Electrolysis Reaction Lab

Chemistry 1225 Lab Write-Up #13

Abstract

Electrolysis is a method of using a direct electric current to drive an otherwise non-spontaneous chemical reaction. During electrolysis, hydrogen atoms (H2) are reduced at the cathode whereas the oxygen atoms (O2) are oxidized at the anode. There were three different solutions used in the experiment in order to have a better understanding of electrolysis reactions. From each solution, reaction equations were produced. The reactions could be observed from the indicators used. The reactions showed that hydrogen ions were produced at the anode, making it acidic and hydroxide ions were produced at the cathode, making it basic. In all parts of the experiment, the reaction that took place at the anode was 2H2O(l)  O2(g) +
4H+ + 4e- and the reaction at the cathode was 4e- + 4H2O(l)  2H2(g) + 4OH-.

Safety

The universal indicator is flammable and must be handled with caution – be sure to keep away from any open flames. Wear splash goggles when handling any chemicals. Before leaving the lab, be sure to thoroughly wash your hands.

Introduction

Electrolysis is a method that is used to separate compounds by attracting ions in aqueous solutions to the opposite electrode. This means the electrode that is attached to the negative pole of the battery (supplies electrons) is called the cathode. The electrode attracted to the positive pole of the battery (accepts electrons) is called the anode. Electrodes that contain graphite are widely used – graphite is typically used because they are inert.

In order to successfully ionize, the aqueous electrolyte containing hydrogen and hydroxide ions, as well as the solute, they must compete with each other at the electrodes. The hydrogen atoms (H2) are reduced at the cathode whereas the oxygen atoms (O2) are oxidized at the anode. When there is an electric current running through an aqueous solution containing an electrolyte, water molecules separate into hydrogen and oxygen. Therefore, the reaction occurs as two distinct half reactions. As a result, each half reaction is accompanied by hydroxide (OH-) and hydronium (H+) ions. The OH- ions will assemble at the cathode and cause the pH to increase. At the anode, H+ ions will gather and therefore cause the pH to decrease.

A universal indicator is a solution that undergoes several color changes over a wide range of pH values. The color is used to indicate pH directly. Universal indicators are usually mixtures of several indicators. The scale for the universal indicator is provided below.

Procedure

Part A: Universal Indicator and Sodium Sulfate

Begin by obtaining a petri dish. Separate the petri dish so you now have two halves in front of you. Pour approximately 5-10 mL of universal indicator to roughly 20-30 mL of sodium sulfate solution. Mix both of them thoroughly. Pour just enough of this solution into the petri dish in order to cover the bottom half of the dish.
Obtain pencil lead. Break the piece of lead in half. Attach one half of lead to one side of the petri dish with an alligator clip and attach the other half of lead to the opposite side with the other alligator clip. Without having the alligator tips submerged in the solution be sure both halves of the lead are submerged. Attach the 9V battery to the snaps that are connected to the battery clip. Allow the experiment to run for about 5 minutes. Note any observations.

Part B: Potassium Iodide and Phenolphthalein

Repeat the procedure from part A; however, instead of adding universal indicator and sodium sulfate, use approximately 20 grams of potassium iodide and roughly 30 drops of phenolphthalein. Allow the experiment to run for about 5 minutes. Note any observations.

Part C: Potassium Chloride and Bromothymol Blue

Repeat the procedure from part A; however, instead of adding universal indicator and sodium sulfate, use approximately 20 grams of potassium chloride and roughly 30 drops of bromothymol blue. Allow the experiment to run for about 5 minutes. Note any observations.

Data

Part A

Color pH Reaction
Anode
Bright Red / Orange
~4.5
2H2O(l)  O2(g) + 4H+ + 4e- (Oxidation)
Cathode
Dark Purple
~10
4e- + 4H2O(l)  2H2(g) + 4OH- (Reduction)

Part B

Color pH Reaction
Anode
Yellowish Clear
~3.5
2H2O(l)  O2(g) + 4H+ + 4e- (Oxidation)
Cathode
Bright Pink
~11.5
4e- + 4H2O(l)  2H2(g) + 4OH- (Reduction)

Part C

Color pH Reaction
Anode
Bright Orange
~3
2H2O(l)  O2(g) + 4H+ + 4e- (Oxidation)
Cathode
Deep Blue
~11
4e- + 4H2O(l)  2H2(g) + 4OH- (Reduction)

Discussion

In part A, a current passed through an aqueous solution of sodium sulfate.
When the current passes through, the water was oxidized at the anode and reduced at the cathode. The equations that were produced show that hydrogen ions are produced at the anode, making it acidic and hydroxide ions were produced at the cathode, making it basic. Sodium sulfate was used because pure water doesn't contain a high enough concentration of ions to produce a current; therefore, the sodium sulfate was acting as the salt bridge in the reaction. This processes was demonstrated in parts B and C as well; however, instead of the sodium sulfate acting as the salt bridge, potassium iodide was substituted in part B and potassium chloride in part C.

In each part of the procedure a different indicator was used – part A was accompanied by a universal indicator, part B was joined with phenolphthalein, and part C was combined with bromothymol blue. The indicators allow one to follow what was happening during the reaction. The indicator would change the color based on how acidic or basic the solution was. From the pH values obtained, one can then decipher the reactions taking place at the cathode and anode.
Conclusion

In part A (sodium sulfate and universal indictor), a bright red color was observed at the anode and a dark purple at the cathode. From the colors observed, we were able to approximate the pH – the pH at the anode was ~4.5 and the pH at the cathode was ~10. In part B (potassium iodide and phenolphthalein), a yellow color was observed at the anode and bright pink at the cathode. From the colors observed, we were able to approximate the pH – the pH at the anode was ~6 and the pH at the cathode was ~10. In part C (potassium chloride and bromothymol blue), a bright orange color was observed at the anode and deep blue at the cathode. From the colors observed, we were able to approximate the pH – the pH at the anode was ~5 and the pH at the cathode was ~9. In all parts of the experiment, the reaction that took place at the anode was 2H2O(l)  O2(g) + 4H+ + 4e- and the reaction at the cathode was 4e- + 4H2O(l)  2H2(g) + 4OH-. These results demonstrate the concept of electrolysis reactions.