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 experiment tested conventional cyclic voltammetry with a 10 ml solution of 0.1 M KNO3, 0.01M HNO3, 0.001M Ru(NH3)6Cl3, and deionized water. This was tested in two environments: unstirred and stirred. The unstirred environment had an initial potential of 0.1 V with a switching potential of -0.5 V at a rate of 0.05 V/s. The stirred environment had an initial potential of 1.1 with a high energy potential of 1.1 V and a low energy potential of -0.5 at a rate of 0.1 V/s. Finally, the third experiment was performed to test a redox system for Randles-Sevcik behavior. The same solution for the previous experiment was used while varying the scan rate to confirm or reject the mathematical relation between scan rate and current as shown below in equation 4.
Data & Results:
The first experiment was testing the model based on the Nernst equation which is shown in equation 3. First, the open circuit potential near E0 was verified using deionized water, as seen in Figure 1. A linear regression could not mathematically be created due to software issues, however, quantitatively the verification is nearly linear with a slight parabolic characteristic. The open circuit potential was expected to decay over time due to the reactants becoming fully reduced if the reaction was carried out long enough. The open circuit potential initially was approximately +0.2472V and a standard potential table found by Brown is +0.36V. The difference is 0.1128V which could have been lost due to many errors such as lower initial concentration, differences in apparatus’, electrode not being cleaned, electrode being damaged, etc. Next, the flow of current was verified by running both a forward scan and a revers scan of a 10 ml diluted solution containing 0.01 M HNO3, 0.0442g K4Fe(CN)6, and 0.1M KNO3. The graph and all of the specifications for the verification can be seen in Figure 2. As the figure shows, after four segments were scanned the anodic and cathodic currents flowed which is verified by no breaks or gaps within the graph. After the open circuit potential near E0 and the current flow were verified a conventional cyclic voltammetry experiment, of both unstirred and stirred environments, was conducted with a 10ml solution of 0.01M HNO3, 0.006g Ru(NH3)6Cl3, and 0.1M KNO3. The unstirred experiment graph and specifications are shown in Figure 3. SP shows the Starting Point of the scans and the EP shows the End Point of the scans. As the reaction proceeds forward the reactants undergo reduction. The reverse reaction begins at point A and the species begin to be oxidized. Notice how the forward reaction reaches a current with magnitude of 10*10-6A but the reverse reaction only reaches an approximate magnitude of 6.5*10-6A. This, according to the University of Cambridge, is due to the reactants being electrochemically destroyed and because less products convert back to reactants. The University of Cambridge study on cyclic voltammetry also inferred a relationship between the magnitude of the reverse peak and the rate constant in that the larger the peak the larger value of the rate constant within the reaction. Further experimentation of this relationship would need to be performed to confirm or reject this correlation. The unstirred solution yielded an average E0 of -0.131V which is 0.031V difference from the known standard potential of -0.10V. This difference is relatively small and could have been caused in the same manner as the previous experiment. The stirred solution yielded an average E0 of -0.139V which is only a difference of 0.039V from the known standard potential. This difference is still relatively small. Figure 4 shows the graph and specification for the stirred experiment of cyclic voltammetry. As can be seen, the overall noise has decreased in part due to the order of magnitude of the current in comparison to Fig. 3 but also due to the diffusion of species at the electrode surface. Notice how the forward reaction reaches a current with magnitude of approximately 1.5*10-5Awhereas the reverse reaction reaches a current with magnitude of approximately 1.0*10-5A. The difference between the two currents has more than doubled. This is due to a more efficient diffusion caused by the movement of the species in the bulk solution. The species are able to interact with the electrode more frequently, thereby causing a larger current as mathematically presented by Randles-Sevcik’s equation, (4) ip=(2.95*105)n3/2AD1/2v1/2C Where ip is the predicted current, n is the electron density, A is the area of the electrode, D is the diffusion coefficient, v is the scan rate, and C is the concentration of the bulk solution. As the individual variables increase, the current is predicted to increase as well; if the solution is stirred then D increases and thus the current increases exponentially. This relationship also shows the relationship between stirring and current in that when the solution is stirred the particles are moving and there is an increased interaction with the electrode and thus a higher current. If the rate of stirring is slow relative to the potential scan rate, a steady-state current will not be seen because with a faster scan rate the stirring becomes a negligible variable and the electrode would become saturated with reduced reactants thereby minimizing the area of the electrode in which species may interact. This mathematical relationship of the Randles-Sevcik behavior was then tested with the scan rate, v, being varied over four trials. The solution used in the previous experiment was used to in this experiment. As shown in Figures 5-8, as the scan rate increased the current increased as well. To confirm this relation with more accuracy, more information would need to be obtained such as area of electrode and electron density. Conclusion:
The voltammetric behavior of redox reactions was observed under two conditions: unstirred and stirred. The data collected gave evidence that a stirred solution was more stable and efficient in producing current in comparison to an unstirred solution. This was concluded due to the average current of the unstirred solution being 8.9265*10-6A that is 3.8435*10-6A less than that of the stirred solution that was 1.277*10-5A. This conclusion was also supported by the Randles-Sevcik equation that mathematically shows a relationship between current and diffusion which is increased with a stirred solution. The relationship between current and scan rate within equation 4 was then examined and confirmed that as the scan rate increases the current increases. The data collected supported the Randles-Sevcik behavior of voltammetric cells. References:
Modified lab experiment of Dr. James Cox, Department of Chemistry and Biochemistry, Miami University, Oxford, Ohio Dr. Scaffidi Jon, Department of Chemistry and Biochemistry, Miami University, Oxford, Ohio “Cyclic Voltammetry: The Investigation of Electrolysis Mechanisms”, Department of Chemistry and Biotechnology, University of Cambridge, Cambridge, Massachusetts Brown, Lemay, Bursten, Murphy, “Chemistry: The Central Science” , Upper Saddle River, New Jersey