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ELECTROCHEMICAL CELLS ADDITIONAL READING The concepts in this experiment are also discussed in sections 18.3 - 18.6 in Principles of Chemistry - A...

Here is a lab that I will be doing this week I need you to solve the PRE-LABORATORY QUESTIONS.

1 ELECTROCHEMICAL CELLS ADDITIONAL READING The concepts in this experiment are also discussed in sections 18.3 – 18.6 in Principles of Chemistry – A Molecular Approach , by Tro. ABSTRACT This experiment is divided into three parts. In the part one you will prepare a Cu||Pb 2+ cell and measure its potential. In part two you will test a voltaic cell that uses an unknown metal electrode and identify the metal. In the part three you will prepare a Ag concentration cell and measure its potential. Using this potential and the Nernst equation you will calculate the Ksp of AgCl BACKGROUND Electrochemistry deals with 1) the chemical changes produced by electric current, and 2) the production of electricity by chemical reactions. In this experiment, you will focus on the second topic. All electrochemical reactions involve the transfer of electrons, and so are oxidation-reduction reactions. In order to measure the electron transfer (voltage, current, etc.), the sites of oxidation and reduction must be separated physically; oxidation occurs at one location, while reduction occurs at the other. Electrochemical processes require some method of introducing a stream of electrons (from the species being oxidized) into a reacting system and some means of withdrawing electrons (by the species being reduced). In most applications, the reacting system is contained in a cell , and an electric current enters or exits at electrodes . Electrochemical cells are classified into types, they are: 1. Electrolytic cells : These cells use electrical energy from an external source, and cause non-spontaneous chemical reactions to occur. 2. Voltaic cells : These cells produce electricity from spontaneous chemical reactions and supply it to an external circuit. When a strip of zinc metal is placed into a solution containing copper ions, a brownish metal coats the strip and the blue color of the original solution begins to fade. This occurs because the copper ions are being reduced yielding brown copper metal, and the zinc metal is being oxidized. The color of the solution fades because the blue copper(II) ions are replaced by colorless zinc ions. The net ionic equation for this reaction is shown in equation (1): Zn(s) + Cu 2+ (aq) Cu(s) + Zn 2+ (aq) (1) The reaction shown in equation (1) is an oxidation-reduction (redox) reaction. Although oxidation may be defined as a loss of electrons, and reduction as a gain of electrons, no electron flow is detected in the system described above. However, we can conclude from this qualitative experiment that zinc metal is a stronger reducing agent than copper metal , and so the zinc will reduce the copper ions to copper metal . Conversely, copper ions are stronger oxidizing agents than zinc ions , and so the copper ions oxidize the zinc metal to zinc ions . Also recall that oxidizing agent causes oxidation by itself undergoing reduction, and that reducing agent causes reduction by itself undergoing oxidation . The Zn||Cu 2+ galvanic cell that will be discussed below is an example of a voltaic cell. The spontaneous oxidation- reduction reaction between zinc and copper produces electrical energy. The two halves of the redox reaction are separated, requiring electron transfer to occur through an external circuit. In this way, useful electrical energy is obtained. The dry cells commonly used in flashlights, iPods, photographic equipment and many toys and appliances are voltaic cells. Automobile batteries consist of voltaic cells connected in series so that their voltages add. If we wish to use the Zn/Cu 2+ (aq) system to generate electricity (i.e., produce a measurable flow of electrons), we must use a more elaborate set-up. This apparatus, an electrochemical cell, would consist of a zinc electrode, placed into a solution containing zinc ions, and a copper electrode placed into a solution containing copper ions. A
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2 wire would connect the electrodes. Putting a voltmeter into this circuit would permit a measurement of the potential (the tendency of electrons to flow from one electrode to the other) of the cell. This electrochemical cell is shown in figure 1 below. Figure 1: The Zn-Cu galvanic H-cell The reduction half-reaction is confined to the right side of the cell; the oxidation half-reaction is confined to the left side. The two sides are connected by a tube which is blocked by a porous glass disc (sometimes called a sintered glass disc or a glass frit). The holes in the disc are large enough to allow ions to pass through, but they are small enough to discourage ionic transfer by means of convection. It is important that the Cu 2+ (aq) ions in the right compartment not be allowed to move towards the Zn(s) in the left compartment; if this happens, the reaction can take place without requiring the electrons to pass though the wire, defeating the purpose of the cell. For obvious reasons (the shape of the cell), the reaction vessel used in this type of operation is often called an H cell. An important feature of this cell is the porous glass disk. In less sophisticated cells the porous glass disk is referred to as a salt bridge.
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