1. Electrochemistry Experiment 12

2. Oxidation – Reduction Reactions Consider the reaction of Copper wire and AgNO 3 (aq) AgNO 3 (aq) Ag(s) Cu(s)

3. Oxidation – Reduction Reactions If you leave the reaction a long time the solution goes blue! The blue is due to Cu 2+ (aq)

4. Oxidation-Reduction Reactions So when we mix Ag + (aq) with Cu (s) we get Ag(s) and Cu 2+ (aq) Ag + (aq) + 1e -  Ag (s) Cu (s)  Cu 2+ (aq) + 2e - The electrons gained by Ag + must come from the Cu 2+ Can’t have reduction without oxidation (redox) Each Cu can reduce 2 Ag + 2Ag + (aq) + 2e -  2Ag (s) Cu (s)  Cu 2+ (aq) + 2e - 2Ag + (aq) + 2e - + Cu (s)  2Ag (s) + Cu 2+ (aq) + 2e - lose electrons = oxidation gain electrons = reduction

5. Redox Cu/Ag Cu Ag + E electron flow

6. Redox Cu/Ag Cu 2 + Ag E ΔE = e.V e = charge on an electron V = Voltage in a electrochemical cell

7. 7 Redox Reactions & Current redox reactions involve the transfer of electrons from one substance to another therefore, redox reactions have the potential to generate an electric current in order to use that current, we need to separate the place where oxidation is occurring from the place that reduction is occurring

8. 8 Electric Current Flowing Directly Between Atoms

9. Tro, Chemistry: A Molecular Approach9 Electric Current Flowing Indirectly Between Atoms

10. Tro, Chemistry: A Molecular Approach10 Electrochemical Cells electrochemistry is the study of redox reactions that produce or require an electric current the conversion between chemical energy and electrical energy is carried out in an electrochemical cell spontaneous redox reactions take place in a voltaic cell – aka galvanic cells nonspontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy

11. Tro, Chemistry: A Molecular Approach11 Electrochemical Cells oxidation and reduction reactions kept separate – half-cells electron flow through a wire along with ion flow through a solution constitutes an electric circuit requires a conductive solid (metal or graphite) electrode to allow the transfer of electrons – through external circuit ion exchange between the two halves of the system – electrolyte

12. Tro, Chemistry: A Molecular Approach12 Electrodes Anode (donates electrons to the cathode) – electrode where oxidation occurs – anions attracted to it – connected to positive end of battery in electrolytic cell – loses weight in electrolytic cell Cathode (attracts electrons from the anode) – electrode where reduction occurs – cations attracted to it – connected to negative end of battery in electrolytic cell – gains weight in electrolytic cell electrode where plating takes place in electroplating

13. Tro, Chemistry: A Molecular Approach13 Voltaic Cell the salt bridge is required to complete the circuit and maintain charge balance

14. Tro, Chemistry: A Molecular Approach14 Current and Voltage the number of electrons that flow through the system per second is the current – unit = Ampere – 1 A of current = 1 Coulomb of charge flowing by each second – 1 A = 6.242 x 10 18 electrons/second – Electrode surface area dictates the number of electrons that can flow the difference in potential energy between the reactants and products is the potential difference (the potential for an electric field to cause an electrical current) – unit = Volt – 1 V of force = 1 J of energy/Coulomb of charge – the voltage needed to drive electrons through the external circuit – amount of force pushing the electrons through the wire is called the electromotive force, emf

15. Tro, Chemistry: A Molecular Approach15 Cell Potential the difference in potential energy between the anode the cathode in a voltaic cell is called the cell potential the cell potential depends on the relative ease with which the oxidizing agent is reduced at the cathode and the reducing agent is oxidized at the anode the cell potential under standard conditions is called the standard emf, E° cell – 25°C, 1 atm for gases, 1 M concentration of solution – sum of the cell potentials for the half-reactions

16. 16 Standard Reduction Potential a half-reaction with a strong tendency to occur has a large + half-cell potential when two half-cells are connected, the electrons will flow so that the half-reaction with the stronger tendency will occur we cannot measure the absolute tendency of a half- reaction, we can only measure it relative to another half- reaction we select as a standard half-reaction the reduction of H + to H 2 under standard conditions, which we assign a potential difference = 0 V – standard hydrogen electrode, SHE

17. Tro, Chemistry: A Molecular Approach17

18. Tro, Chemistry: A Molecular Approach18 Half-Cell Potentials SHE reduction potential is defined to be exactly 0 V half-reactions with a stronger tendency toward reduction than the SHE have a + value for E° red half-reactions with a stronger tendency toward oxidation than the SHE have a + value for E° red ΔE° cell = E° oxidation + E° reduction – E° oxidation = -E° reduction – when adding E° values for the half-cells, do not multiply the half-cell E° values, even if you need to multiply the half-reactions to balance the equation ΔG o cell =-nFΔE° cell

20. Tro, Chemistry: A Molecular Approach20 Electrochemical Cell Summary salt bridge e-e- anode Zn (s)--> Zn 2+ (aq)+ 2e - cathode Cu 2+ (aq)+ 2e - --> Cu(s ) The differing stability of reactants, (Zn(s), Cu 2+ (aq)), and products (Zn 2+, and Cu(s)), creates a potential energy gradient through which the charges migrate (from high energy to low). This manifests as a potential difference E cell, across the electrodes. Where -qE cell is the change in potential energy when an amount of negative charge (-q) passes from the anode to the cathode The cell potential is related to the free energy of the reaction according to the relation  G cell = -nFE cell The cell potential can be calculated knowing the standard reduction potentials. These can be used to find E o red for the reaction at the cathode, and E o ox (= - E o red ). Then E o cell = E o ox + E o red Zn 2+ (aq) + 2e - --> Zn(s)-0.76V Cu 2+ (aq) + 2e - --> Cu(s) E red =0.34V Zn(s) --> Zn 2+ (aq) + 2e - E ox = 0.76V E cell =1.1 V E cell = 0.76V+0.34V = 1.1V

21. Tonight Construction of Voltaic Cells and Measurement of Cell Potentials Use the corresponding 0.1 M metal sulfate of the same metal as the electrode in the half cell Construct a salt bridge Measure the voltage