The objective of this experiment was to get acquainted and study galvanic cells, Nernst equation, and electrochemical reactions and determine the experimental value for the Ecell vs. log Q for Cu/Zn cell.
Galvanic cells are the systems which can produce electric current due to reactions occurring in them. In our common understanding these cells are the batteries. In these cells, the processes of electron transfer and corresponding chemical reactions occur due to the difference in tendency of different substances to gain electrons. The galvanic cell typically consists of two half-cells and a junction between them which is often a salt bridge. Each of these half-cells consists of an electrode and electrolyte. The electrolyte typically contains the ions that the electrode can produce. The simplest half-cell is the metal partially submerged in the solution of its salt. Different metals have different tendency to gain or lose electrons. In the galvanic cell consisting of Cu/Cu and Zn/Zn half-cells, copper electrode wants to gain electrons more than zinc, and therefore, copper electrode is gaining electrons, while the zinc electrode gives them off. It is commonly accepted that reactions where electron is gained are called reduction reactions and the ones where electron is lost – oxidation. So, in this cell, copper electrode undergoes reduction and zinc electrode – oxidation. In equation form these reactions can be written as:
Cu2+ + 2e → Cu0
Zn0 → Zn2+ + 2e
The total equation of the process is
Cu2+ + Zn0 → Cu0 + Zn2+
The ions situated in order of ascendance of their tendency to gain electrons are called electrochemical series. Although there is no objective reference point for these series, it is commonly agreed that the voltage of H+ gaining electron under standard conditions is 0 V.
Electrochemical reactions occur if two half-reactions are physically separated from one another (commonly with a salt bridge), in this case the electrons with which the half-cells are exchanging can be drawn outside the system.
Processes occurring in galvanic cells are best described by the Nernst equation. It shows the relationship between electromotive force of the cell or its voltage and the conditions under which the cell operates:
Ecell=Ecell0-2.303 RTnFlogQ
where R is the universal gas constant, T – temperature in Kelvins, n – number electrons transferred during reaction (according to balanced equation), Q – reaction quotient, Ecell0 is the voltage of the cell under normal conditions.
Data
Experimental data for Part I of the experiment is summarized in Tab.1:
The results for the second part are presented in Tab.2:
Analysis
Judging by the results of the measured voltages compared to Pb/Pb half-cell observed in part I, the most negative voltage, and the highest tendency to oxidation is pertinent to Zn/Zn cell. The second in the ascending order of voltage comes the Pb/Pb2+ half-cell and in this series its tendency to oxidation is medium. The highest voltage and the highest tendency to reduction has the Cu/Cu half-cell. Based on these observations, it can be concluded that the half-cells examined in the experiment can be aligned in order of voltage ascendance as Zn/Zn, Pb/Pb, Cu/Cu, with Cu/Cu half-cell having the highest tendency to reduction, Zn/Zn – to oxidation. Pb/Pb acts differently in both cases – with Zn/Zn the reduction reaction occurs in it, and with Cu/Cu half-cell – oxidation. In Pb/Pb pair no current should occur if their concentrations are equal. These measurements were needed to find out that in the same half-cell both oxidation and reduction reactions can occur depending on the voltage of another half-cell to which it is connected.
The data from the second part of the experiment is best presented in the form of graph:
Figure 1. Ecell vs. log Q Graph
In this graph, the y-axis Ecell reflects the voltage of the cells constructed of Cu/Cu half- cell and Zn/Zn half-cells of different Zn2+ ion concentrations. The x-axis, log Q, represents the logarithm of the ratio of [Zn2+] to [Cu2+] in the corresponding half-cells. The lower is [Cu2+], the lower is the voltage of the cell. In all of those cells the reaction is the same:
Cu2+ + Zn0 → Cu0 + Zn2+
Oxidation reaction occurs in the Zn/Zn half-cell and reduction reaction occurs in Cu/Cu half-cell. The experimental value for the slope of the graph is -0.037 as can be seen from the approximation equation. In essence, the slope of this graph represents the multiplier standing before log Q in Nernst equation. In the basic form it is:
Ecell=Ecell0-0.0591nlogQ
For this particular cell n = 2 (the reaction involves two electrons), and so the Nernst equation will look like:
Ecell=Ecell0-0.05912logQ=Ecell0-0.02955logQ
So, the theoretical value for the slope is -0.2955 and the experimental value is different from the theoretical.
Conclusion
The objectives of this experiment were men and it can be concluded successful. The galvanic cells, Nernst equations, and electrochemical reactions were explored, and the experimental value for the slope of the Ecell vs. log Q graph was found. There was a discrepancy between theoretical (-0.02955) and experimental (-0.037) values of the slope and the percent error comprised
%error=theor.-exp.theor.100%=-0.02955-(-0.037)-0.02955100%=25.21%
The error is not huge but is significant. This error may have occurred due to a number of processes. Among the most possible sources are contamination of the electrodes with outside solutions and the decrease of concentration of ions in the cells due to small volume of the wells and ion depletion in electrochemical reaction (especially Cu2+).
