ElectrochemistryRedox Reactions：electrochemistryredox反应.doc

Electrochemical Cells1. Using reduction potentials answer the following. a. What is the strongest oxidizing agent? b. What is the strongest reducing agent? c. Will Cl2 be able to reduce Cr3+? d. Will Pb2+ oxidize Ni? e. Will Au dissolve in an HCl solution? f. Will Zn dissolve in an HCl solution? g. What can oxidize Al but not Fe?h. What can reduce I2 but not H+?2. Consider the following reaction at 75 C.Pb2+(aq) + 2 Cr2+(aq) Pb(s) + 2 Cr3+(aq)a. Calculate the standard potentialb. Calculate the standard change in free energy c. Calculate the equilibrium constant d. If the above cell is allowed to operate spontaneously, will the voltage increase, decrease or remain the same?e. If the above cell is allowed to operate spontaneously, what will happen to the concentrations at the electrodes?f. Calculate the initial cell voltage for the above reaction if the initial concentrations are Pb2+ = 0.25 M, Cr2+ = 0.20M and Cr3+ = 0.005 M.3. Consider the following cell: Cu (s) | Cu2+ (0.001 M) | Br2 (l) | Br-(aq) (? M) a. Assign the electrodesb. Determine the direction of electron flow and the direction of the currentc. Describe the flow of ions in the salt bridged. Calculate the standard potentiale. At 25 C the measured cell voltage is 0.975 V. Calculate the concentration of the bromide ion.4. Consider the following cell: Al(s) | Al3+ (1.0 M) | | Pb2+ (1.0 M) | Pb(s) Calculate the cell potential after the reaction has operated long enough for the Al3+ to have changed by 0.3 M at 25 C.5. Consider a cell with 0.0001 M Fe2+/Fe (s) and 10 M Fe2+/Fe(s). Determine the direction of electron flow, assign the electrodes and calculate the cell voltage at 25 C. What are the concentrations at equilibrium?6. What mass of Co forms from a solution of Co2+ when a current of 15 amps is applied for 1.15 hours?7. A solution of iron chloride underwent electrolysis for 2 hrs at 10 amps yielding 20.84 g of iron. What is the molecular formula for the iron chloride? 8. Consider a cell with 0.0001 M Fe2+/Fe (s) and 10 M Fe2+/Fe(s). Determine the direction of electron flow, assign the electrodes and calculate the cell voltage at 25 C. 9. What volumes of H2 and O2 at STP are produced from the electrolysis of water by a current of 3.5 amps in 15 minutes? Hierarchy for Assigning Oxidation Numbers/States1. Atoms in their elemental form = 02. Monoatomic ions = charge3. The sum of the oxidation numbers for compounds = 04. The sum of the oxidation numbers for polyatomic ions = net charge5. Hydrogen in a compound or polyatomic ion = +16. Fluorine in a compound = -17. Oxygen in a compound or polyatomic ion = -28. Chlorine, Bromine and Iodine = -19. Nitrogen in a compound or polyatomic ion = -3Red-ox Reactions1. A reaction where there is a transfer of electrons2. Oxidation is when an element loses electrons and is noted by an increase in the oxidation number 3. Reduction is when an element gains electron and is noted by a decrease in the oxidation number4. The number of electrons gained must equal the number of electrons lost5. An oxidizing agent is the substance getting reduced6. A reducing agent is the substance getting oxidizedElectrochemical Cells1. Galvanic or Voltaic spontaneously produce current ( 0)2. Electrolytic or Electrolysis require a current to run a non-spontaneous redox ( 0)3. Anode where oxidation occurs if the electrode is an active metal it will corrode as the cell discharges anions in the salt bridge flow to this side4. Cathode where reduction occurs if a metal ion is reduced to its ground state it will plate out on the electrode cations in the salt bridge flow to the side5. For all cells electrons flow from anode to cathode or the current flows from cathode to anodeBalancing Red-ox Reactions1. Separate into two half reactions2. Balance all elements EXCEPT H and O3. Balance the oxygen by adding H204. Balance the hydrogen by adding H+5. Balance the net charges by adding electrons (e-) to the more positive side6. Make the electrons lost equal to the electrons gained in each half reaction by multiplying the half reactions7. Add two half reactions back together canceling out electrons, H+ and H2O8. If acidic stop. If basic add an OH- to each side for every H+. The OH- cancels out the H+ making water, which then needs to be adjusted for. Table of Standard Reduction PotentialsHalf-Reaction(volts)F2 + 2 e- 2 F-2.87H2O2 + 2 H+ + 2 e- 2 H2O 1.78MnO4- + 8 H+ + 5 e- Mn2+ + 4 H2O 1.49Au3+ + 3 e- Au1.42Cl2 + 2 e- 2 Cl-1.36Cr2O72- + 14 H+ + 6 e- 2 Cr3+ + 7 H2O1.33O2 + 4 H+ + 4 e- 2 H2O1.23Br2 + 2 e- 2 Br-1.06NO3- + 4 H+ + 3 e- NO + 2 H2O0.96Ag+ + e- Ag0.80Fe3+ + e- Fe2+0.77I2 + 2 e- 2 I-0.53Cu+ + e- Cu0.52O2 + 2 H2O + 4 e- 4 OH-0.40Cu2+ + 2 e- Cu0.34Cu2+ + e- Cu+0.162 H+ + 2 e- H20.00Pb2+ + 2 e- Pb-0.13Sn2+ + 2 e- Sn-0.14Ni2+ + 2 e- Ni-0.23Co2+ + 2 e- Co-0.28Fe2+ + 2 e- Fe-0.41Cr3+ + e-