Chapter 16Redox ReactionsSlide 3Problem 16-7Redox Reactions in Electrochemical CellsSlide 6- an array consisting of 2 or 3 electrodes, each of which is in contact with an electrolyte solution. Typically, the electrolytes are in electrical contact through a salt bridge. An external metal conductor connects the electrodes. A cathode is an electrode where reduction occurs An anode is the electrode where oxidation occurs -Electrochemical Cells are either galvanic or electrolytic galvanic cells store and supply electrical energy (ex. Batteries) electrolytic cells require an external source of electrical energy for operationSlide 8Slide 9Electrode PotentialsElectrode PotentialsThe Nerst EquationProblem 16-24Chapter 16Chapter 16Krissy KellockKrissy KellockAnalytical Chemistry 221Analytical Chemistry 221Redox ReactionsRedox Reactions•Redox Reactions-reaction in which electrons are transferred from one reactant to another–Ce4+ + Fe2+ ↔ Ce3+ + Fe3+•Ce is the oxidizing agent or oxidant because it accepts electrons from iron and Fe is the reductant because it donated electrons to Ce.•A redox reaction can be split into 2 half reactions•Ce4+ + e- ↔ Ce3+•Fe2+ ↔ Fe3+ + e-Redox ReactionsRedox Reactions•The 2 half reactions must be balanced just like any other reaction. The number of atoms of each element and the net charge on each side of the equation must be in balance so for the oxidation of Fe2+ by MnO4- the half reactions would be:–MnO4- + 8H+ + 5e- ↔ Mn2+ + 4H2O–5Fe2+ ↔ 5Fe3+ + 5e-•The net charge for the first half reaction is (-1 – 5 + 8) = +2 which is the same on the right side. For the second half reaction it must by multiplied by 5 so that the number of electrons lost by Fe equals the number gained by MnO4. The balanced net equation for the reaction would be:–MnO4- + 5Fe2+ + 8H+ ↔ Mn2+ + 5Fe3+ + 4H2OProblem 16-7•a) 2Fe3+ + Sn2+  2Fe2+ + Sn4+•b) Cr + 3Ag+  Cr3+ + 3Ag•c) 2NO3- + Cu + 4H+  2NO2 + 2H2O + Cu2+•d) 2MnO4- + 5H2SO3  2Mn2+ + 5SO42- + 4H+ + 3H20•e) Ti3+ + Fe(CN)63- + H2O  TiO2+ + Fe(CN)64- + 2H+•f) H2O2 + 2Ce4+  O2 + 2Ce3+ + 2H+•g) 2Ag + 2I- + Sn4+  2AgI + Sn2+•h) UO22+ + Zn + 4H+  U4+ + Zn2+ + 2H2O•i) 5HNO2 + 2MnO4- + H+  5NO3- + 2Mn2+ + 3H2O•j) H2NNH2 + IO3- +2H+ + 2Cl  N2 + ICl2- + 3H2ORedox Reactions in Redox Reactions in Electrochemical CellsElectrochemical Cells•Types- –reaction is performed by direct contact between the oxidant and the reductant in a suitable container.–The reactants do not come in contact with one another, salt bridge employed.- a salt bridge prevents the mixing of the contents of the 2 electrolyte solutions making up electrochemical cellsSource: www.public.asu.edu/~laserweb/woodbury/classes/chm341/lecture_set10/lecture10.html- an array consisting of 2 or 3 electrodes, each of - an array consisting of 2 or 3 electrodes, each of which is in contact with an electrolyte solution. which is in contact with an electrolyte solution. Typically, the electrolytes are in electrical contact Typically, the electrolytes are in electrical contact through a salt bridge. An external metal through a salt bridge. An external metal conductor connects the electrodes.conductor connects the electrodes.A cathode is an electrode where reduction A cathode is an electrode where reduction occursoccursAn anode is the electrode where oxidation An anode is the electrode where oxidation occursoccurs-Electrochemical Cells are either galvanic or -Electrochemical Cells are either galvanic or electrolyticelectrolyticgalvanic cells store and supply electrical galvanic cells store and supply electrical energy (ex. Batteries)energy (ex. Batteries)electrolytic cells require an external source electrolytic cells require an external source of electrical energy for operationof electrical energy for operationElectrochemical CellSource: www.public.asu.edu/~laserweb/woodbury/classes/chm341/lecture_set10/lecture10.htmlSource: www.public.asu.edu/~laserweb/woodbury/classes/chm341/lecture_set10/lecture10.htmlElectrode Potentials•the potential difference that develops between the electrodes of the cell is a measure of the tendency for the reaction to proceed from a nonequilibrium state to the condition of equilibrium.•Gibb’s Free Energy–ΔG = -nFEcellElectrode PotentialsElectrode Potentials•If the reactants and products are in their standard states the resulting cell potential is called the standard cell potential and is related to Gibb’s Free Energy Equation:ΔG˚ = -nFE˚ = -RT ln KeqThe Nerst Equation•E = E˚ - (0.0592 / n) log [C]c[D]d [A]a[B]bProblem 16-24Problem 16-24•E = -0.763 – 0.0296 log 1/[Zn2+]• [ZnY2-] = 3.2x1016  [ZnY2-]•[Y4-][Zn2+] [Y4-]3.2x1016•E = -0.763 – 0.0296 log [Y4-]3.2x1016 [ZnY2-]•when [Y4-] = [ZnY2-] = 1.00 E = E˚ZnY2-E = E˚ZnY2- = -0.763 – 0.0296 log (1.00) (3.2x1016/1.00) = -1.25