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Chapter 20 Electron Transfer Reactions Objectives: 1. Carry out balancing of redox reactions in acidic or basic solutions; 2. Recall the parts of a basic and commercial voltaic cells; 3. Perform cell potential calculations from standard reduction potentials; 4. Classify oxidizing and reducing agents; 5. Apply the Nerst equation to redox problems; 6. Determine K from Ecell; 7. Perform electrolysis calculations. Introduction • NADH + (1/2)O2 + H+ -----> NAD+ + H2O • Medicinal Biochemistry: • http://web.indstate.edu/thcme/mwking/home.html Introduction • Pyruvate + CoA + NAD+ ------> CO2 + acetyl-CoA + NADH + H+ Redox Reactions Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s) Electron transfer reactions are ________________ or redox reactions. Redox reactions can result in the generation of an electric current or be caused by imposing an electric current. Therefore, this field of chemistry is often called _____________________. Review • OXIDATION – _______________________________ • REDUCTION – _______________________________ • OXIDIZING AGENT – _______________________________ • REDUCING AGENT – _______________________________ Redox Reactions Direct Redox Reaction Oxidizing and reducing agents in direct contact. Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s) Redox Reactions Indirect Redox Reaction A battery functions by transferring electrons through an external wire from the reducing agent to the oxidizing agent. Electrochemical Cells • An apparatus that allows a redox reaction to occur by transferring electrons through an external connector. • Product favored reaction ---> ________________cell ----> electric current. • Reactant favored reaction ---> ________________ cell ---> electric current used to cause chemical change. Batteries are voltaic cells Electrochemistry Alessandro Volta, 17451827, Italian scientist and inventor. Luigi Galvani, 1737-1798, Italian scientist and inventor. Why study electrochemistry? • Batteries • Corrosion • Industrial production of chemicals such as Cl2, NaOH, F2 and Al • Biological redox reactions The heme group Balancing Redox Equations • Some redox reactions have equations that must be balanced by special techniques. MnO4- + 5 Fe2+ + 8 H+ ---> Mn2+ + 5 Fe3+ + 4 H2O Mn = +7 Fe = +2 Mn = +2 Fe = +3 Balancing Redox Equations Cu + Ag+ --give--> Cu2+ + Ag Balancing Redox Equations Cu + Ag+ --give--> Cu2+ + Ag Step 1: Divide the reaction into half-reactions, one for oxidation and the other for reduction. Ox Red Step 2: Balance each for mass. Step 3: Balance each half-reaction for charge by adding electrons. Ox Red Balancing Redox Equations Step 4: Multiply each half-reaction by a factor so that the reducing agent supplies as many electrons as the oxidizing agent requires. Reducing agent Oxidizing agent Step 5: Add half-reactions to give the overall equation. The equation is now balanced for both charge and mass. Reduction of VO2+ with Zn Balance the following in ACID solution: VO2+ + Zn ---> VO2+ + Zn2+ Step 1: Write the half-reactions Ox Red Step 2: Balance each half-reaction for mass. Ox Red Add H2O on O-deficient side and add H+ on other side for H-balance. Balancing… Step 3: Balance half-reactions for charge. Ox Red Step 4: Multiply by an appropriate factor. Ox Red Step 5: Add balanced half-reactions Tips on Balancing • Never add O2, O atoms, or O2- to balance oxygen. • Never add H2 or H atoms to balance hydrogen. • Be sure to write the correct charges on all the ions. • Check your work at the • end to make sure mass • and charge are balanced. • PRACTICE! Balance the following in basic solution: MnO4- + NO2- MnO2 + NO3Oxidation half reaction: Balancing… MnO4- + NO2- MnO2 + NO3Reduction half reaction: Balancing… MnO4- + NO2- MnO2 + NO3Oxidation half reaction: Reduction half reaction: Multiply by appropriate factor to cancel e- and add both half-reactions Oxid. X Red. X Sum: Study Exp 11 – Procedure to balance redox reactions – practice and E calculation Chemical Change ---> Electric Current With time, Cu plates out onto Zn metal strip, and Zn strip “disappears.” Electrons are transferred from Zn to Cu2+, but there is no useful electric current. Oxidation: Zn(s) ---> Zn2+(aq) + 2eReduction: Cu2+(aq) + 2e- ---> Cu(s) -------------------------------------------------------Cu2+(aq) + Zn(s) ---> Zn2+(aq) + Cu(s) Chemical Change ---> Electric Current • To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire. • This is accomplished in a GALVANIC or VOLTAIC cell. • A group of such cells is called a _____________. Chemical Change ---> Electric Current • Electrons travel thru external wire. • _____________ allows anions and cations to move between electrode compartments. Zn --> Zn2+ + 2eOxidation Anode Negative Cu2+ + 2e- --> Cu <--Anions Cations--> Reduction Cathode Positive The Cu|Cu2+ and Ag|Ag+ Cell Electrochemical Cell • _________ move from anode to cathode in the wire. • _______ & _________move thru the salt bridge. Terminology Figure 20.6 What Voltage does a Cell Generate? • Electrons are “driven” from anode to cathode by an _______________________or emf. • For Zn/Cu cell, this is indicated by a voltage of 1.10 V at 25 ˚C and when [Zn2+] and [Cu2+] = 1.0 M. 1.10 V Zn and Zn2+, anode Cu and Cu2+, cathode 1.0 M 1.0 M Cell Potential, E For Zn/Cu cell, potential is +1.10 V at 25 ˚C and when [Zn2+] and [Cu2+] = 1.0 M. This is the STANDARD CELL POTENTIAL, Eo —a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25 ˚C. Calculating Cell Voltage • Balanced half-reactions can be added together to get overall, balanced equation. Zn(s) ---> Zn2+(aq) + 2eCu2+(aq) + 2e- ---> Cu(s) --------------------------------------------------------Cu2+(aq) + Zn(s) ---> Zn2+(aq) + Cu(s) • If we know Eo for each halfreaction, we could get Eo for net reaction. • But we need a reference! Cell Potential: SHE • Can’t measure 1/2 reaction Eo directly. Therefore, measure it relative to a ______________________, SHE. 2 H+(aq, 1 M) + 2e- <----> H2(g, 1 atm) Eo = 0.0 V Zn/Zn2+ half-cell hooked to a SHE. Eo for the cell = +0.76 V Zn(s) ---> Zn2+(aq) + 2eNegative electrode Positive electrode Supplier of electrons Acceptor of electrons Zn --> Zn2+ + 2eOxidation Anode 2 H+ + 2e- --> H2 Reduction Cathode Reduction of Protons (H+) by Zn Overall reaction is reduction of H+ by Zn metal. Zn(s) + 2 H+ (aq) --> Zn2+ + H2(g) Eo = +0.76 V Therefore, Eo for Zn ---> Zn2+ (aq) + 2e- is +0.76 V Zn is a (better) (poorer) reducing agent than H2. Cu/Cu2+ and H2/H+ Cell, E0 for the cell = + 0.34 V Cu2+(aq) Eo = +0.34 V + 2e- ---> Cu(s) e- e- Positive Negative Acceptor of electrons Supplier of electrons Cu2+ + 2e- --> Cu Reduction Cathode H2 --> 2 H+ + 2eOxidation Anode Overall reaction is reduction of Cu2+ by H2 gas. • Cu2+ (aq) + H2(g) ---> Cu(s) + 2 H+(aq) • Measured Eo = +0.34 V • Therefore, Eo for Cu2+ + 2e- ---> Cu is + 0.34 V Zn/Cu Electrochemical Cell + Anode, negative, source of electrons • Zn(s) ---> Zn2+(aq) + 2eEo = +0.76 V • Cu2+(aq) + 2e- ---> Cu(s) Eo = +0.34 V --------------------------------------------------------------• Cu2+(aq) + Zn(s) ---> Zn2+(aq) + Cu(s) Cathode, positive, sink for electrons Uses of Eo values • Organize half-reactions by relative ability to act as oxidizing agents. Cu2+(aq) + 2e- ---> Cu(s) Eo = +0.34 V Zn2+(aq) + 2e- ---> Zn(s) Eo = –0.76 V Note that when a reaction is reversed the sign of E˚ is reversed! Cu2+ is better oxidazing agent than Zn2+; Cu2+ will be reduced and Zn will be oxidized Cu2+ reaction (reduction) will occur at the cathode Zn reaction (oxidation) will occur at the anode. E˚net = E˚cathode - E˚anode = (0.43) – (-0.76) = 1.1 V Std. Reduction Potentials • Organize halfreactions by relative ability to act as oxidizing agents. (All references are written as reduction processes). Table 20.1 • Use this to predict direction of redox reactions and cell potentials. Potential Ladder for Reduction HalfReactions Best oxidizing agents Figure 20.14 Best reducing agents Using Standard Potentials, Eo Which is the best oxidizing agent: O2, H2O2, or Cl2? Which is the best reducing agent: Hg, Al, or Sn? oxidizing ability of agent Cu2+ + 2e- Eo (V) Cu +0.34 + 2 H + 2e- H2 0.00 2+ Zn + 2e- Zn -0.76 reducing ability of agent Standard Reduction Potentials Any substance on the right will reduce any substance higher than it on the left. Zn can reduce H+ and Cu+. H2 can reduce Cu2+ but not Zn2+ Cu cannot reduce H+ or Zn2+. Standard Reduction Potentials Ox. agent Cu2+ + 2e- --> Cu + 2H + 2e- --> H2 Zn2+ + 2e- --> Zn +0.34 0.00 -0.76 Red. agent Any substance on the right will reduce any substance higher than it on the left. Northwest-southeast rule: product-favored reactions occur between • reducing agent at southeast corner • oxidizing agent at northwest corner Using Standard Reduction Potentials In which direction do the following reactions go? • Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s) • 2 Fe2+(aq) + Sn2+(aq) ---> 2 Fe3+(aq) + Sn(s) • What is Eonet for the overall reaction? Calculating Cell Potential E˚net = “distance” from “top” half-reaction (cathode) to “bottom” half-reaction (anode) E˚net = E˚cathode - E˚anode Eonet for Cu/Ag+ reaction = +0.46 V Eo for a Cell Cd --> Cd2+ + 2eor Cd2+ + 2e- --> Cd Fe --> Fe2+ + 2eor Fe2+ + 2e- --> Fe All ingredients are present. Which way does reaction proceed? Eo for a Cell From the table, you see • Fe is a better reducing agent than Cd • Cd2+ is a better oxidizing agent than Fe2+ Overall reaction: Fe + Cd2+ ---> Cd + Fe2+ Eo = E˚cathode - E˚anode = = Fe/Fe2+ // Cd2+/Cd More about Eo for a Cell Assume I- ion can reduce water. 2 H2O + 2e- ---> H2 + 2 OHCathode 2 I- ---> I2 + 2eAnode ------------------------------------------------2 I- + 2 H2O --> I2 + 2 OH- + H2 Assuming reaction occurs as written, E˚net = E˚cathode - E˚anode = _________ E˚ means rxn. occurs in ___________ direction It is ____________ favored. Eo at non-standard conditions E = Eo - (RT/nF) lnQ The NERNST E = Eo - 0.0257/n lnQ EQUATION E = potential under nonstandard conditions R = gas constant (8.314472 J/Kmol) T = temperature (K) n = no. of electrons exchanged F = Faraday constant (9.6485338 x 104 C/mol) ln = “natural log” Q = reaction quotent (concentration of products/concentration of reactants to the appropriate power) One ___________ is the quantity of electric charge carried by one mole of electrons. If [P] and [R] = 1 mol/L, then E = E˚ If [R] > [P], then E is ______________ than E˚ If [R] < [P], then E is ______________ than E˚ A voltaic cell is set up at 25oC with the following half-cells: Al3+(0.0010 M)/Al and Ni2+(0.50 M)/Ni. Write an equation for the reaction that occurs when the cell generates an electric current. a) Determine which substance is oxidized (decide which is the better reducing agent). Al is best reducing agent. Then Al is oxidized and Ni2+ is reduced. Ox (Anode): Red (Cathode): b) Add the half-reactions to determine the net ionic equation. Net eq: c) Calculate Eo and use Nernst eq. to calculate E. Eo = Eocathode – Eoanode E = Eo – 0.0257/n ln Q Calculate the cell potential, at 25 °C, based upon the overall reaction: 3 Cu2+(aq) + 2 Al(s) 3 Cu(s) + 2 Al3+(aq) if [Cu2+] = 0.75 M and [Al3+] = 0.0010 M. The standard reduction potentials are as follows: Cu2+(aq) + 2 e- → Cu(s) E° = +0.34 V Cathode Al3+(aq) + 3 e- → Al(s) E° = -1.66 V Anode Eo = Eocathode – Eoanode E = Eo – 0.0257/n ln Q Eo and Thermodynamics DE = q + w The maximum work done by an electrochemical system (ideally) is proportional to the potential difference (volts) and the quantity of charge (coulombs): Wmax = nFE E is the cell voltage nF is the quantity of electric charge transferred from anode to cathode. • Eo is related to ∆Go, the free energy change for the reaction (energy released by the cell); under standard conditions: ∆Go = -nFEo where F = Faraday constant = 9.6485 x 104 J/V•mol of e(or 9.6485 x 104 coulombs/mol) and n is the number of moles of electrons transferred Calculate DGo from Eo • Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s) Eonet for Cu/Ag+ reaction = +0.46 V DGo = -nFEo 1J = 1C * 1V 1000 J = 1kJ Eo and the Equilibrium Constant When Ecell = 0, the reactants and products are at equilibrium, Q = K E = 0 = Eo – 0.0257/n ln K then ln K = n Eo / 0.0257 (at 25oC) For: Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s) Eonet for Cu/Ag+ reaction = +0.46 V Eo and the Equilibrium Constant ∆Go = - n F Eo For a product-favored reaction Reactants ----> Products ∆Go < 0 and so Eo > 0 Eo is positive For a reactant-favored reaction Reactants <---- Products ∆Go > 0 and so Eo < 0 Eo is negative Eo = - DG0 nF Primary batteries Uses redox reactions that cannot be restored by recharge. * Indicate which reaction goes in the anode which in the cathode. Dry cell battery: _____________ Zn ---> Zn2+ + 2e_____________ 2 NH4+ + 2e- ---> 2 NH3 + H2 Alkaline batteries Nearly same reactions as in common dry cell, but under basic conditions. _______________ Zn + 2 OH- ---> ZnO + H2O + 2e_______________ 2 MnO2 + H2O + 2e- ---> Mn2O3 + 2 OH- Secondary batteries • Uses redox reactions that can be reversed. • Can be restored by recharging. Lead storage batteries ___________ Eo = +0.36 V Pb + HSO4- ---> PbSO4 + H+ + 2e___________ Eo = +1.68 V PbO2 + HSO4- + 3 H+ + 2e- ---> PbSO4 + 2 H2O Ni-Cd battery ______________ Cd + 2 OH- ---> Cd(OH)2 + 2e______________ NiO(OH) + H2O + e- ---> Ni(OH)2 + OH- Fuel Cell: H2 as fuel • • • • Reactants are supplied continuously from an external source. Cars can use electricity generated by H2/O2 fuel cells. H2 carried in tanks or generated from hydrocarbons. Used in space rockets. Fuel Cell: H2 as fuel Cathode (red) Anode (ox) O2 (g) + 2 H2O (l) + 4e- 4 OH2H2(g) 4 H+ (aq) + 4e- • Temperature of 70-140oC and produce ~ 0.9 V. • The two halves are separated by a proton exchange membrane (PEM). • Protons combine with hydroxide ions forming water. • The net reaction is then: 2 H2 + O2 2 H2O Electrolysis Electric Energy ----> Chemical Change 2 H2O 2 H2 + O2 ________________ 4 OH- ---> O2(g) + 2 H2O + 4e________________ 4 H2O + 4e- ---> 2 H2 + 4 OHEo for cell = -1.23 V Anode Cathode Electrolysis of Molten NaCl • Electrolysis of molten NaCl. • Here a battery “pumps” electrons from Cl- to Na+. + • NOTE: Polarity of Anode electrodes is reversed from batteries. ________________ 2 Cl- ---> Cl2(g) + 2e________________ Na+ + e- ---> Na electrons BATTERY Cathode Cl- Na+ Electrolysis of Molten NaCl Eo for cell (in water) =E˚c - E˚a = - 2.71 V – (+1.36 V) = - 4.07 V (in water) External energy needed because Eo is (-). Electrolysis of Aqueous NaCl Anode (+) 2 Cl- ---> Cl2(g) + 2eCathode (-) 2 H2O + 2e- ---> H2 + 2 OHEo for cell = -2.19 V Note that H2O (-0.8277) is more easily reduced than Na+ (-2.71). Also, Cl- (1.36) is oxidized in preference to H2O (1.33) because of kinetics. Electrolysis of Aqueous CuCl2 Anode (+) 2 Cl- ---> Cl2(g) + 2eCathode (-) Cu2+ + 2e- ---> Cu Eo for cell = -1.02 V Note that Cu is more easily reduced than either H2O or Na+. Electrolytic Refining of Copper Impure copper is oxidized to Cu2+ at the _________. The aqueous Cu2+ ions are reduced to Cu metal at the _______________. Electrolysis of Aqueous SnCl2 Sn2+(aq) + 2 Cl-(aq) ---> Sn(s) + Cl2(g) Eocell = Eocathode-Eoanode = Al production 2 Al2O3 + 3 C ---> 4 Al + 3 CO2 Charles Hall (1863-1914) developed electrolysis process. Founded Alcoa. Counting electrons • The number of e- consumed or produced in an electron transfer reaction is obtained by measuring the current flowing in the circuit in a given time. • The current flowing is the amount of charge (coulombs, C) per unit time, the unit is the ampere (A). 1 A = 1 C/s then 1C = A *s 1 F = 9.6485338 x 104 C/mol e1 mol e- = 96,500 C 1.50 amps flow thru a Ag+(aq) solution for 15.0 min. What mass of Ag metal is deposited? a) Calculate the charge Charge (C) = current (A) x time (t) b) Calculate moles of e- used c) Calculate the mass The anode reaction in a lead storage battery is Pb(s) + HSO4-(aq) ---> PbSO4(s) + H+(aq) + 2eIf a battery delivers 1.50 amp, and you have 454 g of Pb, how long will the battery last? a) Calculate moles of Pb b) Calculate moles of e- c) Calculate charge (C): d) Calculate time Time (sec) = Charge (C) I (amps) End of Chapter • Go over all the contents of your textbook. • Practice with examples and with problems at the end of the chapter. • Practice with OWL tutor. • Practice with the quiz on CD of Chemistry Now. • Work on your OWL assignment for Chapter 20.