Chapter 14 slides - 2S.pdf - Electrochemistry Oxidation...

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Unformatted text preview: Electrochemistry Oxidation numbers and balancing equations Electrochemical cells and cell parts Energetics 1 Redox Reactions Classifications: Direct – No external circuit Indirect – external circuit Galvanic / voltaic cell -­ Produces electrical energy Electrolytic cell – Requires electrical energy 2 Direct Redox Reactions Oxidizing and reducing agents in direct contact. Cu (s) + 2 Ag+ (aq) Cu 2+ (aq) + 2 Ag (s) 3 Direct Redox Reactions Zn metal Cu2+ ions 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) + 2e– Reduction: Cu2+(aq) + 2e– Cu(s) -­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­-­ Cu2+(aq) + Zn(s) Zn2+(aq) + Cu(s) 4 Indirect Redox Reactions To obtain a useful current, we separate the oxidizing and reducing agents. -­ No direct contact. -­ Electron transfer occurs through an external wire. -­ Convert chemical to electrical energy Electrochemical Cells Electrochemical Cell An apparatus that allows a redox reaction to occur by transferring electrons through an external connector (circuit). Galvanic / voltaic cell Chemical change produces electric current ΔGrxn < 0 Product favored reaction Batteries are voltaic / galvanic cells Electrolytic cell Electric current used to cause chemical change in a non-­ spontaneous reaction ΔGrxn > 0 Reactant favored reaction Chrome plate car bumpers Gold / silver plate jewelry, etc. 6 Galvanic Cell Parts Anode: Zn(s) Zn2+(aq) + 2 e– Cu2+(aq) + 2e– Cu2+(aq)+ Zn(s) Cu(s) Cathode: wire electrons Zn Zn2+(aq)+ Cu(s) Zn2+ ions Left-­hand solution Convention Zn(s) Cu salt bridge Zn2+(aq) + 2 e– Anode: Oxidation occurs at this electrode Cu2+ ions Right-­hand solution Cu2+(aq) + 2e– Cu(s) Cathode: Reduction occurs at this electrode Electrons flow in the wire from anode to cathode Galvanic Cell Parts Ammeter: measures current Solutions [Cu(NO3)2 and AgNO3] are connected by a salt bridge, containing NaNO3. Porous plugs at ends of bridge prevent solutions from mixing but allows ions to pass through. Left: Cu(s) is oxidized -­ Cu2+ dissolves -­ pos. charge builds up -­ ions move Right: Ag+ is reduced -­ Ag precipitates -­ neg. charge builds up -­ ions move 8 Writing Galvanic Cells Anode on the left. Cathode on the right. Electrons flow left to right. Cation in solution Cation in solution at the anode at the cathode Cu(s)|Cu2+(aq)||Ag+(aq)|Ag(s) Anode Salt Bridge Cathode 9 Example Diagram the following galvanic cell, indicating the direction of flow of electrons in the external circuit and the motion of ions in the salt bridge: Pt(s)|Cr2+(aq), Cr3+(aq)||Cu2+(aq)|Cu(s) Write the half-­reactions at each electrode and a balanced equation for the overall reaction in this cell. 10 Electrolytic Cell Downs cell • Increase the electrostatic potential energy of electrons in the cathode to make it flow towards the anode • So anode becomes cathode and vice versa 2 NaCl(l) 2 Na(l) + Cl2(g) 11 Electrochemistry Oxidation numbers and balancing equations Electrochemical cells and cell parts Energetics 12 Energetics • Dual aspect of electrochemical cells realized shortly after cell’s discovery by Alessandro Volta (1800). • Sir Humphrey Davy prepared sodium (Na) and potassium (K) metal using a battery to electrolyze hydroxides. • Faraday (1833) expressed a quantitative relationship between amounts reacted and total electrical charge passing through the cell Faraday’s Laws (of electrolysis) 1. The mass of a given substance produced or consumed at an electrode is proportional to quantity of electrical charge passed through the cell. 2. Equivalent masses of different substances are produced or consumed at an electrode by passing of a given quantity of electric charge. Energetics Charge e on a single electron accurately determined to be: e = 1.60217646 ✕ 10–19 C. Quantity of charge on a mole of electrons: Q = (6.0221420 ✕ 1023 mol–1)(1.60217646 ✕ 10–19 C) = 96,485.34 C mol–1 This quantity of charge is called the Faraday constant: F = 96,485.34 C mol–1 Electric current I (units: amperes) = amount of charge Q (units: coulombs) flowing through circuit per unit time. I = Q / t moles of electrons = I t 96,485 C mol–1 Energetics If the only change is the transfer of electrons, the change in energy of the reaction is the same as the change in energy of the electrons. Δ Energy = Q * V Q is charge V is voltage Q = 1 C V = 1 volt Δ Energy = 1 J Q = 1 electron (1.602 x 10-­19 C) V = 1 volt Δ Energy = 1 eV 1 eV = 1.602 x 10-­19 J Electrochemical energetics are listed by voltage for 1 electron, not energy. If we measure the voltage (voltmeter), current (ammeter) and time then we can calculate -­ number of electrons passed through the electrodes -­ amount of product produced (mass, moles, concentration) -­ amount of reactant consumed (mass, moles, concentration) -­ work performed 15 Example An electrolytic cell is constructed in which the silver ions in silver chloride are reduced to silver at the cathode and copper is oxidized to Cu2+ (aq) at the anode. A current of 0.500 A is passed through the cell for 101 minutes. Calculate the mass of the copper dissolved and the mass of the silver deposited. 16 Practice A galvanic cell is constructed that has a zinc anode immersed in Zn(NO3)2 solution and a platinum electrode immersed in a NaCl solution equilibrated with Cl2(g) at 1 atm and 25°C. A salt bridge connects the two half-­cells. (a) Write a balanced equation for the cell reaction. (b) A steady current of 0.800 A is observed to flow for a period of 25.0 minutes. How much charge passes through the circuit during this time? How many moles of electrons is this charge equivalent to? (c) Calculate the change in mass of the zinc electrode (Zn atomic weight = 65.4). (d) Calculate the volume of gaseous chlorine generated or consumed as a result of the reaction. 17 The Gibb’s Free Energy and Cell Voltage Previously discussed pressure-­volume work involving gases. A new kind of work, electrical work, in electrochemical cells: welec = –Q ΔE Charge Q moves through a potential difference ΔE. Minus sign: convention that work done by the system is negative. Units: Work (joule), charge (coulomb). So ΔE has units J/C. Define volt (V) as joule/coulomb. Compute work for current I flowing for time t: welec = – I t ΔE ΔE is positive for galvanic cells. Gibbs Free Energy and Electrical Work Fundamental relationship between change in free energy at constant T and P and maximum electrical work which reaction can produce: –welec,max = |ΔG| (at constant T and P) Proof: Definition of Gibbs free energy: G = H – TS = E + PV – TS At constant T and P, ΔG = ΔE + P ΔV – T ΔS From 1st law of thermodynamics, ΔE = q + w or ΔE = q + welec – P ΔV (now two kinds of work) PV work: –Pext ΔV = – P ΔV. Electrical work: welec= –Q ΔE Gibbs Free Energy and Electrical Work Substitute into equation for free energy change: ΔG = ΔE + P ΔV – T ΔS = (q + welec – P ΔV) + P ΔV – T ΔS = q + welec – T ΔS Condition of reversibility on galvanic cell: q = qrev = T ΔS Hence, free energy change is: ΔG = welec,rev Maximum electrical work produced when cell operated reversibly. If n moles of electrons (or nF coulombs of charge) passes through circuit when operated reversibly, ΔG = welec = –Q ΔE = –nF ΔE (reversible) Electrical work produced by cell only if ΔG < 0 (equivalently ΔE > 0). Can determine ΔG from measurement of cell voltage. Standard States and Cell Voltage Standard state of a substance: pressure of 1 atm and specified T. • For solutions, standard state of a solute is that for which concentration in an ideal solution is 1 M. • For reactions carried out in electrochemical cells, standard free energy change ΔG° is ΔG° = –nF ΔE° Here, ΔE° is cell voltage (potential difference) of galvanic cell in which reactants and products are in their standard states. Standard cell voltage ΔE° can be computed from ΔG°, and vice versa. Example A galvanic cell, Co|Co2+||Ag+|Ag, is constructed in which the standard cell voltage is 1.08 V. Calculate the free energy change at 25°C when (a) 1 mole, (b) 2.00 g of silver (atomic weight 108) plates out, if all concentrations remain at their standard value of 1.0 M throughout the process. What is the maximum electrical work done by the cell on its surroundings during this experiment? ...
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