Electrochemistry - Introduction

  • Definition: Study of the relationships between electricity and chemical reactions
  • Importance:
    • Electrochemical cells in batteries and fuel cells
    • Electrolysis for metal extraction and purification
    • Corrosion prevention
  • Branches of electrochemistry:
    • Thermodynamics
    • Kinetics
    • Applications

Electrochemical Cells

  • Definition: Devices that convert chemical energy into electrical energy
  • Two types of electrochemical cells:
    1. Galvanic cells (also known as voltaic cells)
    2. Electrolytic cells
  • Components of a galvanic cell:
    • Anode (site of oxidation)
    • Cathode (site of reduction)
    • Salt bridge or porous barrier
    • Electrolyte solutions

Galvanic Cells

  • Spontaneous redox reaction
  • Electrons flow from anode (oxidation) to cathode (reduction)
  • Overall cell reaction: Anode (-) to Cathode (+)
  • Positive cell potential (E°) indicates feasibility of reaction
  • Example: Zn-Cu Galvanic Cell
    • Zn undergoes oxidation at the anode: Zn → Zn²⁺ + 2e⁻
    • Cu²⁺ undergoes reduction at the cathode: Cu²⁺ + 2e⁻ → Cu

Electrolytic Cells

  • Non-spontaneous redox reaction
  • External source of electrical energy is required
  • Electrons flow from cathode (reduction) to anode (oxidation)
  • Overall cell reaction: Cathode (-) to Anode (+)
  • Negative cell potential (E°) indicates non-feasibility of reaction
  • Example: Electrolysis of water
    • 2H₂O(l) → 2H₂(g) + O₂(g)

Standard Electrode Potential (E°)

  • Measure of relative tendency of a metal to lose or gain electrons
  • Standard hydrogen electrode (SHE) has E° = 0 V
  • E° is positive for metals that are good reducing agents (strongly reducing)
  • E° is negative for metals that are good oxidizing agents (strongly oxidizing)
  • Example:
    • Cu²⁺/Cu electrode has E° = +0.34 V (reducing agent)
    • Fe²⁺/Fe electrode has E° = -0.44 V (oxidizing agent)

Cell Potential (Ecell)

  • Measure of the driving force for an electrochemical cell
  • Calculated using the Nernst equation: Ecell = E°cell + (0.0592/n) log(Q)
  • n = number of electrons transferred in the reaction
  • Q = reaction quotient
  • Ecell is positive for spontaneous reactions (E°cell > 0) and negative otherwise
  • Example:
    • E°cell = +1.10 V, n = 2, Q = 0.2
    • Ecell = +1.10 V + (0.0592/2) log(0.2)

Faraday’s Laws of Electrolysis

  • First Law: The amount of substance produced at an electrode is directly proportional to the quantity of electricity passed through the cell.
  • Second Law: The amounts of different substances deposited or liberated by the passage of the same quantity of electricity are directly proportional to their chemical equivalent masses.
  • Example: Electrolysis of molten NaCl
    • 2NaCl(l) → 2Na(l) + Cl₂(g)
    • Mass of Na = (quantity of electricity passed x 2) / 96500

Application: Corrosion

  • Definition: Deterioration of metals due to chemical reactions with the environment
  • Factors affecting corrosion:
    • Presence of oxygen and moisture
    • Electrolyte or corrosive agents
    • Temperature
  • Methods to prevent corrosion:
    • Barrier protection (e.g., paint or coating)
    • Sacrificial protection (e.g., galvanization)
    • Cathodic protection (e.g., impressed current)

Application: Batteries

  • Definition: Portable electrochemical cells that provide electrical energy
  • Types of batteries:
    • Alkaline batteries (e.g., AA or AAA batteries)
    • Lead-acid batteries (e.g., car batteries)
    • Lithium-ion batteries (e.g., smartphone or laptop batteries)
  • Battery reactions involve redox reactions
  • Example: Lead-acid battery
    • Pb(s) + PbO₂(s) + 2H₂SO₄(aq) → 2PbSO₄(s) + 2H₂O(l)
  1. Electrochemical Equilibrium
  • Definition: State when the rate of the forward reaction equals the rate of the reverse reaction in an electrochemical cell
  • Equilibrium constant (K): Ratio of product concentrations to reactant concentrations at equilibrium
  • Nernst equation for equilibrium: E°cell = (0.0592/n) log(K)
  1. Electrochemical Series
  • Definition: Ranking of metals and non-metals based on their standard electrode potentials (E°)
  • Metals with more positive E° values are more likely to be reduced (strong reducing agents)
  • Metals with more negative E° values are more likely to be oxidized (strong oxidizing agents)
  • Non-metals (e.g., halogens) can also be ranked based on their reduction potentials
  1. Concentration Cells
  • Definition: Electrochemical cells with identical electrodes, but different ion concentrations
  • Difference in ion concentration drives the cell reaction
  • Example: Concentration cell using Zn²⁺ ions
    • Zn²⁺ (low concentration) → Zn²⁺ (high concentration) + 2e⁻
    • The electrons flow from the side of low concentration to the side of high concentration
  1. Fuel Cells
  • Definition: Electrochemical cells that convert chemical energy from fuels into electrical energy
  • Oxygen and fuel (e.g., hydrogen) react to produce water and release energy
  • Types of fuel cells:
    • Hydrogen fuel cells
    • Methanol fuel cells
    • Solid oxide fuel cells
  1. Corrosion Prevention - Barrier Protection
  • Definition: Creating a physical barrier between the metal and the environment to prevent corrosion
  • Examples:
    • Applying paint or coating on metal surfaces
    • Using plastic or rubber sleeves for metal wires
    • Galvanizing steel by coating it with a layer of zinc
  1. Corrosion Prevention - Sacrificial Protection
  • Definition: Using a more reactive metal to protect a less reactive metal
  • More reactive metal (sacrificial anode) undergoes oxidation instead of the less reactive metal
  • Examples:
    • Galvanizing iron or steel by coating it with a layer of zinc
    • Using magnesium or aluminum rods in water heaters to protect against corrosion
  1. Corrosion Prevention - Cathodic Protection
  • Definition: Applying an external electrical current to the metal to shift its potential to a more negative value and prevent oxidation
  • Examples:
    • Impressed current cathodic protection in pipelines or underground structures
    • Sacrificial anodes connected to the metal structure
  1. Batteries - Alkaline Batteries
  • Definition: Portable batteries that use an alkaline electrolyte (usually potassium hydroxide) and a zinc-manganese dioxide reaction
  • Example: AA or AAA batteries used in various electronic devices
  • Zn(s) + 2MnO₂(s) + 2NH₄Cl(aq) → ZnCl₂(aq) + 2MnO(OH)(s) + H₂O(l) + 2NH₃(g)
  1. Batteries - Lead-Acid Batteries
  • Definition: Rechargeable batteries commonly used in cars and other vehicles
  • Consist of lead and lead dioxide electrodes in a sulfuric acid electrolyte
  • Example: Car batteries
  • Pb(s) + PbO₂(s) + 2H₂SO₄(aq) → 2PbSO₄(s) + 2H₂O(l)
  1. Batteries - Lithium-Ion Batteries
  • Definition: Rechargeable batteries used in portable electronic devices
  • Consist of lithium cobalt oxide as the cathode, graphite as the anode, and a lithium salt electrolyte
  • Example: Smartphone or laptop batteries
  • C₆Li₈CoO₁₂(s) + 6LiC₆(s) → 8Li(s) + 6CO₂(g) + 12C(s)
  1. Electrochemical Equilibrium
  • Definition: State when the rate of the forward reaction equals the rate of the reverse reaction in an electrochemical cell
  • Equilibrium constant (K): Ratio of product concentrations to reactant concentrations at equilibrium
  • Nernst equation for equilibrium: E°cell = (0.0592/n) log(K)
  • Example:
    • 2Fe³⁺(aq) + 3S²⁻(aq) ⇌ Fe₂S₃(s)
    • E°cell = (0.0592/5) log(K)
  1. Electrochemical Series
  • Definition: Ranking of metals and non-metals based on their standard electrode potentials (E°)
  • Metals with more positive E° values are more likely to be reduced (strong reducing agents)
  • Metals with more negative E° values are more likely to be oxidized (strong oxidizing agents)
  • Non-metals (e.g., halogens) can also be ranked based on their reduction potentials
  • Example: Electrochemical series for metals
    • Li > K > Ba > Ca > Na > Mg > Al > Mn > Zn > Cr > Fe > Ni > Sn > Pb > H
  1. Concentration Cells
  • Definition: Electrochemical cells with identical electrodes, but different ion concentrations
  • Difference in ion concentration drives the cell reaction
  • Example: Concentration cell using Ag⁺ ions
    • Ag⁺ (low concentration) → Ag⁺ (high concentration) + e⁻
    • The electron flows from the side of low concentration to the side of high concentration
  • Applications of concentration cells: pH measurement, ion-selective electrodes
  1. Fuel Cells
  • Definition: Electrochemical cells that convert chemical energy from fuels into electrical energy
  • Oxygen and fuel (e.g., hydrogen) react to produce water and release energy
  • Types of fuel cells:
    • Hydrogen fuel cells (H₂ + ½O₂ → H₂O)
    • Methanol fuel cells (CH₃OH + 3/2O₂ → CO₂ + 2H₂O)
    • Solid oxide fuel cells (various fuel options)
  • Advantages of fuel cells: high efficiency, low pollution, versatility
  1. Corrosion Prevention - Barrier Protection
  • Definition: Creating a physical barrier between the metal and the environment to prevent corrosion
  • Examples:
    • Applying paint or coating on metal surfaces
    • Using plastic or rubber sleeves for metal wires
    • Galvanizing steel by coating it with a layer of zinc
  • Barrier protection methods reduce the contact of metal with corrosive agents (oxygen, moisture)
  1. Corrosion Prevention - Sacrificial Protection
  • Definition: Using a more reactive metal to protect a less reactive metal
  • More reactive metal (sacrificial anode) undergoes oxidation instead of the less reactive metal
  • Examples:
    • Galvanizing iron or steel by coating it with a layer of zinc
    • Using magnesium or aluminum rods in water heaters to protect against corrosion
  • Sacrificial protection involves sacrificing the more reactive metal to protect the less reactive metal
  1. Corrosion Prevention - Cathodic Protection
  • Definition: Applying an external electrical current to the metal to shift its potential to a more negative value and prevent oxidation
  • Examples:
    • Impressed current cathodic protection in pipelines or underground structures
    • Sacrificial anodes connected to the metal structure
  • Cathodic protection reduces the oxidation potential of the metal surface, preventing corrosion
  1. Batteries - Alkaline Batteries
  • Definition: Portable batteries that use an alkaline electrolyte (usually potassium hydroxide) and a zinc-manganese dioxide reaction
  • Example: AA or AAA batteries used in various electronic devices
  • Zn(s) + 2MnO₂(s) + 2NH₄Cl(aq) → ZnCl₂(aq) + 2MnO(OH)(s) + H₂O(l) + 2NH₃(g)
  • Alkaline batteries provide a higher energy density and longer shelf life compared to other battery types
  1. Batteries - Lead-Acid Batteries
  • Definition: Rechargeable batteries commonly used in cars and other vehicles
  • Consist of lead and lead dioxide electrodes in a sulfuric acid electrolyte
  • Example: Car batteries
  • Pb(s) + PbO₂(s) + 2H₂SO₄(aq) → 2PbSO₄(s) + 2H₂O(l)
  • Lead-acid batteries have high power output but are heavy and require regular maintenance
  1. Batteries - Lithium-Ion Batteries
  • Definition: Rechargeable batteries used in portable electronic devices
  • Consist of lithium cobalt oxide as the cathode, graphite as the anode, and a lithium salt electrolyte
  • Example: Smartphone or laptop batteries
  • C₆Li₈CoO₁₂(s) + 6LiC₆(s) → 8Li(s) + 6CO₂(g) + 12C(s)
  • Lithium-ion batteries have high energy density, long lifespan, and low self-discharge rates