Today, STEM All Stars conducted a fun experiment and investigated the electrolysis of water and potassium iodide! Check out our experimental design and analysis of our results below!

Purpose: The purpose of this lab was to measure the reduction potentials of unknown metals relative to copper.

Materials:

  • Multimeter Metal samples M1, M2, M3, M4, and M5
  • Scissors 1 M NaNO3
  • Overhead transparency 1 M solutions of M1, M2, M3, M4, and M5
  • Forceps Sandpaper
  • Filter paper

Procedure:

The multimeter was first set up by connecting the black (−) lead to the ‘com’ port and the red lead (+) to the ‘VΩmA’ port. The dial was turned to 2000m in the DCV region of the multimeter.

A piece of filter paper was obtained, and five small circles were drawn in a pentagonal arrangement with connecting lines and labeled from M1 to M5. Wedges were cut between the circles such that each circle was at the end of one ‘arm’ of the paper. The filter paper was placed on top of the overhead transparency.

Five solutions of metals M1 to M5 were obtained. One piece of each metal was obtained as well, and the surfaces were sanded. Two drops of each solution and each piece of metal were placed on the corresponding circle. Drops of 1 M NaNO3 solution were added as a salt bridge between the metals such that a continuous trail of NaNO3 connected each metal. The NaNO3 trail was periodically refreshed throughout the course of the experiment.

M1, the metal which was copper (as identified by a blue solution and red-brown metal), was used as the reference metal. The potential of four cells was determined by using the leads of the multimeter to connect M1 to each M2, M3, M4, and M5. The voltage reading was recorded; if the reading was negative, the terminals were reversed to yield a positive reading and the reversal was recorded.

The five metals were arranged in order in Data Table 2 from highest to lowest reduction potential. M1, the reference metal, was given an an arbitrary value of 0.00 V. Metals that were connected to the negative terminal were placed above M1, and metals that were connected to the positive terminal were placed below M1. The predicted potential of each of the remaining cell combinations was calculated using the values in Data 2 and recorded in Data Table 3. The experiment was completed by recording the voltage readings of the remaining cell combinations (M2-M3, M2-M4, M2-M5, M3-M4, M3-M5, and M4-M5) using the previously detailed procedure.

At the conclusion of the experiment, each piece of metal was removed from the filter paper using forceps, rinsed, dried, and returned to its place. The filter paper was discarded and the transparency was rinsed.

Data Analysis:

Formulae Used:

Cell Reduction Potential Formula: Eºcell = Eºcathode − Eºanode

Percent Error Formula: % error = actual – theoreticaltheoretical 100

Example Calculations:

M2/M3     Eºcell = 1.00 V − (-0.39 V) = 1.39 V

% error = 1.40V – 1.39V1.39V 100 = 0.01V1.39V100 = 0.7%

The large % error values and different ranking of reduction potentials in Lexa and Sajni’s data are likely due to a significant under-recording of the voltage for M5. See Experimental Sources of Error (question three).

Discussion of Theory:

The key concepts in these labs were redox reactions, half reactions, and the construction of a galvanic/electrolytic cell.  In a redox reaction, two half-reactions occur; one reactant undergoes oxidation and another reactant undergoes reduction. A measure of the tendency for reduction to occur is reduction potential, E, measured in units of volts. At standard conditions, 25 °C and concentrations of 1.0 M for the aqueous ions, the measured voltage of the reduction half reaction is defined as the standard reduction potential. In this experiment, the half cell is a small piece of metal placed into three drops of the corresponding cation solution on a piece of filter paper.

A galvanic cell uses a spontaneous oxidation-reduction reaction to produce electrical energy. A galvanic cell consists of an anode and a cathode. Oxidation occurs at the site of the anode while reduction occurs at the site of the cathode.  The solution of sodium nitrate (NaNO3) placed on the filter paper linking the half-cells is called a salt bridge. Without the salt bridge, there would be charge buildup and an inconsistency of charges on both sides. Using the multimeter, the positive end of the Voltage Sensor makes contact with one metal and the negative end of the Voltage Sensor makes contact with another metal. Thus, the differences of the voltages are able to be seen between the metals and they could be ranked on a number line to make a standard reduction potentials chart with the strongest oxidizing agent to the strongest reducing agent.

An electrolytic cell decomposes a substance using an electric current generated from an external power source. The power source pulls electrons from the positive pole (anode) and pushes electrons into the negative pole (cathode); an electrolytic reaction is nonspontaneous because the power source reverses the redox reaction (cathode is negative rather than positive), though reduction still occurs at the cathode. In this lab, it is important to note that for the electrolysis of KI solution, two oxidation reactions occurred at the anode and two reduction reactions occurred at the cathode because the KI was in aqueous solution (read: water was present). At the anode, iodine ions were oxidized in the half-reaction 2II2+2e and water was oxidized in the half-reaction 2H2O(l) → O2(g) + 4H+(aq) + 4e. Because iodine ions have a higher oxidation potential (lower reduction potential), however, only iodine was actually oxidized at the anode. Similarly, at the cathode, potassium was reduced in the half-reaction K++eK(s)and water was reduced in the half-reaction 2H2O+2eH2(g)+2OH. However, because water has a higher reduction potential, only water was actually reduced at the cathode. Thus, hydrogen gas formation rather than solid potassium formation occurred.

 

Conclusions:

These labs demonstrated the construction of electrochemical cells: a galvanic cell, constructed with filter paper, and an electrolytic cell, constructed using a petri dish. In the galvanic cell, a salt solution (sodium nitrate) was used to connect the electrodes, while electrodes were isolated by creating strips, or ‘arms’, on the filter paper. The solutions only needed to the electrodes, which rested on top. In the petri dish electrolytic cell, two graphite electrodes with opposite charge were used to perform the electrolysis of water and potassium iodide. An indicator was added to the petri dish solution in order to identify the cathode and anode; for example, the oxidation half-reaction of the electrolysis of water produces hydrogen ions, while the reduction half-reaction produces hydroxide ions. The two experiments successfully showed the concepts behind galvanic and electrolytic cells.

Future investigation could include constructing electrochemical cells using beakers rather than filter paper or a petri dish. Beaker-constructed electrochemical cells could better show how the cathode and anode change in size over the course of the reaction, as well as better demonstrate the components that we have learned about, including the salt bridge.

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