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- An electrolysis experiment is performed to determine the value of the Faraday constant (number of coulombs per mole of electrons). In this experiment, 28.8 g of gold is plated out from a AuCN solution by running an electrolytic cell for two hours with a current of 2.00 A. What is the experimental value obtained for the Faraday Constant?arrow_forwardAn electrochemical cell is made by placing a zinc electrode in 1.00 L of 0.200 M ZnSO4 solution and a copper electrode in 1.00 L of 0.0100 M CuCl2 solution. a What is the initial voltage of this cell when it is properly constructed? b Calculate the final concentration of Cu2+ in this cell if it is allowed to produce an average current of 1.0 amp for 225 s.arrow_forwardAn electrode is prepared by dipping a silver strip into a solution saturated with silver thiocyanate, AgSCN, and containing 0.10 M SCN . The cell potential of the voltaic cell constructed by connecting this electrode as the cathode to the standard hydrogen half-cell as the anode is 0.45 V. What is the solubility product of silver thiocyanate?arrow_forward
- An aqueous solution of an unknown salt of vanadium is electrolyzed by a current of 2.50 amps for 1.90 hours. The electroplating is carried out with an efficiency of 95.0%, resulting in a deposit of 2.850 g of vanadium. a How many faradays are required to deposit the vanadium? b What is the charge on the vanadium ions (based on your calculations)?arrow_forwardChlorine, Cl2, is produced commercially by the electrolysis of aqueous sodium chloride. The anode reaction is 2Cl(aq)Cl2(g)+2e How long will it take to produce 2.00 kg of chlorine if the current is 5.00 102 A?arrow_forwardAn electrode is prepared from liquid mercury in contact with a saturated solution of mercury(I) chloride, Hg2Cl, containing 1.00 M Cl . The cell potential of the voltaic cell constructed by connecting this electrode as the cathode to the standard hydrogen half-cell as the anode is 0.268 V. What is the solubility product of mercury(I) chloride?arrow_forward
- Assume the following electrochemical cell simulates the galvanic cell formed by copper and zinc in seawater at pH 7.90 and 25 C. Zn | Zn(OH)2(s) | OH(aq) || Cu(OH)2(s) | Cu(s) a. Write a balanced equation for the reaction that occurs at the cathode. b. Write a balanced equation for the reaction that occurs at the anode. c. Write a balanced chemical equation for the overall reaction. d. Determine the potential (in volts) of the cell.arrow_forwardGiven the following two standard reduction potentials, solve for the standard reduction potential of the half-reaction M3++eM2+ (Hint: You must use the extensive property G to determine the standard reduction potential.)arrow_forwarda Calculate the equilibrium constant for the following reaction at 25C. Sn(s)+Pb2+(aq)Sn2+(aq)+Pb(s) The standard cell potential of the corresponding voltaic cell is 0.010 V. b If an excess of tin metal is added to 1.0 M Pb2+, what is the concentration of Pb2+ at equilibrium?arrow_forward
- What volume of F2 gas, at 25C and 1.00 atm, is produced when molten KF is electrolyzed by a current of 10.0 A for 2.00 h? What mass of potassium metal is produced? At which electrode does each reaction occur?arrow_forwardA galvanic cell is based on the following half-reactions: In this cell, the copper compartment contains a copper electrode and [Cu2+] = 1.00 M, and the vanadium compartment contains a vanadium electrode and V2+ at an unknown concentration. The compartment containing the vanadium (1.00 L of solution) was titrated with 0.0800 M H2EDTA2, resulting in the reaction H2EDTA2(aq)+V2+(aq)VEDTA2(aq)+2H+(aq)K=? The potential of the cell was monitored to determine the stoichiometric point for the process, which occurred at a volume of 500.0 mL H2EDTA2 solution added. At the stoichiometric point, was observed to be 1 .98 V. The solution was buffered at a pH of 10.00. a. Calculate before the titration was carried out. b. Calculate the value of the equilibrium constant, K, for the titration reaction. c. Calculate at the halfway point in the titration.arrow_forwardA voltaic cell is constructed in which one half-cell consists of a silver wire in an aqueous solution of AgNO3.The other half cell consists of an inert platinum wire in an aqueous solution containing Fe2+(aq) and Fe3+(aq). (a) Calculate the cell potential, assuming standard conditions. (b) Write the net ionic equation for the reaction occurring in the cell. (c) Which electrode is the anode and which is the cathode? (d) If [Ag+] is 0.10 M, and [Fe2+] and [Fe3+] are both 1.0 M, what is the cell potential? Is the net cell reaction still that used in part (a)? If not, what is the net reaction under the new conditions?arrow_forward
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