Chapter 18: ElectrochemistryReduction-Oxidation Reaction: RedoxRules for Determining Oxidation NumbersExamplesRedox TerminologyBalancing Redox Reactions (in acidic solutions)Slide 7Electrochemical Cells: Galvanic Cells (Voltaic)Electrochemical CellsSlide 10Inert ElectrodesCell Potential: Output of a Galvanic CellStandard Cell PotentialStandard Half-Cell (Electrode) PotentialsCalculating Cell PotentialsDetermining E°half-cell: The Standard Hydrogen ElectrodeStandard Half-Cell PotentialsSlide 18Calculating Standard Cell PotentialsCell Potential is an Intensive PropertyExampleCell Potential and Free EnergyCell Potential, Free Energy, and EquilibriumSlide 24Chapter 18 1Chapter 18: ElectrochemistryGeorgia Gwinnett CollegeChem 1212KFall 2013(B. Shepler)Chapter 18 2Reduction-Oxidation Reaction: RedoxRedox: an extensive and important class of reactions that is characterized by the transfer of electrons.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).2 Mg(s) + O2(g) → 2 MgO(s)Rules for Determining Oxidation Numbers1. 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 = charge3. Cationsa) Alkali metals (Group I):b) Alkaline Earth metals (Group II):c) Al3+, Zn2+, Ag+4. Fluorine: -15. Hydrogen: +16. Oxygen: -27. Other halogens: -18. Sulfur group: -29. Nitrogen group: -310. Determine other elements based on above rules3Examples•N2O4•Fe2(SO4)3•Ni(NO3)2Chapter 18 4Chapter 4 5Redox Terminology3MnO2(s) 4Al(s)+ 3Mn(s) + 2Al2O3(s)Oxidation Number was ReducedOxidation Number was IncreasedGained Electrons Lost ElectronsMn was reduced Al was oxidizedMnO2 was the oxidizing agentAl was the reducing agentAtomWhole CompoundBalancing 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 H2Oiii. Balance H by adding H+4. Balance the Charge of each half-reaction by adding electrons.5. Make the number of electrons equal by multiplying the half reactions by small integers.6. Add the half-reactions back together. Cancel the electrons and any species as appropriate.7. CHECK THAT THE REACTION IS BALANCED!–CHARGES AND ATOMS!Chapter 18 6ExamplesBalance the following redox reactions occurring in acidic aqueous solution.Chapter 18 7K(s)+Cr3+(aq) ® Cr(s) +K+(aq)SO32-(aq) +MnO4-(aq) ® SO42-(aq) +Mn2+(aq)Electrochemical Cells: Galvanic Cells (Voltaic)8What 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 bridgeElectrochemical Cells•Galvanic (voltaic) Cell: Uses a spontaneous reaction to generate electrical energy (ΔG < 0). System (reaction) does work on the surroundings.•Electrolytic Cell: Use electrical energy to drive a nonspontaneous reactions (ΔG > 0). Surroundings do work on the system (reaction).•In any cell the electrodes–Anode:–Cathode:9Electrochemical Cells10Inert Electrodes11What 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 GraphiteWhat 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 GraphiteCell 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 up electrons)•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.12Higher cellpotential if:1 V = 1 J/CStandard 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. •Example:13 EcelloStandard 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 ΔG, ΔH, ΔS if we reverse the reaction the sign changes.Zn  Zn2+ + 2e− E°half-cell=+0.76 V14Calculating Cell Potentials15Add half-reactions to get overall cell reaction. Add half-cell potentials to get cell potential.Half-cells written as reductionsReverse the oxidation half-cell Ecello = Ecoppero + (- Ezinco) = Ecoppero - Ezi ncoThe standard cell potential is the difference between the half-cell potential of the cathode and the half-cell potential of the anode. Ecello = Ecathodeo - EanodeoDetermining E°half-cell: The Standard Hydrogen ElectrodeDefine the following half-cell to have a potential of zero (E°referenceď= 0.00 V)2H+(aq, 1 M) + 2e−  H2(g, 1 atm) E°reference = 0.00 VH2(g, 1 atm)  2H+(aq, 1 M) + 2e− −E°reference = 0.00 VMeasure any other half-cell relative to the Standard Hydrogen Electrode16Standard Half-Cell Potentials17Standard Half-Cell Potentials18Calculating Standard Cell Potentials192Al(s) + 3Sn2+(aq) ® 2Al3+(aq) + 3Sn(s)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