinc electrode in a 1.0 M Zn2+ solution and a copper electrode in a 1.0 M Cu2+ solution at 25°C. (E°Zn2+/Zn = -0.76 V, E°Cu2+/Cu = +0.34 V)

Answer: Ecell = 1.10 V

Explanation:

For a galvanic cell, Ecell = Ecathode - Eanode

Cathode (reduction): Cu2+ + 2e- → Cu (E° = +0.34 V)

Anode (oxidation): Zn → Zn2+ + 2e- (E° = -0.76 V)

Ecell = +0.34 V - (-0.76 V) = 1.10 V

Since both solutions are at standard conditions (1.0 M), the cell potential equals the standard cell potential.

8. Explain why copper cannot displace hydrogen from dilute acids, while zinc can.

Answer:

The ability of a metal to displace hydrogen from acids depends on its position in the electrochemical series relative to hydrogen.

For a metal to displace hydrogen from an acid, it must have a more negative electrode potential than hydrogen, meaning it must be a stronger reducing agent than hydrogen.

Looking at standard electrode potentials:

- Zn2+/Zn: E° = -0.76 V

- H+/H2: E° = 0.00 V

- Cu2+/Cu: E° = +0.34 V

Since zinc has a more negative electrode potential than hydrogen, it can reduce H+ ions to H2 gas while being oxidized to Zn2+. The reaction Zn + 2H+ → Zn2+ + H2 is spontaneous with a positive Ecell value of 0.76 V.

Copper, however, has a more positive electrode potential than hydrogen, meaning it is a weaker reducing agent than hydrogen. The reaction Cu + 2H+ → Cu2+ + H2 would have a negative Ecell value of -0.34 V, indicating that it is non-spontaneous. Therefore, copper cannot displace hydrogen from dilute acids.

9. Describe the process of corrosion of iron and explain one method to prevent it.

Answer:

Process of Corrosion of Iron (Rusting):

Rusting is an electrochemical process that requires the presence of water and oxygen. It occurs in these steps:

  1. Anodic Reaction (Oxidation): At the anode site, iron is oxidized to iron(II) ions, releasing electrons:
    Fe → Fe2+ + 2e-
  2. Cathodic Reaction (Reduction): At the cathode site, oxygen is reduced in the presence of water and electrons:
    O2 + 4H+ + 4e- → 2H2O
  3. Further Oxidation: Fe2+ ions are further oxidized by oxygen to form iron(III) ions:
    4Fe2+ + O2 + 4H+ → 4Fe3+ + 2H2O
  4. Formation of Rust: Fe3+ ions combine with hydroxide ions (from water) to form iron(III) hydroxide, which then dehydrates to form iron(III) oxide (rust):
    Fe3+ + 3OH- → Fe(OH)3 → Fe2O3·nH2O (rust)

Method to Prevent Corrosion:

Galvanization is an effective method to prevent iron from corroding. It involves coating iron or steel with a layer of zinc. Zinc acts as a sacrificial metal because it is more reactive than iron (more negative electrode potential). Even if the zinc coating is scratched and the iron is exposed, zinc will still preferentially oxidize (corrode) instead of iron. This is called cathodic protection. The zinc forms a protective layer of zinc oxide/zinc carbonate that further slows down the corrosion process.

10. Identify the oxidizing agent and reducing agent in the following reaction and explain your answer:

2KMnO4 + 5H2C2O4 + 3H2SO4 → K2SO4 + 2MnSO4 + 10CO2 + 8H2O

Answer:

Oxidizing Agent: KMnO4 (potassium permanganate)

Reducing Agent: H2C2O4 (oxalic acid)

Explanation:

To identify the oxidizing and reducing agents, we need to track the changes in oxidation numbers:

In KMnO4, the oxidation number of Mn is +7 (K = +1, O = -2, and the compound is neutral).

In MnSO4, the oxidation number of Mn is +2 (S = +6, O = -2, and the compound is neutral).