# Enduring Understanding 3.C.2: Electrochemistry

• Electrochemistry is the study of redox reactions that occur in electrochemical cells.
• Galvanic cells are where a redox reaction is used to generate an electric current
• Electrolytic cells are where an electric current is used to drive a redox reaction.
• A half-cell is a reactant placed in an electrolyte solution. When two half-cells are connected by a conducting wire and a salt bridge, an electrochemical cell is created.

• In electrochemical cells,
• Oxidation occurs at the anode
• Reduction occurs at the cathode.
• The electrical potential can be calculated by identifying the oxidation and reduction half reactions, and using a table of standard reduction potentials.

• Example: A Zn-Cu battery:
• The half reactions are:
• Cu2+(aq) + 2e- → Cu(s)
• Zn(s) → Zn2+(aq) + 2e-
• At the anode, the Zn is oxidized, and at the cathode, the Cu is reduced.
• The standard reduction potentials are:
• Cu2+(aq) + 2e- → Cu(s) ; Eo = 0.339 V
• Zn2+(aq) + 2e- → Zn(s) ; Eo = - 0.76 V
• Since the Zn is being oxidized, the reduction potential for Zn must be reversed, so +0.76 V
• The standard potential of the cell is 0.339 + 0.76 = 1.10 V

• The charge on an electron is 1.602x10-19 C (coulombs).
• The charge on one mole electrons is called a Faraday (F), which is: ?
• F = (6.022045x1023 / mol) x (1.602x10-19 C) = 96485 C/mol
• One ampere is a measure of an amount of electrical charge going through an electric circuit in one second: 1 C/sec
• The free energy of an electrochemical cell is given by: ΔG = nFEo

• Sample Question: In an electroplating experiment, 10 amps are applied to a Cu2+ solution for 96.4 seconds. How many moles of Cu(s) would be deposited?
• 10 amps = 10 C/s
• 10 C/s x 96.5 s = 965 C
• 964 C / 96385 C/mol = 0.01 mol electrons
• There are 2 electrons required to reduce Cu2+ to Cu(s)
• Therefore, 0.01 mol/2 or 0.005 mol, or 5 mmol of Cu would be deposited.

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