Electrochemistry - Electrochemical Cell

  • Topic: Electrochemical Cell

  • Definition: A device that converts chemical energy into electrical energy

  • Consists of two half-cells: an oxidation half-cell and a reduction half-cell

  • Oxidation half-cell: where the oxidation reaction occurs

  • Reduction half-cell: where the reduction reaction occurs

  • Cell potential (Ecell): the measure of the voltage or electric potential difference between two electrode half-cells

  • Standard cell potential (E°cell): the cell potential measured under standard conditions

  • Standard conditions: 1 M concentration, 1 atm pressure, and 298 K temperature

  • Ecell can be positive (spontaneous reaction) or negative (non-spontaneous reaction)

  • Spontaneous reactions have positive standard cell potentials (E°cell > 0)

  • Half-cell potential (E°red): the potential difference between an electrode and its reduction reaction

  • Standard hydrogen electrode (SHE): the reference electrode used to measure E°red

  • Reduction potentials are measured relative to the SHE

  • Positive E°red values indicate a strong reducing agent

  • Negative E°red values indicate a strong oxidizing agent

  • Nernst equation: relates the cell potential to the concentration of reactants and products

  • Ecell = E°cell - (0.0592/n) * log(Q)

  • Ecell: cell potential

  • E°cell: standard cell potential

  • n: number of electrons transferred in the balanced equation

  • Q: reaction quotient (products/reactants)

  • Relationship between cell potential and Gibbs free energy change (∆G°):

  • ∆G° = -n * F * E°cell

  • ∆G°: Gibbs free energy change

  • n: number of electrons transferred in the balanced equation

  • F: Faraday’s constant (96,485 C/mol)

  • E°cell: standard cell potential

  • Electrolytic cell: a type of electrochemical cell used for non-spontaneous reactions

  • Requires an external power source to drive the reaction in the opposite direction

  • Anode: the electrode where oxidation occurs

  • Cathode: the electrode where reduction occurs

  • Positive terminal of the power source is connected to the anode

  • Electrolysis: the process of using an electric current to drive a non-spontaneous chemical reaction

  • The ions in the electrolyte migrate towards the respective electrodes

  • Cations move towards the cathode, while anions move towards the anode

  • At the electrodes, ions gain or lose electrons to form elements or compounds

  • Electroplating: an application of electrolysis in which a metal coating is deposited on an object

  • The object to be coated is used as the cathode

  • The metal to be plated is used as the anode

  • When the power source is applied, metal ions are reduced at the cathode, forming a metal coating

  • Corrosion: the gradual destruction of materials by chemical or electrochemical reactions with their environment

  • An example of spontaneous oxidation-reduction reactions

  • Factors influencing corrosion: moisture, dissolved oxygen, pH, temperature, and the presence of other ions

  • Methods to prevent corrosion: coatings, sacrificial anodes, and controlling the environment

  • Batteries: portable electrochemical cells used to power electronic devices

  • Primary batteries: non-rechargeable and designed for single use

  • Secondary batteries: rechargeable and can be repeatedly charged and discharged

  • Common types of batteries: alkaline, lithium-ion, lead-acid, and nickel-cadmium

  • Batteries play a crucial role in portable devices and electric vehicles

Electrochemistry - Electrochemical Cell

  • Electrochemical cell: device that converts chemical energy into electrical energy
  • Oxidation half-cell: where the oxidation reaction occurs
  • Reduction half-cell: where the reduction reaction occurs
  • Consists of two electrodes: anode (oxidation) and cathode (reduction)
  • Ions flow through an electrolyte to complete the circuit

Electrochemistry - Cell Potential

  • Cell potential (Ecell): measure of voltage between two electrode half-cells
  • Standard cell potential (E°cell): cell potential under standard conditions
  • Standard conditions: 1 M concentration, 1 atm pressure, 298 K temperature
  • Ecell can be positive (spontaneous) or negative (non-spontaneous)
  • Spontaneous reactions have positive standard cell potentials (E°cell > 0)

Electrochemistry - Half-Cell Potential

  • Half-cell potential (E°red): potential difference between an electrode and its reduction reaction
  • Standard hydrogen electrode (SHE): reference electrode for measuring E°red
  • Reduction potentials measured relative to SHE
  • Positive E°red values indicate strong reducing agents
  • Negative E°red values indicate strong oxidizing agents

Electrochemistry - Nernst Equation

  • Nernst equation: relates cell potential to concentration of reactants and products
  • Ecell = E°cell - (0.0592/n) * log(Q)
  • Ecell: cell potential
  • E°cell: standard cell potential
  • n: number of electrons transferred in balanced equation
  • Q: reaction quotient (products/reactants)

Electrochemistry - Gibbs Free Energy Change

  • Relationship between cell potential (E°cell) and Gibbs free energy change (ΔG°)
  • ΔG° = -n * F * E°cell
  • ΔG°: Gibbs free energy change
  • n: number of electrons transferred in balanced equation
  • F: Faraday’s constant (96485 C/mol)
  • E°cell: standard cell potential

Electrochemistry - Electrolytic Cell

  • Electrolytic cell: for non-spontaneous reactions
  • Requires external power source to drive reaction in opposite direction
  • Anode: electrode where oxidation occurs
  • Cathode: electrode where reduction occurs
  • Positive terminal of power source connected to anode

Electrochemistry - Electrolysis

  • Electrolysis: using electric current to drive non-spontaneous reaction
  • Ions in electrolyte migrate towards respective electrodes
  • Cations move towards cathode, anions move towards anode
  • At electrodes, ions gain or lose electrons to form elements or compounds

Electrochemistry - Electroplating

  • Electroplating: deposition of metal coating on object using electrolysis
  • Object to be coated used as cathode
  • Metal to be plated used as anode
  • Metal ions reduced at cathode, forming metal coating

Electrochemistry - Corrosion

  • Corrosion: gradual destruction of materials by chemical or electrochemical reactions with environment
  • Example of spontaneous oxidation-reduction reactions
  • Factors influencing corrosion: moisture, dissolved oxygen, pH, temperature, presence of other ions
  • Methods to prevent corrosion: coatings, sacrificial anodes, controlling environment

Electrochemistry - Batteries

  • Batteries: portable electrochemical cells used to power devices
  • Primary batteries: non-rechargeable, single use
  • Secondary batteries: rechargeable, can be repeatedly charged and discharged
  • Types of batteries: alkaline, lithium-ion, lead-acid, nickel-cadmium
  • Important for portable devices, electric vehicles

Electrochemistry - Example of Spontaneous Reaction

  • Consider the reaction: Fe(s) + Cu²⁺(aq) -> Fe²⁺(aq) + Cu(s)
  • Standard cell potential (E°cell) = +0.338 V (positive value)
  • This is a spontaneous reaction because E°cell is positive
  • Flow of electrons occurs from Fe electrode (anode) to Cu electrode (cathode)

Electrochemistry - Example of Non-Spontaneous Reaction

  • Consider the reaction: Ag(s) + Fe²⁺(aq) -> Ag⁺(aq) + Fe(s)
  • Standard cell potential (E°cell) = -0.337 V (negative value)
  • This is a non-spontaneous reaction because E°cell is negative
  • External power source is required to drive the reaction in the opposite direction

Electrochemistry - Electrolyte in Electrolytic Cell

  • Electrolyte: a solution or molten salt that conducts electricity
  • Contains ions that can move towards the respective electrodes
  • Must be present in an electrolytic cell for the completion of the circuit
  • Examples of electrolytes include aqueous solutions of acids, bases, and salts

Electrochemistry - Faraday’s Constant

  • Faraday’s constant (F) is a conversion factor between electrical charge and the amount of substance produced or consumed in a reaction
  • F = 96,485 C/mol
  • One Faraday of charge (1 F) is equivalent to the charge of one mole of electrons (6.022 × 10^23 electrons)

Electrochemistry - Charging and Discharging of Batteries

  • Charging: process of converting electrical energy into chemical energy in a battery
    • Current flows from an external power source, forcing electrons into the battery
    • Reactants are converted into products through reduction and oxidation reactions
  • Discharging: process of converting chemical energy into electrical energy in a battery
    • Reactants in the battery are consumed, converting them into products
    • Electrons flow from the battery to an external device, generating an electrical current

Electrochemistry - Redox Reaction in Batteries

  • Batteries rely on redox reactions to generate electricity
  • In a lithium-ion battery:
    • Anode: Li⁺ ions oxidize to release electrons
    • Cathode: Li⁺ ions reduce to accept electrons
  • Redox reactions occur at the electrodes, allowing for the flow of electrons through the external circuit

Electrochemistry - Role of Electrolytes in Batteries

  • Electrolytes in batteries allow for the movement of ions within the cell
  • Facilitate the flow of ions between the electrodes
  • Enhance the efficiency of redox reactions at the electrodes
  • Common electrolytes used in batteries include lithium salts, sulfuric acid, and alkaline solutions

Electrochemistry - Factors Affecting Battery Performance

  • Several factors influence the performance of batteries:
    1. Temperature: High temperatures can increase reaction rates, but excessive heat can damage the battery.
    2. State of Charge: The amount of charge remaining in the battery affects its voltage and capacity.
    3. Current Drain: Higher currents drain the battery more quickly, reducing its overall capacity.
    4. Internal Resistance: Higher internal resistance can lead to voltage drop and reduced performance.
    5. Aging: Over time, battery capacity decreases due to internal chemical changes.

Electrochemistry - Applications of Electrochemical Cells

  • Electrochemical cells have various applications:
    • Batteries: portable power sources for electronic devices, vehicles, and backup power
    • Fuel cells: convert chemical energy into electrical energy using an external fuel supply
    • Sensors: used in analytical chemistry for detecting and measuring substances
    • Electroplating: metal coating applications in industries, jewelry, and decorative items

Electrochemistry - Summary

  • Electrochemical cells convert chemical energy into electrical energy.
  • Cell potential is the measure of voltage between electrode half-cells.
  • Nernst equation relates cell potential to the concentration of reactants and products.
  • Gibbs free energy change is related to the cell potential.
  • Electrolytic cells require an external power source for non-spontaneous reactions.
  • Batteries are portable electrochemical cells used to power devices.
  • Corrosion is the gradual destruction of materials by chemical or electrochemical reactions.
  • Electroplating and electrolysis are applications of electrolysis.
  • Understanding electrochemistry is essential for various industries and energy storage applications.