Overview
This lecture explains the fundamentals of galvanic cells, focusing on their structure, operation, electron and ion flow, calculation of cell potential, the significance of standard conditions, and the use of cell notation. It also includes detailed worked examples and practical laboratory observations.
Key Principles of Galvanic Cells
- Galvanic cells convert chemical energy from redox (oxidation-reduction) reactions into electrical energy.
- Oxidation occurs at the anode, where a metal loses electrons to form positive ions.
- Reduction occurs at the cathode, where ions in solution gain electrons to form solid metal deposits.
- Electrons flow from the anode (which becomes negatively charged) to the cathode (which becomes positively charged) through an external wire.
- The separation of oxidation and reduction into different locations allows the chemical energy of redox reactions to be harnessed as electrical energy.
- The overall process involves the simultaneous occurrence of oxidation and reduction, known as a redox reaction.
Structure and Components
- Each half-cell consists of a metal electrode immersed in a solution containing its own ions (e.g., a copper electrode in copper(II) nitrate solution).
- Two half-cells are set up in separate beakers and connected by a wire and a voltmeter, which measures the total cell potential (voltage).
- A salt bridge, typically a glass tube filled with an ionic solution or filter paper soaked in a soluble salt like sodium or potassium nitrate, connects the two solutions. This allows ions to move between the half-cells and maintain electrical neutrality.
- The anode is the electrode where oxidation occurs; the cathode is where reduction occurs.
- The salt bridge is essential for maintaining the charge balance in both half-cells, allowing continuous electron flow and preventing the buildup of charge that would otherwise stop the reaction.
- In laboratory setups, the salt bridge may be a glass tube with pores or a piece of filter paper soaked in an ionic solution, such as potassium nitrate, bent to connect the two beakers.
Cell Potential and Standard Conditions
- The total cell potential (E°cell) is calculated by combining the standard reduction potentials of the two half-cells, using values from the standard reduction potential table:
- The reduction potential for the reduction half-cell (cathode) is used as listed.
- The reduction potential for the oxidation half-cell (anode) is reversed in sign to represent oxidation.
- The standard cell potential is measured under standard conditions:
- 1 M concentration of all aqueous species in both half-cells.
- Temperature of 298 K (25°C).
- Pressure of 1 atmosphere (100 kPa).
- If any of these conditions are not met, the measured cell potential will differ from the standard value provided in the reduction potential table.
- The voltmeter connected between the electrodes displays the total cell potential, which is the sum of the individual half-cell potentials.
- Standard reduction potentials are measured under these standard conditions, and deviations in concentration, temperature, or pressure will affect the observed cell potential.
Electron and Ion Flow
- Electrons move from the anode (site of oxidation) to the cathode (site of reduction) through the external circuit (wire and voltmeter).
- The anode becomes negatively charged as it loses electrons; the cathode becomes positively charged as it gains electrons.
- As the cell operates, electrons accumulate at the cathode and leave the anode, causing the charge difference between the electrodes to decrease. If uncorrected, this would eventually stop electron flow.
- The salt bridge allows:
- Anions (e.g., nitrate ions) to flow into the anode compartment to balance the increasing positive charge from metal ions formed during oxidation.
- Cations (e.g., sodium or potassium ions) to flow into the cathode compartment to balance the increasing negative charge as metal ions are reduced and removed from solution.
- This movement of ions through the salt bridge maintains electrical neutrality in both half-cells and sustains the flow of electrons.
- The direction of electron flow is always from the anode (oxidation half-cell) to the cathode (reduction half-cell).
- Over time, as more electrons flow, the anode becomes less negative and the cathode becomes less positive, reducing the potential difference. The salt bridge counteracts this by allowing ion migration to maintain the initial charge difference and keep the cell operating.
Observations in Galvanic Cells
- The mass of the anode decreases over time as metal atoms are oxidized to ions and enter the solution.
- The mass of the cathode increases as metal ions in solution are reduced and deposit as solid metal on the electrode.
- The color of the solution in the oxidation half-cell may intensify (e.g., copper(II) ions make the solution more blue) as more ions are produced during oxidation.
- The voltmeter shows a measurable potential difference as long as the cell operates and the salt bridge is intact.
- In practical setups, the salt bridge may be a glass tube with an ionic solution or a filter paper soaked in a soluble salt, such as potassium nitrate, connecting the two half-cells.
- The salt bridge is crucial for maintaining electrical neutrality and allowing the galvanic cell to function continuously.
Cell Notation (Shorthand)
- Galvanic cell notation provides a concise way to represent the setup without drawing a full diagram.
- The notation always lists the oxidation half-cell (anode) on the left and the reduction half-cell (cathode) on the right.
- The anode (solid metal) is written first, followed by a single vertical line separating it from its ion in solution.
- A double vertical line (||) represents the salt bridge.
- The cathode’s ion in solution is written next, followed by a single vertical line and then the cathode (solid metal).
- Single vertical lines separate different physical states (solid/aqueous); commas separate species of the same state.
- It is good practice to include the concentrations of aqueous species if known, as this affects cell potential.
- Example: For a zinc/nickel cell under standard conditions:
Zn(s) | Zn²⁺(aq, 1 M) || Ni²⁺(aq, 1 M) | Ni(s)
- If there are multiple species in the same state, they are separated by commas in the notation.
- The notation always starts with the oxidation half-cell on the left, followed by the reduction half-cell on the right, with the salt bridge in between.
Worked Example: Zn/Ni Galvanic Cell
- Oxidation (at anode):
Zn(s) → Zn²⁺(aq) + 2e⁻
- Reduction (at cathode):
Ni²⁺(aq) + 2e⁻ → Ni(s)
- Identifying electrodes:
- Anode: Zinc metal (Zn)
- Cathode: Nickel metal (Ni)
- Direction of electron flow:
From zinc (anode) to nickel (cathode) through the external wire.
- Calculating cell potential:
- Standard reduction potential for Zn²⁺/Zn: –0.76 V (reverse sign for oxidation: +0.76 V)
- Standard reduction potential for Ni²⁺/Ni: –0.24 V
- E°cell = 0.76 V – 0.24 V = 0.52 V
- Ion flow in salt bridge:
- Cations (e.g., Na⁺ or K⁺) flow into the cathode compartment to balance increasing negative charge.
- Anions (e.g., NO₃⁻) flow into the anode compartment to balance increasing positive charge.
- Cell notation:
Zn(s) | Zn²⁺(aq) || Ni²⁺(aq) | Ni(s)
- Summary of steps for analyzing a galvanic cell:
- Identify the oxidation and reduction reactions using the standard reduction potential table.
- Assign the anode and cathode based on which species is oxidized and which is reduced.
- Determine the direction of electron flow (from anode to cathode).
- Calculate the total standard cell potential by combining the half-cell potentials.
- Describe the direction of ion flow in the salt bridge to maintain electrical neutrality.
- Write the cell notation, listing the oxidation half-cell first, followed by the reduction half-cell.
Key Terms & Definitions
- Galvanic Cell: A device that converts chemical energy from redox reactions into electrical energy.
- Anode: The electrode where oxidation occurs (loses electrons); becomes negatively charged.
- Cathode: The electrode where reduction occurs (gains electrons); becomes positively charged.
- Salt Bridge: A medium (glass tube with ionic solution or soaked filter paper) that allows ion flow to maintain electrical neutrality between half-cells.
- Standard Cell Potential (E°cell): The voltage of a cell measured under standard conditions (1 M, 298 K, 1 atm).
- Cell Notation: Shorthand representation of the components and reactions in a galvanic cell, showing the order of species and the salt bridge.
- Standard Conditions: Conditions under which standard cell potentials are measured: 1 M concentration, 298 K temperature, and 1 atm pressure.
- Half-Cell: A compartment containing a metal electrode and a solution of its ions, where either oxidation or reduction occurs.
Action Items / Next Steps
- Review the standard reduction potential table to become familiar with common half-cell potentials and their values.
- Practice writing cell notations for various galvanic cell combinations, ensuring correct order and notation.
- Set up a simple galvanic cell in the laboratory, observe changes in electrode mass and solution color, and measure the cell voltage using a voltmeter.
- Consider how deviations from standard conditions (concentration, temperature, pressure) affect cell potential in practical experiments and be able to explain these effects.
- Reinforce understanding by working through additional examples and problems involving galvanic cells, redox reactions, and cell potential calculations.