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Chapter 24 ­ Coulometry

The Faraday constant is 96,500 coulombs per mole of electrons, which is also equal to 96,500 amp seconds. If we measure the time and current accurately, then we can calculate the amount of an electrochemical reaction that takes place. For example, if we electroplate Cu from CuSO4 solution, we can determine the mass of Cu plated if we know both the current and the time. Let's work one. What mass of Cu will be deposited on the cathode from a CuSO4 solution if a current of 0.100 amps flows for 2.22 minutes?

0.100amp * 2.22 min* 60 sec 1coul 1mole e - 1mole Cu 63.55 g Cu * * * * = 0.00438 g Cu 1 min 1amp sec 96,500coul 2mole e - 1mole Cu

HW Problem 24#1: How long must a current of 0.250 amps flow to deposit 1.00 gram of Cr from a Cr3+ solution? Since chemicals get used up in electrochemical reactions, you can't have both the voltage and the current constant. If the voltage is constant, the current decreases as the chemical concentration decreases. If the current is constant, the voltage changes as chemicals get depleted. This usually results in another electrochemical reaction starting. For example, if we deposit Cu at the cathode under constant current conditions, the Cu2+ ion concentration will decrease as the reaction progresses. Eventually there will not be enough Cu2+ available to maintain the current and something else has to happen. Under normal conditions the electrolysis of water starts and some of the electrons go to that reaction, while others continue the Cu2+ + 2 e- à Cu (s) reaction. Under these conditions the number of coulombs cannot be used to calculate the mass of Cu deposited. In the commercial electroplating world they get around some of these difficulties by the way they set up the system. If we deposit Cu from CuSO 4 solution and at the same time oxidize Cu at a Cu metal anode, the reaction should take place as long as the voltage is sufficient and the Cu anode does not disappear. The voltage required is not 0 even though the same (but opposite) reaction is taking place at the cathode and anode. The applied voltage must be great enough to overcome the double layer capacitance and junction potentials. To make the current larger, the voltage is increased. They also stir the solution vigorously to make the reaction convection limited and not diffusion limited. Many commercial processes are based on electrolysis. In our part of the country the electrolysis is used to produce Al. Alcoa, TN, is a town whose existence depends on the

reduction of Al-containing ores to Al metal. Many of the dams in western NC and eastern TN were built to supply electricity for the Alcoa plant. Since Al is far more reactive than Cu, the electrolysis has to be carried out in the absence of water. A molten salt, cryolite, is used. Copper Hill, TN, is a town where Cu was produced. Electrolysis was used to purify the Cu. Na, K, Mg, Cl2, Br2, and I2 are common chemicals that are also produced by electrolysis. To make coulometry practical for analytical determinations, we use different techniques depending on whether the method is constant current or controlled potential. In order to maintain a constant current while allowing only one type of reaction to take place, we have to set the conditions so electrolysis of water or other undesirable components does not take place. An electrochemically generated reagent is frequently used as the intermediate reactant. For example, if we want to oxidize Fe2+ to Fe3+ using a constant current, we would have troubles as the Fe2+ concentration approached zero. Water would be electrolyzed. So, we carry out the reaction in the presence of a large excess of Br-. 2 Fe2+ + Br2 à 2 Fe3+ + 2 BrThe Br2 oxidizes the Fe2+ and since there is a large excess of the Br-, the reaction does not change during the reaction. The net effect is like titrating the Fe2+ with electrons and the coulomb count can be used to calculate the amount of Fe2+. HW Problem 24#2: 20 mL of unknown Fe2+ solution is treated with 10 mL of 1 M KBr and titrated with a constant current of 100 mA. What is the Fe2+ concentration if the titration takes 15.77 minutes? Electrochemically generated reagents have been used to quantify reactions where an unstable reactant is used. This allows the reactant to be used immediately, so the instability is usually not a problem. For example Cl2 in water is a good reagent for many organic reactions, but you can't store it since the Cl2 comes out as a gas. HW Problem 24#3: How long must a current of 0.500 amp flow to complete the reaction of 1.00 g of p-cresol as indicated below? Assume Cl- is added in excess.

CH3 Cl CH3

+ 2 Cl 2

HO HO Cl

+ 2 HCl

When a controlled potential is applied rather than a constant current, the current versus time plot shows a decrease in the current, asymptotically approaching zero as the chemicals get used up. By setting the potential at selected voltages, the chemistry taking place can be somewhat controlled. For example, if a solution contains both Cu2+ and Ag+, we can set the voltage at 0.400 V versus a SHE and most of the Ag+ will be reduced to Ag (s) without any of the Cu2+ being reduced. The electronic device for controlling the potential of a working electrode with respect to a fixed reference electrode is called a potentiostat. It automatically adjusts the voltage applied between a counter electrode and the working electrode to keep the potential of the working electrode at a set value from the reference electrode. As the current changes the applied voltage has to change to maintain the voltage of the working electrode. HW Problem 24#4: Assume we have a solution that is 0.100 M in both Cu2+ and Ag+. If a 0.400 voltage is controlled at the cathode, what is the concentration of Ag+ when the current goes to zero? In the real world the above problem would not work as stated because of liquid junction potentials, etc. Most of the time we think of electrochemistry from a thermodynamic point of view and not a kinetics point of view. There are, however, a few interesting phenomena that show up in electrochemistry that are based on unusual kinetics. The reduction of H+ to form H2 (g) is not kinetically favored at some metallic surfaces. It takes an extra voltage (negative) to make such a reduction take place at any appreciable rate. This is called an overvoltage. The hydrogen overvoltage on Hg is far more than the overvoltage on Pt. Zn2+ can be reduced to Zn (Hg) at a Hg electrode from an acidic solution but if you try to reduce Zn2+ on a Pt electrode, H2 gas is evolved and the Zn does not plate. The overvoltage is greatest for gaseous products at a Hg electrode. See page 585, 587, and 623.

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