Chapter 18 Electrochemistry Georgia Gwinnett College Chem 1212K Fall 2013 B Shepler Chapter 18 1 Reduction Oxidation Reaction Redox Redox an extensive and important class of reactions that is characterized by the transfer of electrons 2 Mg s O2 g 2 MgO s Magnesium is oxidized it gives up electrons as the charge on its atoms increases from zero to 2 Oxygen is reduced it gains electrons as the charge on its atoms decreases from zero to 2 i e becomes more negative Chapter 18 2 Rules for Determining Oxidation Numbers 1 The Oxidation numbers of the atoms in a neutral molecule must all add up to zero those in an ion must add up to the charge of the ion 2 Free Elements 0 Monatomic ions oxidation state charge 3 Cations a Alkali metals Group I b Alkaline Earth metals Group II c Al3 Zn2 Ag 4 Fluorine 1 5 Hydrogen 1 6 Oxygen 2 7 Other halogens 1 8 Sulfur group 2 9 Nitrogen group 3 10 Determine other elements based on above rules 3 Examples N2O4 Fe2 SO4 3 Ni NO3 2 Chapter 18 4 Redox Terminology 3MnO2 s 4Al s 3Mn s 2Al2O3 s Oxidation Number was Reduced Oxidation Number was Increased Gained Electrons Lost Electrons Mn was reduced Al was oxidized MnO2 was the oxidizing agent Al was the reducing agent Chapter 4 Atom Whole Compound 5 Balancing Redox Reactions in acidic solutions 1 Assign oxidation states to all atoms and identify the substances being oxidized and reduced 2 Separate the overall reaction into two half reactions one for oxidation and the other for reduction 3 Balance each half reaction with respect to mass i Balance elements that are NOT O and H ii Balance O by adding H2O 4 5 6 7 iii Balance H by adding H Balance the Charge of each half reaction by adding electrons Make the number of electrons equal by multiplying the half reactions by small integers Add the half reactions back together Cancel the electrons and any species as appropriate CHECK THAT THE REACTION IS BALANCED CHARGES AND ATOMS Chapter 18 6 Examples Balance the following redox reactions occurring in acidic aqueous solution K s Cr3 aq Cr s K aq 23 4 24 2 SO aq MnO aq SO aq Mn aq Chapter 18 7 Electrochemical Cells Galvanic Cells Voltaic What is the overall reaction Reduction half reaction Oxidation half reaction Define anode Define cathode Define salt bridge Show movement of e Show movement of ions in salt bridge 8 Electrochemical Cells Galvanic voltaic Cell Uses a spontaneous reaction to generate electrical energy G 0 System reaction does work on the surroundings Electrolytic Cell Use electrical energy to drive a nonspontaneous reactions G 0 Surroundings do work on the system reaction In any cell the electrodes Anode Cathode 9 Electrochemical Cells 10 Inert Electrodes What happens if a half reaction doesn t have a solid or metal that we can make an electrode with Use an inert electrode Pt or Graphite 11 Cell Potential Output of a Galvanic Cell The purpose of a galvanic cell is to convert the free energy change of a spontaneous chemical reaction into the kinetic energy of electrons moving through an external circuit electrical energy This electrical energy can do work proportional to the difference in electric potential of the two electrodes Stronger reducing agent at the anode more willing to give Higher cell up electrons potential if Stronger oxidizing agent at the cathode more willing to accept electrons SI unit of electrical potential Volt V SI unit of charge Coulomb C Two electrodes that differ by 1 volt do 1 joule of work for every coulomb passed between them 1 V 1 J C 12 Standard Cell Potential The measured potential of a galvanic cell depends on the concentration of the reactants products Standard Cell Potential The cell potential measured at a specified temperature 298 K if not stated otherwise and all components at their standard states E o cell Example 13 Standard Half Cell Electrode Potentials The reaction at each electrode makes up half of the overall reaction The potential of each half cell makes up half of the overall potential Standard Half Cell Potential The potential associated with a given half cell reaction with all components in their standard states Standard Half Cell Potentials are by convention written as reductions Zn2 2e Zn s E half cell 0 76 V Just as with G H S if we reverse the reaction the sign changes Zn Zn2 2e E half cell 0 76 V 14 Calculating Cell Potentials Half cells written as reductions Reverse the oxidation half cell Add half reactions to get overall cell reaction Add half cell potentials to get cell potential o o o o o E cell E copper E zinc E copper E zinc The standard cell potential is the difference between the half cell potential of the cathode and the half cell potential of the anode o o o E cell E cathode E anode 15 Determining E half cell The Standard Hydrogen Electrode Define the following half cell to have a potential of zero E reference 0 00 V 2H aq 1 M 2e H2 g 1 atm E reference 0 00 V H2 g 1 atm 2H aq 1 M 2e E reference 0 00 V Measure any other half cell relative to the Standard Hydrogen Electrode 16 Standard Half Cell Potentials 17 Standard Half Cell Potentials 18 Calculating Standard Cell Potentials 2Al s 3Sn2 aq 2Al 3 aq 3Sn s 19 Cell Potential is an Intensive Property An Extensive property depends on the amount of substance Examples are mass and energy An Intensive property does not depend on the amount of substance Examples are density and cell potential 20 Example Given the two half cells above design a galvanic cell and calculate the cell potential 21 Cell Potential and Free Energy A Galvanic cell has a positive potential A spontaneous reaction G 0 does work on the surroundings An electrolytic cell has a negative potential A nonspontaneous reaction G 0 is forced to proceed by work done on the system by the surroundings 22 Cell Potential Free Energy and Equilibrium 23 Example Use the table of standard cell potentials to calculate the standard Gibbs free energy and equilibrium constant for the following reaction at 298 K Pb s 2Ag aq Pb2 aq 2Ag s 24