Electrochemistry: A Comprehensive Study Guide for CXC Chemistry

This study guide covers the electrochemistry section of the CXC Chemistry syllabus for 2024-2025. It includes definitions, key concepts, experimental procedures, and self-assessment questions with answers.

Introduction to Electrochemistry

Electrochemistry is the branch of chemistry that studies the relationship between electricity and chemical reactions. It involves the study of chemical changes caused by electrical energy and the production of electrical energy by chemical reactions.

Key Concepts in Electrochemistry

Oxidation and Reduction

Redox reactions are the foundation of electrochemistry. Understanding the basic concepts of oxidation and reduction is crucial.

Defining Oxidation and Reduction

A helpful mnemonic to remember: OIL RIG - Oxidation Is Loss, Reduction Is Gain of electrons.

Identifying Oxidation and Reduction

Examples of Redox Reactions

Consider the reaction between zinc and copper(II) sulfate:

Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)

Zn Cu Zn²⁺ Cu²⁺ electrons

Figure 1: Electron transfer in a Zn-Cu redox reaction

Electrolytes and Electrolysis

Electrolytes

Electrolytes are substances that conduct electricity when dissolved in water or when melted, due to the presence of ions.

Electrolysis

Electrolysis is the process of using electrical energy to drive a non-spontaneous chemical reaction. It involves the decomposition of an electrolyte by passing an electric current through it.

Components of an Electrolytic Cell

Anode (+) Cathode (-) Electrolyte DC Power Anions Cations

Figure 2: Basic setup of an electrolytic cell

Important Electrolysis Processes

Electrolysis of Water

2H2O(l) → 2H2(g) + O2(g)

Electrolysis of Aqueous Sodium Chloride

2NaCl(aq) + 2H2O(l) → 2NaOH(aq) + H2(g) + Cl2(g)

In aqueous electrolysis, water can be preferentially discharged over certain ions. This depends on the position of ions in the electrochemical series and their concentration.

Electrolysis of Copper(II) Sulfate

With inert electrodes (carbon):

With copper electrodes:

Experiment: Electroplating with Copper

Aim: To electroplate an object with copper.

Materials:

Procedure:

  1. Connect the copper anode to the positive terminal of the power supply
  2. Connect the object to be plated to the negative terminal
  3. Immerse both electrodes in copper(II) sulfate solution
  4. Turn on the power supply at a low voltage (about 2-6V)
  5. Allow the process to continue for 15-20 minutes
  6. Remove the plated object, rinse with water, and dry

Observations: A layer of copper forms on the object connected to the negative terminal. The copper anode gradually dissolves.

Explanation: Copper ions from the solution are reduced and deposited on the cathode. At the anode, copper atoms are oxidized to copper ions, maintaining the concentration of the solution.

Faraday's Laws of Electrolysis

  1. First Law: The mass of a substance deposited or liberated at an electrode is directly proportional to the quantity of electricity passed.
  2. Second Law: The masses of different substances deposited or liberated by the same quantity of electricity are proportional to their chemical equivalent weights.

The mathematical form of Faraday's Law:

m = (M × I × t) / (n × F)

Where:

The Faraday constant (F) represents the charge on one mole of electrons and equals 96,500 coulombs/mol.

Electrochemical Cells

Electrochemical cells convert chemical energy into electrical energy (galvanic/voltaic cells) or electrical energy into chemical energy (electrolytic cells).

Galvanic (Voltaic) Cells

Galvanic cells produce electricity from spontaneous redox reactions.

Components of a Galvanic Cell

Note that in galvanic cells, the anode is negative because it supplies electrons, while in electrolytic cells, the anode is positive.

V Zn Anode (-) Cu Cathode (+) ZnSO₄ CuSO₄ Salt Bridge e⁻

Figure 3: Zn-Cu galvanic cell (Daniell cell)

Standard Electrode Potentials (E°)

The standard electrode potential is the potential difference developed by a half-cell compared to the standard hydrogen electrode (SHE) at standard conditions (25°C, 1 atm, 1 M concentration).

Half-Cell 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
Sn2+ + 2e- → Sn -0.14
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

Cell Potential (EMF)

The cell potential or electromotive force (EMF) of a galvanic cell is calculated as:

Ecell = Ecathode - Eanode

For a Zn-Cu cell:

Ecell = ECu²⁺/Cu - EZn²⁺/Zn = +0.34 - (-0.76) = +1.10 V

Practical Applications of Electrochemistry

Commercial Batteries

Electroplating

Electroplating is the process of coating an electrically conductive object with a layer of metal using electrical current. Applications include:

Electrolytic Refining

Electrolytic refining is used to purify metals. For example, in copper refining:

Electrochemical Corrosion

Corrosion is an electrochemical process in which metals are oxidized by environmental agents. Rust formation on iron is a common example:

4Fe(s) + 3O2(g) + 2H2O(l) → 2Fe2O3·H2O(s) (rust)

Corrosion Prevention Methods

Zn Seawater Iron Hull e⁻ Sacrificial anode Protected metal

Figure 4: Cathodic protection using a sacrificial zinc anode

Nernst Equation and Electrochemical Equilibrium

The Nernst equation relates the cell potential to the standard cell potential and the reaction quotient:

E = E° - (RT/nF)ln Q

At 25°C, this simplifies to:

E = E° - (0.0592/n)log Q

Where:

Effect of Concentration on Cell Potential

The relationship between E° and Gibbs free energy: ΔG° = -nFE°. A positive E° indicates a spontaneous reaction (negative ΔG°).

Electrochemical Applications in Industry

Chlor-Alkali Process

The chlor-alkali process is used to produce chlorine gas and sodium hydroxide by the electrolysis of brine (concentrated sodium chloride solution).

Overall reaction: 2NaCl(aq) + 2H2O(l) → Cl2(g) + H2(g) + 2NaOH(aq)

Extraction of Aluminum

Aluminum is extracted from bauxite ore by the Hall-Héroult process, which involves electrolysis of aluminum oxide dissolved in molten cryolite.

Al2O3(l) → 2Al(l) + 3/2 O2(g)

Electrochemical Sensors

Electroplating Techniques

Electroplating involves depositing a metal layer on an object to improve appearance, prevent corrosion, or enhance properties.

Experiment: Determining an Unknown Metal by Electroplating

Aim: To determine the identity of an unknown metal using Faraday's laws of electrolysis.

Materials:

Procedure:

  1. Weigh the cathode before the experiment
  2. Set up the electrolysis apparatus with the unknown solution
  3. Pass a known current for a fixed time period
  4. Remove the cathode, rinse with distilled water, dry, and weigh
  5. Calculate the mass of metal deposited
  6. Using Faraday's law, calculate the molar mass of the metal
  7. Compare with known values to identify the metal

Calculations:

Using m = (M × I × t) / (n × F), rearrange to find M (molar mass)

M = (m × n × F) / (I × t)

Electrochemical Series and Predictions

Using the Electrochemical Series

The electrochemical series can be used to predict:

Predicting Chemical Reactions

Rules for predicting reactions:

Displacement Reactions

Examples of displacement reactions:

The more reactive metal (more negative electrode potential) will always displace the less reactive metal (more positive electrode potential) from its salt solution.

Glossary of Electrochemistry Terms

Electrochemistry
The branch of chemistry that deals with the relationship between electrical energy and chemical reactions.
Oxidation
The loss of electrons by a species, resulting in an increase in oxidation state.
Reduction
The gain of electrons by a species, resulting in a decrease in oxidation state.
Redox Reaction
A chemical reaction involving the transfer of electrons between species, resulting in changes in oxidation states.
Electrolyte
A substance that conducts electricity when dissolved in water or when melted, due to the presence of ions.
Electrolysis
The process of using electrical energy to drive a non-spontaneous chemical reaction.
Electrode
A conductor through which electric current enters or leaves an electrolyte in an electrochemical cell.
Anode
The electrode at which oxidation occurs. In a galvanic cell, it is the negative electrode; in an electrolytic cell, it is the positive electrode.
Cathode
The electrode at which reduction occurs. In a galvanic cell, it is the positive electrode; in an electrolytic cell, it is the negative electrode.
Galvanic (Voltaic) Cell
An electrochemical cell that converts chemical energy into electrical energy through spontaneous redox reactions.
Electrolytic Cell
An electrochemical cell that uses electrical energy to drive a non-spontaneous redox reaction.
Standard Electrode Potential (E°)
The potential difference developed by a half-cell compared to the standard hydrogen electrode under standard conditions.
Cell Potential (EMF)
The potential difference between the cathode and anode in an electrochemical cell.
Faraday's Laws of Electrolysis
Laws that relate the amount of substance produced during electrolysis to the quantity of electricity passed through the cell.
Faraday Constant (F)
The charge carried by one mole of electrons, equal to 96,500 coulombs per mole.
Salt Bridge
A device containing an electrolyte that connects the two half-cells in a galvanic cell, allowing ion flow while preventing mixing of the solutions.
Electroplating
The process of coating an object with a thin layer of metal using electrolysis.
Corrosion
The electrochemical degradation of metals due to reactions with their environment.
Nernst Equation
An equation that relates the cell potential to the standard cell potential and the reaction quotient under non-standard conditions.
Half-Cell
A system consisting of an electrode and an electrolyte, representing either the oxidation or reduction half of a redox reaction.
Electrochemical Series
A list of half-cells arranged in order of their standard electrode potentials, used to predict the direction of redox reactions.

Self-Assessment Questions

Multiple Choice Questions

1. In a galvanic cell, the anode is:
  1. Positively charged and where oxidation occurs
  2. Negatively charged and where oxidation occurs
  3. Positively charged and where reduction occurs
  4. Negatively charged and where reduction occurs

Answer: b. Negatively charged and where oxidation occurs

Explanation: In a galvanic cell, oxidation (loss of electrons) occurs at the anode, making it negatively charged as it supplies electrons to the external circuit.

2. During the electrolysis of aqueous copper(II) sulfate using inert electrodes, what is produced at the cathode?
  1. Copper metal
  2. Hydrogen gas
  3. Oxygen gas
  4. Sulfate ions

Answer: a. Copper metal

Explanation: At the cathode, Cu²⁺ ions are reduced to copper metal: Cu²⁺ + 2e⁻ → Cu

3. Which of the following is NOT a correct statement about the electrochemical series?
  1. Metals with more negative electrode potentials are stronger reducing agents
  2. Metals with more positive electrode potentials will displace metals with more negative potentials from their salt solutions
  3. Metals above hydrogen in the series will react with acids to produce hydrogen gas
  4. The standard hydrogen electrode is assigned a potential of 0.00 V

Answer: b. Metals with more positive electrode potentials will displace metals with more negative potentials from their salt solutions

Explanation: The opposite is true. Metals with more negative electrode potentials (more reactive) will displace metals with more positive potentials (less reactive) from their salt solutions.

4. When calculating the amount of substance produced during electrolysis, which of the following is NOT required?
  1. Current (in amperes)
  2. Time (in seconds)
  3. Number of electrons transferred
  4. Temperature (in Kelvin)

Answer: d. Temperature (in Kelvin)

Explanation: Faraday's law relates the mass of substance produced to the quantity of electricity passed (I × t), the molar mass, and the number of electrons transferred. Temperature is not directly involved in this calculation.

5. The process of galvanization involves:
  1. Coating iron with tin
  2. Coating iron with zinc
  3. Coating iron with copper
  4. Coating iron with chromium

Answer: b. Coating iron with zinc

Explanation: Galvanization is the process of coating iron or steel with a layer of zinc to protect it from corrosion. The zinc acts as a sacrificial anode, protecting the iron even if the coating is scratched.

Short Answer Questions

6. Explain why zinc can displace copper from copper(II) sulfate solution, but copper cannot displace zinc from zinc sulfate solution.

Answer: Zinc has a more negative standard electrode potential (E° = -0.76 V) than copper (E° = +0.34 V), which means zinc is a stronger reducing agent than copper. In a displacement reaction, the stronger reducing agent (zinc) can give electrons to the ions of the weaker reducing agent (copper ions), resulting in the displacement of copper from its solution. The reverse reaction is not spontaneous because copper, being a weaker reducing agent, cannot donate electrons to zinc ions.

7. During the electrolysis of dilute sulfuric acid using inert electrodes, what products are formed at the anode and cathode? Write the half-equations for the reactions.

Answer:

At the cathode (reduction): 2H⁺ + 2e⁻ → H₂

At the anode (oxidation): 2H₂O → O₂ + 4H⁺ + 4e⁻

Hydrogen gas is produced at the cathode, and oxygen gas is produced at the anode.

8. Calculate the mass of copper deposited at the cathode when a current of 2.0 A is passed through a copper(II) sulfate solution for 30 minutes. (Cu = 63.5 g/mol, F = 96,500 C/mol)

Answer:

Using Faraday's Law: m = (M × I × t) / (n × F)

Where:
m = mass of copper deposited (g)
M = molar mass of copper = 63.5 g/mol
I = current = 2.0 A
t = time = 30 × 60 = 1800 s
n = number of electrons required = 2 (for Cu²⁺ + 2e⁻ → Cu)
F = Faraday constant = 96,500 C/mol

m = (63.5 × 2.0 × 1800) / (2 × 96,500) = 1.18 g

Therefore, 1.18 g of copper will be deposited.

9. Why is a salt bridge necessary in a galvanic cell?

Answer: A salt bridge is necessary in a galvanic cell for three main reasons:
1. It completes the electrical circuit by allowing ions to flow between the two half-cells.
2. It maintains electrical neutrality in both half-cells by allowing anions to flow toward the anode and cations toward the cathode.
3. It prevents the direct mixing of the electrolyte solutions in the two half-cells, which would result in direct reaction rather than electron transfer through the external circuit.
Without a salt bridge, the cell would quickly stop functioning as charge would build up in the half-cells.

10. Explain how the Nernst equation shows the relationship between cell potential and concentration.

Answer: The Nernst equation, E = E° - (RT/nF)ln Q, relates the cell potential (E) to the standard cell potential (E°) and the reaction quotient (Q). At 25°C, it simplifies to E = E° - (0.0592/n)log Q. The equation shows that as the concentration of products increases (increasing Q), the cell potential decreases. Conversely, as the concentration of reactants increases (decreasing Q), the cell potential increases. This relationship reflects Le Chatelier's principle: changes in concentration affect the position of equilibrium, which in turn affects the cell potential.

Extended Response Questions

11. Describe the process of extracting aluminum by electrolysis, explaining why it requires so much energy and why recycling aluminum is economically advantageous.

Answer:

The extraction of aluminum from its ore involves several energy-intensive steps:

1. Bauxite mining and processing: Bauxite ore is mined and then processed using the Bayer process to produce pure aluminum oxide (alumina, Al₂O₃).

2. Electrolytic reduction (Hall-Héroult process): Alumina is dissolved in molten cryolite (Na₃AlF₆) at around 950-980°C. The mixture is electrolyzed using carbon anodes and a carbon-lined steel container as the cathode.

3. Electrode reactions:
At the cathode: Al³⁺ + 3e⁻ → Al(l)
At the anode: 2O²⁻ → O₂(g) + 4e⁻
The carbon anodes are also oxidized: C(s) + O²⁻ → CO₂(g) + 2e⁻

4. Collection of molten aluminum: The liquid aluminum collects at the bottom of the cell and is periodically tapped.

Energy requirements: The process requires enormous amounts of electricity (about 13-16 kWh per kg of aluminum) for several reasons:
- High temperatures (950-980°C) must be maintained to keep the cryolite molten
- Aluminum has a very negative electrode potential (-1.66V), requiring significant electrical energy to reduce Al³⁺ ions
- The breaking of the strong Al-O bonds in alumina requires substantial energy
- The process runs continuously, with industrial cells operating at currents of 100,000-200,000 amperes

Economic advantages of recycling:
- Recycling aluminum requires only about 5% of the energy needed for primary production
- This represents a 95% energy saving, significantly reducing costs
- No need for mining or processing bauxite ore
- Reduces waste and environmental impact
- Aluminum can be recycled indefinitely without loss of quality
- Recycling one aluminum can saves enough energy to run a TV for three hours

The high energy cost of primary aluminum production is what makes recycling so economically advantageous, as well as environmentally beneficial.

12. Compare and contrast galvanic and electrolytic cells, explaining their similarities, differences, and applications.

Answer:

Similarities:

Differences:

Galvanic (Voltaic) Cells Electrolytic Cells
Convert chemical energy to electrical energy Convert electrical energy to chemical energy
Involve spontaneous redox reactions (ΔG < 0) Involve non-spontaneous redox reactions (ΔG > 0)
Anode is negative, cathode is positive Anode is positive, cathode is negative
No external power source required External power source required
Electrons flow from anode to cathode in external circuit Electrons flow from power source to cathode
Half-cells are typically separated Can operate with electrodes in the same solution

Applications of Galvanic Cells:

Applications of Electrolytic Cells:

In essence, galvanic cells harness spontaneous chemical reactions to produce useful electrical energy, while electrolytic cells use electrical energy to drive desirable but non-spontaneous chemical reactions. Both types of cells are fundamental to modern technology and industry, with applications ranging from portable power sources to large-scale metal production.