Electrochemistry - Electrochemical Cell
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Topic: Electrochemical Cell
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Definition: A device that converts chemical energy into electrical energy
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Consists of two half-cells: an oxidation half-cell and a reduction half-cell
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Oxidation half-cell: where the oxidation reaction occurs
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Reduction half-cell: where the reduction reaction occurs
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Cell potential (Ecell): the measure of the voltage or electric potential difference between two electrode half-cells
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Standard cell potential (E°cell): the cell potential measured under standard conditions
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Standard conditions: 1 M concentration, 1 atm pressure, and 298 K temperature
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Ecell can be positive (spontaneous reaction) or negative (non-spontaneous reaction)
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Spontaneous reactions have positive standard cell potentials (E°cell > 0)
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Half-cell potential (E°red): the potential difference between an electrode and its reduction reaction
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Standard hydrogen electrode (SHE): the reference electrode used to measure E°red
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Reduction potentials are measured relative to the SHE
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Positive E°red values indicate a strong reducing agent
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Negative E°red values indicate a strong oxidizing agent
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Nernst equation: relates the cell potential to the concentration of reactants and products
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Ecell = E°cell - (0.0592/n) * log(Q)
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Ecell: cell potential
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E°cell: standard cell potential
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n: number of electrons transferred in the balanced equation
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Q: reaction quotient (products/reactants)
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Relationship between cell potential and Gibbs free energy change (∆G°):
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∆G° = -n * F * E°cell
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∆G°: Gibbs free energy change
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n: number of electrons transferred in the balanced equation
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F: Faraday’s constant (96,485 C/mol)
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E°cell: standard cell potential
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Electrolytic cell: a type of electrochemical cell used for non-spontaneous reactions
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Requires an external power source to drive the reaction in the opposite direction
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Anode: the electrode where oxidation occurs
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Cathode: the electrode where reduction occurs
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Positive terminal of the power source is connected to the anode
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Electrolysis: the process of using an electric current to drive a non-spontaneous chemical reaction
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The ions in the electrolyte migrate towards the respective electrodes
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Cations move towards the cathode, while anions move towards the anode
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At the electrodes, ions gain or lose electrons to form elements or compounds
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Electroplating: an application of electrolysis in which a metal coating is deposited on an object
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The object to be coated is used as the cathode
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The metal to be plated is used as the anode
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When the power source is applied, metal ions are reduced at the cathode, forming a metal coating
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Corrosion: the gradual destruction of materials by chemical or electrochemical reactions with their environment
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An example of spontaneous oxidation-reduction reactions
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Factors influencing corrosion: moisture, dissolved oxygen, pH, temperature, and the presence of other ions
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Methods to prevent corrosion: coatings, sacrificial anodes, and controlling the environment
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Batteries: portable electrochemical cells used to power electronic devices
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Primary batteries: non-rechargeable and designed for single use
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Secondary batteries: rechargeable and can be repeatedly charged and discharged
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Common types of batteries: alkaline, lithium-ion, lead-acid, and nickel-cadmium
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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:
- Temperature: High temperatures can increase reaction rates, but excessive heat can damage the battery.
- State of Charge: The amount of charge remaining in the battery affects its voltage and capacity.
- Current Drain: Higher currents drain the battery more quickly, reducing its overall capacity.
- Internal Resistance: Higher internal resistance can lead to voltage drop and reduced performance.
- 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.
