Galvanic Cells - Practical Report PDF

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Galvanic Cells Introduction Electrochemistry is the study of chemical reactions which involve the transfer of electrons. It also includes the study of relationships between chemical reactions and electricity. In able to perform a beneficial electrical work, galvanic cells bind the electrical energy available from the electron transfer in a redox reaction(SparkNotes Editors, n.d.). This is by splitting the oxidation and reduction reactions, the transfer of electrons takes place through an external route instead of straight between reactants. Two half cells are linked by a wire, so that the electrons are able to flow through that wire. In order for the circuit to be completed, a salt bridge is needed. Galvanic cells can also be referred to as voltaic cells which general uses spontaneous redox reactions to generate electricity. Salt bridges complete a circuit and are usually made up of salt solutions like NaNO 3 or KNO3. These don’t interfere with the redox reactions and this allows the flow of electrons by keeping half-cells neutral which avoids charge build up. Anions flow towards the anode where negative ions go towards the negative terminal to counteract electrons that are lost in oxidation. The cathode is where cations flow and it is where positive ions go towards the positive terminal to neutralise electrons being gained in reduction. Any galvanic cell measuring the cell potential, Ecell, in volts with the help of a voltmeter. It isn’t possible to directly measure the potential of each individual half cell however chemists have formulated a method to measure the ability of a chemical species to reduce another by accumulating tables of standard reduction potentials(SparkNotes Editors, n.d.). The o in Eo, shows that the reaction is at a standard state subjectively allocating a charge of exactly zero to the potential of the standard electrode to measure the Eo of any half-reaction. The dissimilarity between the potentials of two given electrodes that are dipped into the equivalent or distinctive solution would be able to be measured (Lower 2014). An oxidation-reduction half cell is made by every electrodesolution pair and the sum of the two half-cell potentials are then measured(Lower 2014). This is called a galvanic cell. An example is shown on the right of the usual setup of a galvanic cell.

Materials & Method o o o o o

100 mL of each metal solution (0.1M ZnSO4, 0.1M CuSO4 and 0.1M Pb(NO3)2) Metal electrodes (Zn, Cu and Pb) Voltmeter with alligator clips Three 150mL beaker Filter paper and 1.0M KNO3 for the salt bridge

Each 150mL beaker was used and filled with 100 mL of each metal solution (0.1M Zinc Sulfate, 0.1M Copper(II) Sulfate and 0.1M Lead Nitrate). Each of the metal electrodes were scrubbed with sandpaper removing any residue or impurities they might have had. The beakers were placed side by side with either of the combinations below. Combination 1: Combination 2: Combination 3:

Zn2+/Zn(s) and Cu2+/Cu(s) Zn2+/Zn(s) and Pb2+/Pb(s) Cu2+/Cu(s) and Pb2+/Pb(s)

The salt bridge was prepared by placing a strip of the filter paper in the 1.0M KNO3 solution to be absorbed. The prepared salt bridge was removed from the solution it was soaked in and draped, connecting the beakers without them physically touching. Both ends of the strips was then immersed in a beaker filled with the chosen solution creating the salt bridge. Each corresponding metal electrode was connected to a voltmeter via the alligator clips connecting the (-) black lead to the anode and the (+) red lead to the cathode. Each electrode was placed into its matching beaker without it touching the salt bridge avoiding a short circuit of the cell if they were to touch. The voltmeter was then switched on, putting it at 20 volts. The potential difference of the cell was shown and recorded after a stable reading on the voltmeter. The cell was left to settle for around five minutes to see if there were any qualitative observations and any visual changes were recorded. The procedure was repeated for the remaining combinations by connecting new pairs of electrodes together. Diagrams were constructed showing the galvanic cell of the combinations done. Below is an example of a galvanic cell and it usually includes:      

The half-cell that was the cathode and the half-cell that is the anode The metal that was the anode and the metal that was the cathode The solutions in each beaker The direction that the electrons flow through the wire The directions ions would migrate in the salt bridge when the cell is operating The relevant balanced half-equations underneath each cell.

After all combinations were finished and results already recorded, all equipment were cleaned and returned where found. The solutions that were disposed of into the ‘HEAVY METAL’ waste container. Gloves were worn as protection to the solutions and electrodes and it might have been harmful to the skin. Salt bridges were used as it completes the circuit.

Results

The different galvanic cell combinations are shown in figures 1-3. Table 1 shows the theoretical potential (V) and the measured potential (V). The cells are spontaneous as looking at the results of the calculated voltage, the galvanic cell potential are all greater than zero. This means that these cells are able to be conductors of electricity. Below are the different combinations used in the experimentCombination 1: Combination 2: Combination 3:

Galvanic cell combination

Zn2+/Zn(s) and Cu2+/Cu(s) Zn2+/Zn(s) and Pb2+/Pb(s) Cu2+/Cu(s) and Pb2+/Pb(s)

Anode half reaction (V) (±0.005)

Zinc and Copper -0.763 Zinc and Lead -0.763 Copper and Lead -0.13

Cathode half reaction (V) (±0.005)

Theoretical potential (V) (±0.005)

Measured potential (V) (±0.005)

Visual change

0.337 -0.126 0.34

1.1 0.637 0.47

1.06 0.60 0.47

No Change No Change No Change

Table 1 Theoretical and calculated voltages

Referring to Table 1 , there is minimal differences between the theoretical potential to the measured potential. However to get the best results possible, thorough cleaning of equipment before use is essential to ensure there isn’t any unwanted residue left on the equipment. If the cells where left to sit for longer, there may have been a chance of possible visual change but there was no sign of changes for the time it was left.

A redox reaction is spontaneous if the standard electrode potential for the redox reaction, Eo(redox reaction), is positive (AUS-e-TUTE). Eo(redox reaction) = Eo(reduction reaction) + Eo(oxidation reaction) Eo(redox reaction) > 0 that is, Eo(redox reaction) is positive

Different cell combinations

Figure 1 Galvanic cell of Zinc and Copper

Overall redox reaction in galvanic cell Cu2+ + Zn(s)  Zn2+ + Cu(s) In Figure 1, shows the reaction between zinc and copper in the galvanic cell. The anode in this case was Zinc and it was being oxidised. This means that zinc ions were being produced due to loss of electrons. The cathode was Copper and it was being reduced which means there was gaining of electrons in order to produce copper atoms.

Figure 2 Galvanic cell of Zinc and Lead

Overall redox reaction in galvanic cell Zn(s) + Pb2+(aq)  Zn2+(aq) + Pb(s) In Figure 2, shows the reaction between zinc and lead. The anode in this case was Zinc and it was being oxidised. Zinc ions were being produced due to loss of electrons. The cathode was lead and it as being reduced, meaning there was gaining of electrons in able to produced lead atoms.

Figure 3 Galvanic cell of Copper and Lead

Overall redox reaction in galvanic cell Pb2+(aq) + Cu  Cu2+ + Pb(s) In Figure 2, shows the reaction between lead and copper. The anode in this case was lead and it was being oxidised. lead ions were being produced due to loss of electrons. The cathode was copper and it as being reduced, meaning there was gaining of electrons in able to produced copper atoms. Displacement reactions involve a transfer of electrons, galvanic cells use cell notation and use of tables of standard reduction potentials. Later on predictions were made if spontaneous reactions would occur with a given species and in this case all pairs of the half cells simultaneously reacted. Potential difference was observable but there was no significant current that can flow as well as no significant chemical change. The flow of electrons creates the electrical current. The conventional way the electrochemical cell is the oxidation half reaction written on the left and the reduction on the right-hand side(Lower 2014). For example :

Zn(s) + Cu2+ → Zn2+ + Cu(s) we write Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

Conclusion The galvanic cell reactions were spontaneous and are able to be conductors of electricity. Two half-cells were set up in different beakers, being connected by the salt bridge. The anode was negative and the cathode is the more positive electrode. The reaction that occurs at the anode is oxidation and at the cathode reduction occurs. Electrons being supplied by the substances getting oxidised where they move from the anode to the cathode in the circuit. The oxidation and reduction half-reactions are joined by a wire, so that the electrons flow through that wire. A current is sent through the circuit which explains why they could be used for a number of electrical purposes. When two half-cells are combined by a salt bridge that allows ions to pass between two sides maintain electroneutrality(Lower 2014). These cells are essential because they are the foundation for the batteries that fuel modern society. Two or more cells that are connected together is what forms a battery. That is the principles of cells in order to make electrical batteries.

References Lower, S. 2014, LibreTexts Chemistry: Galvanic cells and electrodes, California, viewed 23 May 2017,

SparkNotes Editors. (n.d.), Galvanic Cells, Spark notes, viewed 24 May 2017,

UTS Insearch Lecture 5 A&B Slides Spontaneous Redox Reactions and non-spontaneous redox reactions tutorial, AUS-e-TUTE, viewed 21 May 2017, < http://www.ausetute.com.au/redoxspon.html> Lower, S. 2014, 24.2: Galvanic cells and electrodes, LibreTexts Chemistry, California, viewed 23 May 2017,...