JHU_GL12_Lab (1).docx - 2 ELECTROCHEMISTRY ELECTROLYTIC...

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12. ELECTROCHEMISTRY - ELECTROLYTIC CELLSJHU INTRO. CHEM. LABSUMMER 2017In this experiment you will coat a brass key withcopper using electroplating with an electrolytic cell.This is a common manufacturing technique used, forexample, to plate silver on utensils, silver and goldon jewelry, and chrome on bumpers. A brief reviewof electrolytic cells is presented here and in the on-line prelab lecture. You should also refer to yourtextbook (Atkins, Jones, and Laverman, ChemicalPrinciples, 6th ed., Ch. 14.11). In addition, there aresome excellent animations of electrolytic cells on theInternet. One link is given in the information postedon Blackboard. In particular, as you prepare for theexperiment, you should consider which terminal onthe battery should be connected to the brass key inorder to plate copper onto the key.ObjectivesConstruct an electrolytic cell and identify its parts.This includeso the power sourceo the direction of electron flowo the chemical half-reactions that take place at the electrodeso the identity of the anode and the cathodeo the contents of the beaker.Calculate the total charge consumed in theelectrolytic cell.Calculate the identity and amount of materialplated or consumed at the electrodes.Apply the understanding of the chemicalprocesses and procedures of the experiment to anerror analysis of your experiment.Distinguish the similarities and differencesbetween electrolytic cells and galvanic cells. Asummary of these similarities and differences canbe found on Blackboard and also in the onlinepre- lab video lecture.British chemist, first showed that plating one mole ofsingly charged ions out of solution requires one moleof electrons. For example, in the reaction Ag+(aq) + e-Ag(s) occurring at the cathode in the electrolysis of AgNO3, one mole of electrons is required to neutralizeBACKGROUND INFORMATIONCurrent is the name for the flow of electrons.These electrons can cause half-reactions to occur inan electrochemical cell. Passing an electric currentthrough a solution causes chemical reactions to occurat the electrodes. Michael Faraday, a 19th century
2the positive charge on one mole of silver ions andplate one mole of silver metal on the electrode.However, electricity is not measured in number ofmoles of electrons but in units of coulombs. Onecoulomb (C) of electric charge, Q, is the quantity ofelectricity that flows through a wire in one second ifthe electric current, I, is one ampere (A). Thenumber of coulombs is the product of the current inamperes times the time in seconds:Q = I·t(1)The number of coulombs required to plate one moleof singly charged ions out of solution is known asFaraday's constant (F) or "the Faraday". TheFaraday, 96,485.3399 C/mol, is the total electricalcharge carried by one mole of electrons.I·t = moles of electrons(2)FThe value of the Faraday has been measuredexperimentally by determining the number ofcoulombs required to plate a known mass of ionsout of solution. Faraday's Laws and the Faraday

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