Lab #4 Instructions

Lab #4 Instructions - EXPERIMENT IV Oxidation-Reduction...

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EXPERIMENT IV Oxidation-Reduction Reactions PURPOSE: Determine the cell voltage for several oxidation-reduction reactions and determine how concentration and pH affect the potential (free energy) of these reactions. INTRODUCTION: This experiment is related to material that will be covered in lecture over the next few weeks. The goal of the exercises is to provide practical experience with electrochemical cells, to give some background knowledge for the concepts and calculations introduced in lecture. To accomplish this, you will: Construct an electrochemical cell Learn to use your pH meter is a voltmeter; Measure the cell voltages of several electrochemical cells; Determine the effect of pH on cell voltage; Determine the effect of concentration on cell voltage; Determine Ksp for a salt by an electrochemical approach Part 1: Building an Electrochemical Cell An electrochemical cell consists of two half-cells connected via an external circuit and an internal electrolyte (salt bridge). An electrochemical cell allowed to run spontaneously is referred to as a galvanic cell . In a galvanic cell, the anode material (usually a metal) will oxidize more readily than the cathode material. The classic example is the zinc/copper cell depicted below. Zinc oxidizes more readily than copper. Therefore the anode reaction is: Zn(s) Zn 2+ (aq) + 2e - anode reaction The electrons given off by zinc at the anode travel through the external circuit to the copper cathode. At the cathode the excess electrons reduce copper ions to copper metal. Cu 2+ + 2e - Cu(s) cathode reaction If the galvanic cell is allowed to reach equilibrium most of the Cu 2+ ions will have been plated out on the copper cathode, and the zinc anode will have been partially corroded away. At equilibrium the potential of the cell will be zero. A charged galvanic cell will exhibit a positive potential. We can
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measure this potential by placing a voltmeter between the anode and cathode in the external circuit. The voltmeter not only measures the potential, it also drastically slows down the flow of electrons. This enables a stable voltage to be measured. The portable pH meter you have been using is really a voltmeter. When attached to a pH electrode the instrument measures the difference in potential between a silver/silver chloride half-cell and a proton concentration cell (more on this later). When you calibrate the pH meter using pH 4, 7 and 10 buffer solutions, you are really setting the reference potential for the electrochemical cell. Potential, like enthalpy or free energy, doesn’t have an absolute value. It must always be referenced to some reproducible standard state. From here on out, we will refer to potential using the symbol E . Therefore any measurement is understood to mean E E ref (potential relative to the reference potential).
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Lab #4 Instructions - EXPERIMENT IV Oxidation-Reduction...

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