Galvanic Cell Diagram
Electron Flow & Current Direction
Ecell vs Reaction Quotient Q
Concentration Effect on EMF
Cell Parameters
Standard Electrode Potentials
Ion Concentrations (M)
Temperature
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Common Galvanic Cells
Galvanic Cell Equations
What is a Galvanic Cell?
A galvanic cell (or voltaic cell) is an electrochemical cell that converts chemical energy into electrical energy through spontaneous redox reactions. It consists of two half-cells connected by a salt bridge and an external circuit. Electrons flow from the anode (oxidation) to the cathode (reduction), generating an electric current that can do work.
Nernst Equation
Nernst Equation: E = E° - (RT/nF)ln(Q), where E is the cell potential under non-standard conditions, E° is the standard cell potential, R is the gas constant (8.314 J/mol·K), T is temperature in Kelvin, n is the number of electrons transferred, F is Faraday's constant (96485 C/mol), and Q is the reaction quotient.
At 298 K: E = E° - (0.0592/n)log(Q). This simplified form is commonly used at room temperature.
Applications: Predicting cell voltage under various conditions, determining equilibrium constants, and calculating concentration effects.
Electrode Processes
Anode (Oxidation): The electrode where oxidation occurs. Metal atoms lose electrons and enter the solution as ions. Electrons are released into the external circuit. The anode is negatively charged in a galvanic cell.
Cathode (Reduction): The electrode where reduction occurs. Ions from the solution gain electrons and deposit as metal atoms. Electrons are consumed from the external circuit. The cathode is positively charged in a galvanic cell.
Salt Bridge: Maintains electrical neutrality by allowing ion flow between half-cells, completing the circuit.
Concentration Effects
Le Chatelier's Principle: Increasing reactant concentration (or decreasing product concentration) shifts the equilibrium toward products, increasing cell potential. Decreasing reactant concentration has the opposite effect.
Q < K: When Q < K, ΔG < 0 and E > 0, the reaction proceeds spontaneously in the forward direction.
Q = K: At equilibrium, ΔG = 0 and E = 0, no net reaction occurs.
Q > K: When Q > K, ΔG > 0 and E < 0, the reaction would proceed spontaneously in reverse.
Temperature Effects
Temperature affects the cell potential through the (RT/nF) term in the Nernst equation. Higher temperatures increase the magnitude of the correction term, making the cell potential more sensitive to concentration changes. The temperature dependence varies with the entropy change of the reaction: ΔG = ΔH - TΔS. Exothermic reactions (ΔH < 0) typically have decreased spontaneity at higher temperatures.
Real-World Applications
Batteries: All commercial batteries are based on galvanic cell principles. From alkaline batteries to lithium-ion batteries in phones and electric vehicles.
Corrosion Prevention: Understanding galvanic cells helps prevent unwanted corrosion in structures like ships, pipelines, and bridges.
Biological Systems: Nerve impulses and muscle contractions involve electrochemical gradients similar to galvanic cells.
Industrial Electroplating: Uses external voltage to drive non-spontaneous reactions opposite to galvanic cell processes.
pH Measurements: Glass electrodes work on potentiometric principles related to cell potentials.
Common Galvanic Cells
Daniell Cell: Zn|Zn²⁺||Cu²⁺|Cu, E° = 1.10 V. One of the earliest practical batteries.
Voltaic Pile: Zn|Ag, historical first battery invented by Volta.
Lemon Battery: Zn|Cu using citric acid as electrolyte, educational demonstration.
Lead-Acid Battery: Pb|PbSO₄||PbSO₄|PbO₂, used in automobiles, E° ≈ 2.0 V.
Fuel Cells: Continuous galvanic cells using external fuel supply, like hydrogen fuel cells.