Electrochemistry: Electrode Potential

Abstract

Chemical reactions can be used to produce electricity and electricity can be used to cause chemical reactions through oxidation-reduction reactions. The first part of the experiment measures the standard electrode potentials of five various half cells against the Cu2+(1M)|Cu half cell. The last three half cells are prepared through electrolysis. The electrode potentials of all reactions are positive which means that they are spontaneous. There is a significant percent difference from the theoretical and the experimental standard electrode potentials. The sources of error are incorrect solution preparation and contaminated materials. The second part of the experiment uses standard electrode potentials to predict whether or not a reaction will occur and eventually test that prediction. All the standard electrode potentials are positive which means that they are spontaneous, as directly observed. There is also a percent difference from the theoretical and experimental cell potentials. The sources of error are also incorrect solution preparations, contaminated glasswares and materials, and the resistance of the electrodes.

Keywords: cell potential, half-cell, electrode potential, oxidation-reduction reactions, standard electrode potential

Introduction

Chemical reactions can be used to produce electricity and electricity can be used to cause chemical reactions. The practical applications of electrochemistry are countless, ranging from batteries and fuel cells as electric power sources, to the manufacture of key chemicals, to the refining of metals, and to the methods of controlling corrosion.

The objective of this experiment is to measure the standard electrode potentials of five various half cells against the Cu2+(1M)|Cu half cell and to use the standard electrode potentials to predict whether or not a reaction will occur. The standard electrode potential, EÐ'Ñ"cell, is the electric potential that develops on an electrode when the oxidized and reduced forms of some substance are in their standard states. When used in electrochemical studies, a strip of metal is called the electrode. An electrode immersed in a solution containing ions of the same metal is called a half-cell. A salt bridge is used to join two half-cells in an electrochemical cell. It salt bridge permits the flow of ions between two half-cells.

The first part of experiment measures the standard electrode potentials of five various half-cells against the Cu2+(1M)|Cu half cell. The last three half cells are prepared through electrolysis. Electrolysis is the decomposition of a substance, either in the molten state or in an electrolyte solution, by means of an electric current.

The second part of the experiment uses the standard electrode potential to predict whether or not a reaction will occur. A reaction is said to occur if the standard electrode potential is greater than 0. That reaction is called a spontaneous reaction. If the standard electrode potential is less than 0, the reaction is said not to occur. That reaction is a called non-spontaneous reaction.

Experimental Detail

The Cu2+ (0.1 M)|Cu half cell is set-up by immersing a copper electrode in 0.10 M CuSO4. The 0.10 M CuSO4 was prepared by diluting 10 mL of 1.0 M CuSO4 in a 100 mL volumetric flask. The Zn2+ (0.1 M)|Zn half cell is set-up by immersing a zinc electrode in 0.10 M ZnSO4. The 0.10 M ZnSO4 is prepared by the laboratory. The Fe2+ (0.5 M), Fe3+ (0.5 M)|C half-cell is prepared by immersing a graphite electrode in a solution prepared by mixing equal volumes of 0.1 M FeSO4 and 0.1 M FeCl3. The graphite electrode is obtained from a pencil.

The Cu2+|Cu half cell is connected to the other two half-cells in separate set-ups using a salt bridge between the two half cells and a voltmeter to measure the emf. The salt bridge is prepared by adding a salt solution in a fabricated glass U-tube. The two ends of the U-tube are covered with cotton balls. The voltmeter reading is recorded.

The Cl (1M), Cl2|C, Br (1M), Br2|C and the I (1M), I2|C are prepared through electrolysis. The halogen X2 is generated by electrolyzing for about a minute a 1 M solution of potassium halide KX between graphite electrodes using a current from two 1.5 V dry cells connected in series.

The X-, X2|C half cells is connected to the Cu2+|Cu half-cell. The meter is in place before generating the halogen in order not to disturb the X, X2|C half-cell after its preparation. The dry cells are disconnected before connecting the two half-cells.

The second part of the experiment is the applications of the electrochemical cells. Any voltaic cell in which the net cell reaction involves only a change in the concentration of some species is called a concentration cell. A concentration cell consists of two half-cells with identical electrodes but different ion concentrations. The concentration cell is set up by this cell notation:

Cu(s)|Cu2+(aq)(0.01 M)||Cu2+(aq)(0.1 M)|Cu(s).

The second set-up is a redox reaction involving complexes. The electrochemical cell is set up by this cell notation:

Zn(s)|Zn2+(aq)(0.1 M)||[Cu(NH3)4(0.33M), NH3(aq)(1.0 M)|Cu(s).

The [Cu(NH3)4] solution is prepared by mixing 10 mL of 0.1 M Cu(SO4)2 with 20.0 mL 1.0 M NH3.

The final set-up is reaction of partially soluble solid in the electrochemical cell. The electrochemical cell is set up by this cell notation:

Zn(s)|Zn2+(aq)(0.1M)||OH-(0.10M)|Cu(OH)2(s),Cu(s)

The Cu(OH)2 is prepared by mixing 10 mL of 0.10 M Cu(SO4)2 with 30.0 mL of 1.0 M NaOH.

Results and Discussion

As shown in Figure 1 is a simple voltaic cell. A voltaic cell is an electrochemical cell in which a spontaneous chemical reaction produces electricity. The anode is the electrode at which the oxidation occurs while the cathode is the electrode at which reduction occurs. The electrons are joined by a metal wire to permit the flow of electrons from the anode to the cathode. The flow of electric current between the solutions is in the form of migration of ions. The ends of the salt bridge are plugged with a porous material that allows ions to migrate but prevents the bulk flow of fluid.

Figure 1. A simple voltaic (galvanic) cell

As shown in Table 1 are