CHM 51Chapter 18-Electrochemistry
18.5 - Cell Potentials and Free-Energy and the Equilibrium constant
ΔG = nFEor ΔGo= nFEo
Where F = Faraday constant = 96,500 C/mol e
Example
Calculate the cell potential at standard state (Eocell) for the following reaction. Then write the half reactions
I2(s) + 2 Br-(aq) 2I-(aq) + Br2(l) ΔGo = 1.1 x 105J
Example: Use tabulated electrode potentials to calculate G for the reaction. Is the reaction spontaneous?
2 Na(s) + 2 H2O(l) H2(g) + 2-OH(aq) + 2 Na+(aq)
Cell Potential When Ion Concentrations Are Not 1 M
•We know there is a relationship between the reaction quotient, Q; the equilibrium constant, K; and the free energy change, ΔGº.
•Changing the concentrations of the reactants and products so they are not 1 M will affect the standard free energy change, ΔGº.
•Because ΔGº determines the cell potential, Ecell, the voltage for the cell will be different when the ion concentrations are not 1 M.
The Nernst Equation - describe the relationship between Ecell and the concentration of species involved in the cell reaction
What is the cell potential at 25oC for the following short hand redox reaction?
Ni(s)/Ni2+ (1.0M)||Sn2+ (1.0 x 10-4M)/Sn(s)
Given
At equilibrium: Ecell = 0
Example
What is the value of Ecell for the voltaic cell below:
Pt(s)|Fe2+(0.1M),Fe3+(0.2M)||Ag+(0.1M)|Ag(s)
Examples
Calculate the equilibrium constant, Keq, for the reaction below
Zn2+(aq) + 2e- Zn(s) Eored = -0.76 V
Sn2+(aq) + 2e- Sn(s)Eored = -0.14 V
18.6 – Cell Potential and Concentration
•It is possible to get a spontaneous reaction when the oxidation and reduction reactions are the same, as long as the electrolyte concentrations are different.
•The difference in energy is due to the entropic difference in the solutions.
–The more concentrated solution has lower entropy than the less concentrated solution.
•Electrons will flow from the electrode in the less concentrated solution to the electrode in the more concentrated solution.
–Oxidation of the electrode in the less concentrated solution will increase the ion concentration in the solution; the less concentrated solution has the anode.
–Reduction of the solution ions at the electrode in the more concentrated solution reduces the ion concentration; the more concentrated solution has the cathode.
Example:Determine the cell potential for an electrochemical cell based on the following two half-reactions:
Oxidation: Cu(s) Cu2+(aq, 0.010M) + 2e-
Reduction: MnO4-(aq, 2.0M) + 4 H+(aq, 1.0M) + 3e- MnO2(s) + 2 H2O(l)
Example: Calculate the concentration of cadmium ion in the galvanic cell below
Cd(s)|Cd2+(aq)(?M)||Ni2+(aq)(0.100M)|Ni(s)
18.8 – Electrolysis: Driving Nonspontaneous Chemical reactions with Electricity
•In all electrochemical cells, oxidation occurs at the anode, reduction occurs at the cathode.
•In voltaic cells
–Anode is the source of electrons and has a (−) charge.
–Cathode draws electrons and has a (+) charge.
•In electrolytic cells
–Electrons are drawn off the anode, so it must have a place to release the electrons—the positive terminal of the battery.
–Electrons are forced toward the anode, so it must have a source of electrons—the negative terminal of the battery.
Electrolysis and Electrolytic Cells
•Electrolysis: The process of using an electric current to bring about chemical change.
• Process occurring in galvanic cell and electrolytic cells are the reverse of each other
• In an electrolytic cell, two inert electrodes are dipped into an aqueous solution
•The reaction that takes place is the opposite of the spontaneous process.
2 H2(g) + O2(g) 2 H2O(l) spontaneous
2 H2O(l) 2 H2(g) + O2(g) electrolysis
•Some applications are (1) metal extraction from minerals and purification, (2) production of H2 for fuel cells, and (3) metal plating.
•The electrical energy is supplied by a direct current power supply.
•AC alternates the flow of electrons so the reaction won’t be able to proceed.
•Some electrolysis reactions require more voltage than Ecell predicts. This is called the overvoltage.
•The source of energy is a battery or DC power supply.
•The positive terminal of the source is attached to the anode.
•The negative terminal of the source is attached to the cathode.
•Electrolyte can be either an aqueous salt solution or a molten ionic salt.
•Cations in the electrolyte are attracted to the cathode and anions are attracted to the anode.
•Cations pick up electrons from the cathode and are reduced; anions release electrons to the anode and are oxidized.
Electroplating:
18.7 Batteries: Using Chemistry to Generate Electricity
Acidic Dry Cell Battery:
Alkaline Dry Cell
Lead Storage Battery
Lithium ion battery
Fuel Cells
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