Voltommetric Behavior was examined in two different environments: unstirred and stirred. It was confirmed that a stirred solution of electrolytes produces a more stable and efficient current in a voltaic cell. This was concluded due to the average current of the unstirred solution being 8.9265*10-6A which is 3.8435*10-6A less than that of the stirred solution that was 1.277*10-5A. The Randles-Sevcik behavior helped to conclude this and then was verified when varying the scan rate while holding everything else constant. Introduction:
Redox reactions contain a great ability to produce currents within electrolytes in a solution. Redox reactions involve the loss or gain of protons and/or electrons from an atom. Consider the redox reaction of iron; (1) Fe3+→e-+Fe2+ E0=0.777 V The loss of an electron, oxidation, from Fe3+ creates an electric potential of 0.777 volts. A “cell” is created when an electric potential is cyclically moving throughout a system. A half-cell reaction at the anode electrode is similar to the oxidation reaction shown above. (2) Pt/0.1 M Fe3+,0.01M Fe2+ This oxidation reaction creates different masses of each state of iron where Fe3+ is favored at the anode. This creates an “open circuit” potential which is the electric potential in the absence of an applied electric current. The open circuit potential can be predicted by the Nernst equation: (3) Ecell=Ecell∅-0.059logFe2+Fe3+ Where the Ecell∅ is the electric potential from reaction 1 of 0.777 and 0.059 is a constant. Thus, the Ecell is 0.836 volts, however, if an electric current of 0.895 volts is applied and the Ecell must remain constant then the concentrations must be altered by a factor of 10. From equation 3 and from the fact that any change in concentration at the electrode surface must result from electrolysis, three predictions can be made. The first is, when a plot of current vs. applied potential is made, the plot will show zero open circuit potential. The second is, the plot will show an increasing value of cathodic current as the graph goes negative of open circuit potential. The third and final prediction is that the anodic current will increase as the graph goes positive of the open circuit potential. These predictions all depend on a current-value which is in units of moles per unit time and depends on the concentration of the species being oxidized or reduced as shown in formula 3. Notice that current is in units of moles per unit but is not measured in moles, this is converted using a Faraday with units of 96,500 coulombs/equivalent with the number of “equivalents” is the number of moles multiplied by the electrons transferred per mole.
Cyclic voltammetry is used to characterize redox-reactions and to quantify redox-active solutes. Cyclic voltammetry follows the above predictions, however, not all the species in the half reaction are present. Along with those two specifications, the potential of a working electrode is scanned between two limiting values. Furthermore, only one oxidation state, Fe3+ or Fe2+, are present in the solution in order to make an initial current of zero and the potential scan initiated at least 0.1 V before the redox reaction occurs. Lastly, a mathematical correction must be noted for the standard potential. For cyclic voltammetry, E0 is replaced by the formal potential of E0’ which is equal to the standard potential times a constant. Procedure:
Three experiments involving voltammetry where performed to better understand voltammetric behavior. The first two experiments performed were of the voltammetry model based on the Nernst equation. A 10 ml solution of 0.1 M KNO3, 0.01M HNO3, 0.001M K3Fe(CN)6, and deionized water was created then the potential energy of the open circuit potential was verified before verifying the anodic and cathodic currents at 50 mV. The second...
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