Chapter 4: Electrochemistry

Overview

Electrochemistry is the study of the relationship between electrical energy and chemical reactions. This fundamental branch of chemistry explores how chemical reactions can produce electricity and how electricity can drive chemical reactions. From batteries and fuel cells to metal extraction and corrosion protection, electrochemical processes are essential to modern technology and industry. This chapter covers redox reactions, electrochemical cells, electrolysis, and practical applications that demonstrate the importance of electrochemistry in our daily lives.

Learning Objectives

After studying this chapter, you should be able to:

Define and identify oxidation and reduction processes
Understand the electrochemical series and its applications
Differentiate between voltaic and electrolytic cells
Analyze the operation of electrochemical cells like the Daniell cell
Explain electrolysis processes and their industrial applications
Understand metal extraction methods based on reactivity
Apply corrosion prevention techniques in various contexts
Solve problems involving electrochemical calculations

4.1 Oxidation and Reduction

What is Redox?

Redox reactions (reduction-oxidation) are chemical reactions where both oxidation and reduction processes occur simultaneously.

Definitions of Oxidation and Reduction

Process	Oxygen Definition	Hydrogen Definition	Electron Definition	Oxidation Number Definition
Oxidation	Gain of oxygen	Loss of hydrogen	Loss of electrons	Oxidation number increases
Reduction	Loss of oxygen	Gain of hydrogen	Gain of electrons	Oxidation number decreases

Key Concepts

Oxidation

Loss of electrons: Mg → M$g^2$⁺ + 2e⁻
Increase in oxidation number: Mg from 0 to +2
Occurs at the anode

Reduction

Gain of electrons: C$u^2$⁺ + 2e⁻ → Cu
Decrease in oxidation number: Cu from +2 to 0
Occurs at the cathode

Redox Agents

Oxidizing Agent

Causes oxidation in another substance
Gets reduced itself (gains electrons)
Example: C$u^2$⁺ oxidizes Mg to M $g^2$ ⁺ while being reduced to Cu

Reducing Agent

Causes reduction in another substance
Gets oxidized itself (loses electrons)
Example: Mg reduces C $u^2$ ⁺ to Cu while being oxidized to M $g^2$ ⁺

Example: Redox Reaction

Reaction: Mg(s) + CuS$O_4$(aq) → MgS$O_4$(aq) + Cu(s)

Half-reactions:

Oxidation: Mg → M$g^2$⁺ + 2e⁻ (Mg is oxidized, acts as reducing agent)
Reduction: C$u^2$⁺ + 2e⁻ → Cu (C $u^2$ ⁺ is reduced, acts as oxidizing agent)

Oxidation number changes:

Mg: 0 → +2 (increase = oxidation)
Cu: +2 → 0 (decrease = reduction)

Did You Know?

The human body relies heavily on redox reactions for energy production. Cellular respiration is essentially a redox process where glucose is oxidized and oxygen is reduced, producing ATP (energy currency of cells) as the primary product.

4.2 Electrochemical Cells

What is an Electrochemical Cell?

Electrochemical cell (also known as galvanic cell or voltaic cell) is a device that converts chemical energy into electrical energy through spontaneous redox reactions.

Daniell Cell Construction and Operation

Components

Electrodes: Two different metals, e.g., Zinc (Zn) and Copper (Cu)
Electrolytes: Solutions of metal ions for each electrode, e.g., ZnS $O_4$ and CuS $O_4$
Salt Bridge: U-shaped tube containing inert electrolyte (e.g., KCl or KN $O_3$ ) to connect solutions and allow ion flow

Operation Principles

Anode (Negative Terminal):

More electropositive metal (higher in electrochemical series)
Oxidation occurs: Zn(s) → Z$n^2$⁺(aq) + 2e⁻
Electrons flow through external circuit to cathode

Cathode (Positive Terminal):

Less electropositive metal (lower in electrochemical series)
Reduction occurs: C$u^2$⁺(aq) + 2e⁻ → Cu(s)
Electrons enter from external circuit

Cell Voltage: Potential difference between electrodes, measured using voltmeter Calculation: E°cell = E°cathode - E°anode

Electrochemical Series

The electrochemical series lists metals (and hydrogen) arranged by their standard electrode potentials (E°).

Key principles:

Higher position: More easily oxidized (stronger reducing agent)
Lower position: More easily reduced (stronger oxidizing agent)
Voltage calculation: E°cell = E°(higher position) - E°(lower position)

Common electrochemical series:

K+ | K (-2.92 V)
C$a^2$⁺ | Ca (-2.87 V)
Na+ | Na (-2.71 V)
M$g^2$⁺ | Mg (-2.37 V)
A$l^3$⁺ | Al (-1.66 V)
Z$n^2$⁺ | Zn (-0.76 V)
F$e^2$⁺ | Fe (-0.44 V)
H⁺ | H (0.00 V)
C$u^2$⁺ | Cu (+0.34 V)
Ag⁺ | Ag (+0.80 V)
A$u^3$⁺ | Au (+1.50 V)

Key Terms

Anode: Electrode where oxidation occurs
Cathode: Electrode where reduction occurs
Standard Electrode Potential (E°): Voltage produced when half-cell is connected to standard hydrogen electrode
Salt Bridge: Allows ion migration to complete electrical circuit
Electrolyte: Substance that conducts electricity in molten or aqueous state

SPM Exam Tips

Remember: Anode = Oxidation, Cathode = Reduction (A.O.C.R.)

In electrochemical series: Higher position = stronger reducing agent

Cell voltage = E°(cathode) - E°(anode) or E°(higher) - E°(lower)

At anode: metal → metal ion + electrons

At cathode: metal ion + electrons → metal

4.3 Electrolytic Cells

What is Electrolytic Cell?

Electrolytic cell is a device that uses electrical energy to produce non-spontaneous redox reactions. The process is called electrolysis.

Construction

DC power source
Two electrodes (usually inert like carbon or platinum)
Electrolyte (ionic compound in molten or aqueous state)

Operation Principles

Cathode (Negative Terminal):

Positive ions (cations) are attracted to cathode
Reduction occurs: cations gain electrons
Example: Na⁺ + e⁻ → Na (molten) or 2H⁺ + 2e⁻ → $H_2$ (aqueous)

Anode (Positive Terminal):

Negative ions (anions) are attracted to anode
Oxidation occurs: anions lose electrons
Example: 2Cl⁻ → C$l_2$ + 2e⁻ (molten) or 4OH⁻ → $O_2$ + 2$H_2$O + 4e⁻ (aqueous)

Factors Affecting Electrolysis Products (for Aqueous Solutions)

At Cathode

Rule: Cation with lower position in electrochemical series is preferentially discharged Example: Between Na⁺ and H⁺, H⁺ is discharged: 2H⁺ + 2e⁻ → $H_2$

At Anode

Rule: Anion with lower position in electrochemical series is preferentially discharged Example: Between S $O_4$ ²⁻ and OH⁻, OH⁻ is discharged: 4OH⁻ → $O_2$ + 2$H_2$O + 4e⁻

Special Cases

Concentrated halide ions: If Cl⁻, Br⁻, or I⁻ are concentrated, they are preferentially discharged over OH⁻
Active electrodes: If electrode is not inert (e.g., Cu), the electrode itself oxidizes: Cu → C$u^2$⁺ + 2e⁻

Key Terms

Electrolysis: Decomposition of chemical compounds by electric current
Electrolyte: Substance that conducts electricity in molten or aqueous state and undergoes chemical change
Inert electrode: Electrode that does not participate in reaction (e.g., carbon, platinum)

Safety Reminder

When working with electrochemical experiments:

Use proper eye protection and lab coats

Handle electrical equipment carefully

Be cautious with electrolytes (acids, bases, salts)

Ensure proper ventilation when gases are produced

Never use mains electricity - always use low-voltage DC power supplies

Follow proper waste disposal procedures for chemicals

4.4 Metal Extraction from Ores

What is Metal Extraction?

Metal extraction is the process of obtaining metals from their naturally occurring compounds (ores). The method used depends on the reactivity of the metal.

Extraction Methods Based on Reactivity

Very Reactive Metals (K, Na, Ca, Mg, Al)

Method: Electrolysis of molten compounds
Process: Requires large amount of energy
Example: Aluminum extraction from bauxite in Hall-Héroult process

2Al_2O_3(l) → 4Al(l) + 3O_2(g)

Moderately Reactive Metals (Zn, Fe, Sn, Pb)

Method: Reduction of metal oxides using carbon (coke) in blast furnace
Process: Carbon is cheaper and acts as reducing agent
Example: Iron extraction from iron ore

Fe_2O_3(s) + 3C(s) → 2Fe(l) + 3CO(g)

Less Reactive Metals (Cu, Hg, Ag, Au, Pt)

Method: Exist as elements or easily decomposed compounds by heating alone
Process: Simple thermal decomposition
Example: Mercury extraction from cinnabar

HgS(s) → Hg(l) + S(s)

Environmental Considerations

Problems:

Air pollution (acidic gases like S $O_2$ )
Water pollution
Habitat destruction

Solutions:

Pollution control devices
Recycling of metals
Sustainable mining practices

Key Terms

Ore: Naturally occurring rock containing metal minerals in economically extractable quantities
Blast Furnace: Furnace used for extracting iron from its ore
Reducing Agent: Substance that removes oxygen from another substance

4.5 Rusting (Corrosion)

What is Rusting?

Rusting is a redox reaction that occurs on iron when exposed to oxygen and water. It is a specific form of corrosion that damages iron structures.

Requirements for Rusting

Iron (Fe)
Oxygen ( $O_2$ )
Water ( $H_2$ O)

Electrochemical Mechanism

Anode (Center of water droplet)

Iron is oxidized: Fe → F$e^2$⁺ + 2e⁻

Cathode (Edge of water droplet)

Oxygen is reduced: $O_2$ + 2$H_2$O + 4e⁻ → 4OH⁻

Rust Formation

F $e^2$ ⁺ combines with OH⁻ to form Fe(OH)₂
Fe(OH)₂ is further oxidized by oxygen to form hydrated iron(III) oxide: $Fe_2O_3·xH_2O$ (rust)

Factors Accelerating Rusting

Presence of electrolytes: Salt or acidic pollutants dissolve in water to form better conducting solutions
Contact with less electropositive metals: When iron touches less reactive metal (e.g., copper), iron becomes anode and corrodes faster

Rust Prevention Methods

1. Surface Protection

Prevents iron from contact with water and oxygen
Examples:
- Painting
- Oiling/greasing
- Plastic coating
- Metal plating (e.g., tin on food cans)

2. Sacrificial Protection

Connect iron to more electropositive metal (e.g., zinc or magnesium)
More reactive metal corrodes first, protecting iron
Used in ship hulls and underground pipes
Example: Zn → Z$n^2$⁺ + 2e⁻ (sacrificial anode)

3. Alloying

Mix iron with other metals like chromium and nickel
Produces stainless steel, which is highly corrosion resistant

Key Terms

Rusting: Specific corrosion of iron
Corrosion: Gradual destruction of metals by chemical reaction with environment
Sacrificial Protection: Method of corrosion prevention where more reactive metal is sacrificed

Did You Know?

The Statue of Liberty is made of copper and has its characteristic green color due to patina formation - a protective layer of copper carbonate formed by oxidation over time. This layer actually protects the underlying metal from further corrosion.

4.6 Industrial Applications of Electrochemistry

1. Batteries

Primary batteries: Non-rechargeable (e.g., zinc-carbon, alkaline)
Secondary batteries: Rechargeable (e.g., lead-acid, lithium-ion)
Fuel cells: Convert chemical energy directly to electrical energy

2. Electroplating

Process of depositing one metal onto another
Applications:
- Decorative plating (gold, silver, chrome)
- Corrosion protection (zinc, nickel)
- Electrical conductivity (silver, copper)

3. Anodizing

Electrolytic process that thickens natural oxide layer on metals
Example: Aluminum anodizing for corrosion resistance and decorative finishes

4. Electrowinning

Process of extracting metals from solutions using electrolysis
Example: Copper extraction from copper sulfate solution

5. Electrorefining

Process of purifying metals using electrolysis
Example: Purification of copper

4.7 Electrochemical Calculations

Example 1: Calculating Cell Voltage

Problem: Calculate the cell voltage for a Zn-Cu cell using standard electrode potentials.

Given:

E°Z $n^2$ ⁺/Zn = -0.76 V
E°C $u^2$ ⁺/Cu = +0.34 V

Solution:

E°cell = E°cathode - E°anode
E°cell = E°C$u^2$⁺/Cu - E°Z$n^2$⁺/Zn
E°cell = +0.34 V - (-0.76 V)
E°cell = +0.34 V + 0.76 V
E°cell = +1.10 V

Example 2: Calculating Products of Electrolysis

Problem: Calculate the products when aqueous NaCl is electrolyzed.

Analysis:

Cathode: Na⁺ vs H⁺ → H⁺ discharged: 2H⁺ + 2e⁻ → $H_2$
Anode: Cl⁻ vs OH⁻ → Cl⁻ discharged if concentrated: 2Cl⁻ → C$l_2$ + 2e⁻

Overall reaction: 2NaCl(aq) + 2$H_2$O(l) → 2NaOH(aq) + $H_2$(g) + C$l_2$(g)

Summary

Key Concepts

Redox reactions involve simultaneous oxidation and reduction
Electrochemical cells convert chemical energy to electrical energy
Electrolytic cells use electrical energy to drive non-spontaneous reactions
Metal extraction methods depend on metal reactivity
Rusting is an electrochemical process requiring iron, oxygen, and water
Corrosion prevention includes surface protection, sacrificial protection, and alloying

Problem-Solving Strategy

Identify redox processes: Determine what is oxidized and reduced
Write half-reactions: Separate oxidation and reduction processes
Apply electrochemical series: Predict cell voltage and discharge products
Consider experimental conditions: Account for concentration, temperature, and electrode type
Calculate quantities: Use stoichiometry and electrochemical principles

Practical Applications

Batteries: Portable power sources
Electroplating: Metal coating and protection
Metal extraction: Industrial production of metals
Corrosion protection: Extending lifespan of metal structures
Water purification: Electrochemical treatment methods

Safety Considerations

Use appropriate protective equipment
Handle electrical equipment safely
Properly dispose of chemicals
Ensure adequate ventilation for gas production
Follow laboratory safety protocols

Practice Questions

Explain the difference between voltaic and electrolytic cells with examples.
Calculate the cell voltage for a Mg-Zn cell and predict the reactions at each electrode.
Describe how sacrificial protection works to prevent rusting of iron structures.
Explain why concentrated NaCl solution produces different electrolysis products than dilute NaCl solution.

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