August 1, 2013
Part A: Voltaic Cells- use a voltmeter to measure the cell potential (E˚cell) between various 1.0 M aqueous reactant solutions, then using balanced half-cell rxns, calculate theoretical cell potential values and compare to experimental. Part B: Concentration Cell: Measure the cell potential of CuSO4 (aq) of two differing Molarities – one concentrated and one dilute- then use the Nernst equation to determine the theoretical cell potential , comparing to experimental. Part C: Electrolytic Cells- use an external source of electricity to drive the chemical reactions of normally nonspontaneous redox reactions and observe what happens at each electrode as electrolysis is carried out then write balanced half reactions and compare to observations.
Part A/B- Aquire a U-tube with flexible tubing at bottom, and close tightly with a clamp to add different cells before mixing. Prepare metal strips of lead, copper, zinc, and iron as electrodes by cleaning them with steel wool. Fill each side of you tube with proper cell and electrode for each test and record cell potential with a voltmeter. Part C – Use an all glass U tube and various cell solutions to be tested with proper electrodes. When ready to begin reaction, connect power supply to generate electricity and take careful observations and read voltages with a voltmeter.
A. Reactant cells
Experimental Cell Potential values
Theoretical Cell Potential values
Pb| Pb(NO3)2 (aq) ||
Cu | Cu(NO3)2 (aq)
Cu | CuSO4 (aq) ||
Fe | Fe (NH4)2(SO4)2 (aq)
Cu | CuSO4 (aq) ||
Zn | ZnSO4 (aq)
B. Concentration Cell
Cu | CuSO4 (aq) || [1.0] con
Cu | CuSO4 (aq) [1x10-6] di
Part A/B- Voltaic Cells- See appendix A for Voltaic NIE and Concentration Cell potential Part C- See appendix B for Electrolytic Cell NIE
In Part A : Voltaic Cells (Galvanic) are spontaneous redox reactions- reactions will spontaneously transfer electrons and release energy. This energy is what was measured with a voltmeter and is the cell potenial, or Ecell The cell potential is the potential difference between the anode and the cathode in a cell known as electromotive force. Electrons spontaneously flow only one way in a redox reaction from higher to lower potential energy. In each of these three cells, the voltages measured were relatively close to the standard state cell potentials. Electrons were spontaneously transferred from the anode side to the cathode side and this is evident in the net ionic equations.
In Part B- Concentration Cell- The CuSO4 compound of different Molarities waswas used in the U-tube and – one concentrated with a [1.0 ] M and the other with [1x10 -6 ] M. According to the standard cell potential value the voltage will always equal zero for a concentration cell and this was true. However, a voltage was measured. The Nernst equation implies that a cell could be created with the same substance at both electrodes. Using this equation I was able to calculate a theoretical cell potential of 0.2 V and compare it to my measured 0.709 V. Though there is a significant difference in the two, this difference could be due to the actual Molarities of the two CuSO4 or the voltmeter that was used.
In part C- Electrolytic cells- a source of power was applied to cells to drive normally nonspontaneous reactions. Unlike Voltaic cells, these reactions need a current of energy to go. For cell 1. Na2SO4- the net ionic equation theoretically implies the reaction should produce oxygen gas and hydrogen gas. This supports my observations because I saw bubbles and pink on the cathode side where OH was produced and bubbles on the anode side where oxygen was produced. For cell 2- KI- The net ionic equation theoretically should produce I2 solid , H2 gas, and 2OH-. According to my observations this is exactly what happened because at the anode side the stainless steel electrode began to bleed a dark red I2 (s) and at the cathode pink pheno indicated the presence of OH and bubbles formed. In cell three- CuBr-at the anode side yellow precipitate formed turning the solution light blue to light green. At the cathode end rust formed on the stainless steel. This rust is from the reduction of the copper 2+ (aq) to copper (s). In cell 4- ogars favorite- the net ionic equation implies cadmium 2+ should form as well as oxygen gas. This is true because at the anode end, bubbles were observed and at the cathode side, when turned up to max voltage, a crystal structure of cadmium formed on the copper electrode. Copper does not react in this reaction. In cell 5- Sulfuric acid- the net ionic equation theoretically implies Hydrogen gas should form and this was true. Vigorous bubbling was observed on the cathode (-) side.