Investigating the Nernst Equation

HOW DOES THE EMF OF A CHANGE WITH THE CHANGE IN CONCENTRATION OF COPPER SULPHATE ELECTROLYTE? RAAHISH KALARIA Contents

1 Introduction ...... 2 2 Research Question ...... 5 3 Hypothesis ...... 5 4 Variables...... 5 4.1 Independent Variable ...... 5 4.2 Dependent Variable ...... 6 4.3 Controlled Variables ...... 6 5 Apparatus ...... 8 6 Pre-Lab Preparation ...... 9 6.1 Preparing the Salt Bridge ...... 9 6.2 Zinc sulphate (1 mol dm-3, 250 cm3) ...... 10 6.3 Copper sulphate (1 mol dm3, 250 cm3) ...... 11 6.4 Copper sulphate (0.1 mol dm-3, 250 cm3) ...... 12 6.5 Preparing the Electrodes ...... 14 7 Procedure ...... 15 7.1 Setting up the Apparatus ...... 15 7.2 Investigation Procedure ...... 16 7.3 Safety Considerations ...... 17 8 Data Collection and Processing ...... 18 8.1 Qualitative Data ...... 18 8.2 Raw Data Collection ...... 18 8.3 Data Processing ...... 21 9 Conclusion and Evaluation ...... 29 9.1 Conclusion ...... 29 9.2 Evaluation ...... 30 10 Further Investigation...... 31 11 Bibliography ...... 32

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1 Introduction

This investigation aims to study chemical reactions occurring in the Galvanic cell. A Galvanic cell typically consists of Zinc and Copper electrodes, in different containers (half cells), immersed in an aqueous solution containing the salt of the corresponding metal1. The two containers are connected with a salt bridge, which acts like a pathway for the cations of one solution and the anions of the other solution, completing the circuit. The cell, due to the redox reaction taking place, generates a difference in between the electrodes, thus generating an (EMF). This investigation studies the effect of the concentration of one electrolyte on the overall EMF produced. In theory, the Nernst Equation is used to calculate the potential of a cell. This investigation also aims to verify this equation.

Figure 1.1 – Diagram of a Galvanic Cell

Photo taken from: http://upload.wikimedia.org/wikipedia/commons/thumb/a/a5/Galvanic_cell_labeled.svg/2000px- Galvanic_cell_labeled.svg.png

Redox reactions are reactions in which there in a transfer of from one reactant to another2. In these reactions, two half-reactions occur, in which one reactant loses electrons (oxidizes), and the other reactant gains electrons (reduces). In this experiment, the two metals used are Zinc, which is the anode, and Copper, which is the cathode. The respective electrolytes used were Zinc Sulphate and Copper Sulphate, of which the concentration of Zinc sulphate was kept constant and the concentration of Copper sulphate was varied. The salt bridge consisted of Potassium Nitrate in aqueous solution which is acts as a good conductor of the in the cell.

1http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_2%3A_Galvanic_cells_ and_Electrodes 2http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Redox_Chemistry/Oxidation- Reduction_Reactions

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Thus, the reaction occurring at the anode is:

푍푛(푠) → 푍푛2+(푎푞) + 2푒− The reaction occurring at the cathode is:

퐶푢2+(푎푞) + 2푒− → 퐶푢(푠) Due to the positive charge on the Zinc half-cell, the negative sulphate ions are attracted to that side of the cell, and move through the salt bridge, while the positive Zinc ions are attracted to the copper side of the cell, and also move through the salt bridge to the copper half-cell.

For any electrode, at standard conditions of 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 , E°.

The reduction potentials of Zinc and Copper are given in the table below3:

Table 1.1 – Standard Reduction potentials (E0) of Zinc and Copper. Electrode Standard Reduction potential (V) Zinc -0.76 Copper 0.34

The greater the reduction potential, the greater is the tendency for the reduction to occur. So Cu2+ has a greater tendency to be reduced than Zn2+.

The cell potential, Ecell, which is a measure of the EMF that the battery can provide, is calculated from the standard half-cell reduction potentials4:

0 0 0 퐸푐푒푙푙 = 퐸푐푎푡ℎ표푑푒 − 퐸푎푛표푑푒

0 Thus, for the reaction taking place in the cell, the E cell is:

0 퐸푐푒푙푙 = (0.34) − (−0.76) 0 퐸푐푒푙푙 = 1.10 푉 As this is positive, we can deduce that the reaction in the cell is spontaneous, from the Gibb’s 0 5 0 equation, ΔG = -nFE cell , as when the E cell is positive, the value of ΔG is negative, which means the reaction is spontaneous. Therefore, there would be a current produced in the cell.

The Nernst Equation is used to calculate the potential of a cell. The equation is as follows6: 푅푇 퐸 = 퐸0 − ln (푄) 푐푒푙푙 푐푒푙푙 푛퐹

3 https://www.chem.umn.edu/services/lecturedemo/info/Cu-Zncell.html 4 Pearson Baccalaureate: Higher Level Chemistry for the IB Diploma (Print) 5http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_and_Thermodynamics 6 https://www.chem.tamu.edu/class/majors/tutorialnotefiles/nernst.htm

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Where:

0 E cell is the Standard reduction potential of the cell; R is the Universal , which is 8.314 J K-1 mol-1; 7 T is the temperature of the cell, in ; n is the number of electrons transferred in the redox reaction, which is 2 in this case; F is the , which is 96485.3 J V-1 mol-1; 8 Q is the ratio between the concentration of the anode metal ions and the cathode metal ions in the respective electrolytes. [푍푛2+] 푄 = [퐶푢2+] Using the Nernst Equation, a Nernst plot can be obtained on a graph, which would express the relation between the concentration of one of the electrolyte and the Ecell. To do this, we need to get the Nernst Equation in y = mx + c form.

The equation can be written as:

푅푇 [푍푛2+] 퐸 = 퐸0 − ln ( ) 푐푒푙푙 푐푒푙푙 푛퐹 [퐶푢2+] 푅푇 퐸 = 퐸0 − (ln([푍푛2+]) − ln ([퐶푢2+])) 푐푒푙푙 푐푒푙푙 푛퐹 푅푇 푅푇 퐸 = 퐸0 − ln([푍푛2+]) + ln ([퐶푢2+]) 푐푒푙푙 푐푒푙푙 푛퐹 푛퐹 Since we are keeping the concentration of Zinc sulphate electrolyte constant, and varying the 푅푇 concentration of Copper sulphate, the 퐸0 − ln([푍푛2+]) part becomes constant. Thus becomes 푐푒푙푙 푛퐹 푅푇 the ‘y-intercept’, or ‘c’ in the equation y = mx + c. Our ‘x’ is ln ([퐶푢2+]), and ‘m’ or the slope is , a 푛퐹 constant.

푦 = 푚푥 + 푐 푅푇 푅푇 퐸 = ln([퐶푢2+]) + (퐸0 − ln([푍푛2+])) 푐푒푙푙 푛퐹 푐푒푙푙 푛퐹 Thus, we can deduce that the Nernst plot is a straight line, which increases with the increase in the value of ‘x’, as the slope is positive. The value of ln([Cu2+]) increases if the value of [Cu2+] increases.

7 http://hyperphysics.phy-astr.gsu.edu/hbase/kinetic/idegas.html 8 http://chemistry.about.com/od/chemistryglossary/g/Faraday-Constant-Definition.htm

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2 Research Question

How does the EMF of a Galvanic cell change with the change in concentration of Copper sulphate electrolyte?

3 Hypothesis

According to the Nernst Equation, an increase in the concentration of Copper sulphate will lead to greater EMF generation. This is because when the concentration increases, the value of Q would decrease, consequently decreasing the value of ln(Q), and thus, decreasing the subtracted value 0 from E cell, and overall, increasing the value of Ecell. However, the change in value of concentration would have to be very large, because only then would the change in the value of ln(Q) be significant enough to cause a measurable change to the Ecell.

4 Variables

4.1 Independent Variable Variable Unit How it was measured Concentration of mol dm-3 The concentration of Copper sulphate was changed by Copper sulphate making individual solutions of the required concentration. To get a significant impact on the Ecell, the concentrations needed to have a huge difference between each other. Hence, the concentrations were varied by the 10-1th power, so that it would make some significant change to ln(Q) in the Nernst Equation. The concentrations used are shown below: Table 4.1 – Concentrations of Copper sulphate used Reading Set no. Concentration (mol dm-3) 1 1 2 0.1 3 0.01 4 0.001 5 0.0001 The solutions were prepared using serial dilution method, and thus, the uncertainty for each solution is different. The individual uncertainties are discussed in detail later on in the Pre-Lab preparation section.

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4.2 Dependent Variable Variable Unit How it was measured Potential Difference (V) ± 0.0005 The Potential Difference in the cell, between the electrodes, was measured using a digital multimeter, set to the 2000mV setting, which means that it measures voltage to the least count of 1mV. Thus, the uncertainty is ±0.5mV, or ±0.0005V. The multimeter ends were attached to the electrodes directly with insulated wires and clips, so that no external factor could influence the Voltage readings. The readings in the voltmeter were logged for 10 seconds manually, and all the readings that were recorded were averaged out to give a final reading for one trial. This was repeated five times for each concentration value of the electrolyte. Then again, these values were averaged, and one final value for the corresponding concentration was obtained. The data and calculation is shown further.

4.3 Controlled Variables Variable Unit How it was controlled Concentration of Zinc mol dm-3 The concentration of Zinc sulphate, the electrolyte of the sulphate anode, was kept constant, because the effect of change in concentration of one electrolyte. This was done by adding a fixed mass of the salt to a fixed amount of water, to get the desired concentration. The concentration was kept fixed at 1 mol dm-3 which would make calculations easier. The uncertainty in the concentration is discussed further in the Pre-Lab preparation section. Temperature Celsius (˚C) The temperature was kept constant as according to the Nernst Equation, the temperature at which the reaction is occurring, has a significant impact on the Ecell. Thus, the reaction was conducted in a controlled, air conditioned environment, with minimum temperature change. Also, the temperature of the cell was constantly monitored using a Pasco temperature probe of the uncertainty 0.05˚C. The value of temperature during one reading was noted. The values of temperatures were then averaged out. The variation in temperature was observed to be very small, around 0.1˚C, so keeping a mechanism to keep the temperature constant would not be totally justified.

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Salt bridge size The structure and material of the salt bridge was kept the same in all readings. The width of the U-shaped tubes used were also kept the same. This is because increasing the width would enable greater transfer of ions, while reducing the diameter would increase the resistance in the circuit. Immersed surface area of cm2 Throughout all the readings, the immersed surface area electrodes of the electrolytes was kept constant, by marking a specific height on the electrode, and then immersing the electrode to that level in all readings. This was important as if the surface area of the electrode immersed changes, there will be more atoms of the respective electrode reacting in the cell.

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5 Apparatus

Apparatus Capacity Uncertainty Quantity Glass beakers 300 cm3 - 3 Glass beaker 50 cm3 - 2 Copper Plate - - 1 Zinc Plate/Rod - - 1 Crocodile Clip wires - - 2 Digital Multimeter - - 1 Stopwatch - - 1 Pasco® Temperature Probe - - 1 Pasco® Data logger - - 1 U-Shaped glass tube - - 5 Plastic Syringe >20 cm3 - 1 Digital Balance - - 1 Measuring Cylinder 10 cm3 ± 0.1 cm3 1 Measuring Cylinder 250 cm3 ± 1 cm3 1 Volumetric flask9 250 cm3 ± 0.1 cm3 5 Dropper - - 1 Spatula - - 1 Stirrer - - 1 Bunsen Burner - - 1 Tripod Stand - - 1 Wire Mesh - - 1 Safety gloves - - 1 Magnetic Stirrer and Capsule - - 1 Pipette10 25 cm3 ± 0.03 cm3 1 Sand paper - - 1 Stands - - 2 Holders - - 2 Clamps for holder and stand - - 2

9, 9 http://academics.wellesley.edu/Chemistry/Chem105manual/Appendices/uncertainty_volumetric.html

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6 Pre-Lab Preparation

6.1 Preparing the Salt Bridge

6.1.1 Preparation 1. Weigh 0.4 grams of Agar in the Digital Balance and add it to a 50 cm3 beaker which is cleaned and dried. 2. Measure 10 cm3 of distilled water in a 10 cm3 measuring cylinder and add it to the 50 cm3 beaker with the Agar. 3. Stir the solution well. Put the beaker aside for now. 4. Weigh 1.011 grams of Potassium nitrate, as accurately as possible with a Digital Balance, and add it to another 50 cm3 beaker. 5. Measure 10 cm3 of distilled water in the 10 cm3 measuring cylinder and add it to the beaker with Potassium nitrate. 6. Stir well, until the salt completely dissolves. 7. Add the Potassium nitrate solution to the other beaker containing Agar. 8. Mix the solution well. 9. Set up the Bunsen burner, by placing a tripod stand over it, and a wire mesh on the stand. 10. Switch the burner on, and place the beaker with the solution of Agar and Potassium nitrate, on the mesh. 11. Stir the solution while heating, until the colour of the solution becomes completely transparent, from the hazy-white that it is, indicating that all the Agar has dissolved. 12. Switch the burner off immediately, and remove the beaker from over the stand. 13. Using a Syringe, remove the solution from the beaker, and start injecting it into the U- shaped tube, at a slow pace. 14. Make sure that no air bubbles remain in the process, and fill the tube until it is full from both ends. 15. Remove the remaining solution from the syringe, back into the beaker. 16. Take the U-shaped tube and place it carefully in a refrigerator, upright, so that the solution does not drip. 17. Wash both the 50 cm3 beakers with distilled water, and dry them. 18. Repeat Steps 1 to 16 for four more tubes, and place all of them similarly in the refrigerator.

6.1.2 Calculations Mass of Agar used Percentage of Agar: 2% Volume required: 20 cm3 Therefore, Mass of Agar required is: 2 × 20 푐푚3 = 0.4 푔푟푎푚푠 100 The uncertainty of this is ignored, because the mass of Agar used can be approximate, ranging from 2-5%.

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Mass of Potassium nitrate Molar Mass: 101.1032 g mol-1 Concentration required: 1 mol dm-3 Volume required: 10 cm3 Therefore, Mass required for 10 cm3 1 mol dm-3 solution is: (101.1032 푔 푚표푙−1) × 10 푐푚3 × 1 푚표푙 = 1.011032 푔 ≈ 1.011 푔 1000 푐푚3

In this case, the uncertainty can be ignored, as a small variation in the concentration would not have a drastic change in the resistance of the circuit. However, the concentration does need to be kept constant to some extent.

6.2 Zinc sulphate (1 mol dm-3, 250 cm3)

6.2.1 Preparation 1. Weigh 40.368 grams of Zinc sulphate salt as accurately as possible, using a Digital balance, and add it to a clean and dry 300 cm3 beaker. 2. Measure 250 cm3 of distilled water in the 250 cm3 measuring cylinder, and add it to the beaker with Zinc sulphate. 3. Place the beaker on a Magnetic stirrer, drop the stirring capsule in the beaker, and start the stirring process. 4. Stop the stirring when there is no sign of solid salt particles remaining in the solution. 5. Cover the beaker and keep it aside, to ensure that any remaining particles dissolve 6. Repeat Steps 1 to 5, for four more solutions, so that there are now five beakers.

6.2.2 Calculations Mass of Zinc sulphate Molar Mass of Zinc sulphate: 161.47 g mol-1 Volume required: 250 cm3 Concentration required: 1 mol dm-3 Therefore, the Mass of Zinc sulphate required to make 250 cm3 of 1 mol dm-3 solution is: 1 푚표푙 161.47 푔 푚표푙−1 × × 250 푐푚3 = 40.3675 푔 ≈ 40.368 푔 1000 푐푚3

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Uncertainty: Uncertainty in Mass = ±0.0005 g Uncertainty in Volume = ± 1 cm3 Percentage Uncertainty in Concentration – 0.0005 1 [( × 100) + ( × 100)] = [0.00124 + 0.4] = 0.40124% ≈ 0.40% 40.368 250 Absolute uncertainty – 0.4 × 1 푚표푙 푑푚−3 = ±0.004 푚표푙 푑푚−3 100

Final Concentration = 1 ± 0.004 mol dm-3

6.3 Copper sulphate (1 mol dm3, 250 cm3)

6.3.1 Preparation 1. Weigh 39.902 grams of anhydrous Copper sulphate as accurately as possible, on a Digital balance, and add it to a clean and dry 300 cm3 beaker 2. Measure 250 cm3 of distilled water in a 250 cm3 measuring cylinder and pour it in the beaker with the Copper sulphate. 3. Place the beaker over a Magnetic stirrer and drop the stirring capsule in the beaker. 4. Start the stirring process. 5. Turn off the stirring when there is no sign of any salt particles in the solution, and it is completely blue in colour. 6. Remove the capsule from the beaker using a spatula, and place the beaker aside for some time, to let any remaining particles dissolve completely. 7. Take a 250 cm3 Volumetric flask, and pour the salt solution in the flask until the solution is a little below the indication mark. 8. Take a clean dropper, and using it, fill the remaining solution into the flask from the beaker, until it is exactly touching the 250 cm3 mark in the flask. 9. Close the flask with a stopper. 10. Label the flask with the value of concentration of solution in it.

6.3.2 Calculations Mass of Anhydrous Copper sulphate Molar Mass: 159.609 g mol-1 Volume required: 250 cm3 Concentration required: 1 mol dm-3 Therefore, Mass of anhydrous Copper sulphate required for 250 cm3 of 1 mol dm-3 solution is: 1 푚표푙 159.609 푔 푚표푙−1 × × 250 푐푚3 = 39.90225 푔 ≈ 39.902 푔 1000 푐푚3

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Uncertainty: Uncertainty in Mass = 0.0005 g Uncertainty in Volume = ±1 cm3 Percentage Uncertainty in Concentration – 0.0005 1 [( × 100) + ( × 100)] = [0.00125 + 0.4]% = 0.40125% ≈ 0.40% 39.902 250 Absolute Uncertainty – 0.4 × 1 푚표푙 푑푚−3 = ±0.004 푚표푙 푑푚−3 100 Uncertainty of Volumetric Flask = ± 0.1 cm3

Final Concentration = 1 ± 0.004 mol dm-3 Final Volume = 250 ± 0.1 cm3

6.4 Copper sulphate (0.1 mol dm-3, 250 cm3)

6.4.1 Preparation 1. Take a clean and dry 25 cm3 pipette, and from the Volumetric flask containing 1 mol dm-3 Copper sulphate, pipette out exactly 25 cm3 of the solution. Be as precise as possible in reaching the 25 cm3 mark in the pipette. 2. In a new 250 cm3 volumetric flask, pour out the solution from the pipette. Make sure the solution completely drains out into the flask. 3. Fill a 250 cm3 measuring cylinder with 200 cm3 of distilled water, and pour it into the volumetric flask, which has the solution which you pipetted out. 4. Using a clean dropper, carefully fill in the remaining portion of the flask, up to the 250 cm3 mark, with distilled water. 5. Close the flask with a stopper. 6. Label the flask with the value of concentration of solution in it.

6.4.2 Calculations Volume of Copper sulphate pipetted out 3 Concentration of Copper sulphate present (M1): 1 mol dm -3 Concentration of Copper sulphate desired (M2): 0.1 mol dm 3 Volume of Copper sulphate desired (V2): 250 cm Volume of Copper sulphate needed to make 250 cm3 of 0.1 mol dm-1 solution from 1 mol dm-3 solution: V1

푀1푉1 = 푀2푉2 푀 푉 2 2 푉1 = 푀1

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0.1 푚표푙 푑푚−3 × 250 푐푚3 ∴ 푉 = = 25 푐푚3 1 1 푚표푙 푑푚−3

Uncertainty: -3 Uncertainty in M1 = ±0.004 mol dm 3 Uncertainty in V1 = ±0.03 cm (Uncertainty of pipette) 3 Uncertainty in V2 = ±0.1 cm (Uncertainty of Volumetric Flask)

Therefore, Percentage uncertainty in M2 – 0.004 0.03 0.1 [( × 100) + ( × 100) + ( × 100)] = [0.4 + 0.12 + 0.04]% = 0.56% 1 25 250 Absolute Uncertainty – 0.56 × 0.1 푚표푙 푑푚−3 = ±0.00056 푚표푙 푑푚−3 100

Final Concentration = 0.1 ± 0.00056 mol dm-3

Repeat the serial dilution process to create Copper sulphate solutions of the concentrations of 0.01, 0.001, and 0.0001 mol dm-3. The uncertainties for each concentration are calculated below.

Uncertainty for 0.01 mol dm-3: -3 Uncertainty in M1 = ± 0.00056 mol dm 3 Uncertainty in V1 = ±0.03 cm (Uncertainty of pipette) 3 Uncertainty in V2 = ±0.1 cm (Uncertainty of Volumetric Flask)

Therefore, Percentage uncertainty in M2 – 0.00056 0.03 0.1 [( × 100) + ( × 100) + ( × 100)] = [0.56 + 0.12 + 0.04]% = 0.72% 0.1 25 250 Absolute Uncertainty – 0.72 × 0.01 푚표푙 푑푚−3 = ±0.000072 푚표푙 푑푚−3 100

Final Concentration = 0.01 ± 0.000072 mol dm-3

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Uncertainty for 0.001 mol dm-3: -3 Uncertainty in M1 = ±0.000072 mol dm 3 Uncertainty in V1 = ±0.03 cm (Uncertainty of pipette) 3 Uncertainty in V2 = ±0.1 cm (Uncertainty of Volumetric Flask)

Therefore, Percentage uncertainty in M2 – 0.000072 0.03 0.1 [( × 100) + ( × 100) + ( × 100)] = [0.72 + 0.12 + 0.04]% = 0.88% 0.01 25 250 Absolute Uncertainty – 0.88 × 0.001 푚표푙 푑푚−3 = ±0.0000088 푚표푙 푑푚−3 100

Final Concentration = 0.001 ± 0.0000088 mol dm-3

Uncertainty for 0.0001 mol dm-3: -3 Uncertainty in M1 = ±0.0000088 mol dm 3 Uncertainty in V1 = ±0.03 cm (Uncertainty of pipette) 3 Uncertainty in V2 = ±0.1 cm (Uncertainty of Volumetric Flask)

Therefore, Percentage uncertainty in M2 – 0.0000088 0.03 0.1 [( × 100) + ( × 100) + ( × 100)] = [0.88 + 0.12 + 0.04]% = 1.04% 0.001 25 250 Absolute Uncertainty – 1.04 × 0.0001 푚표푙 푑푚−3 = ±0.00000104 푚표푙 푑푚−3 100

Final Concentration = 0.0001 ± 0.00000104 mol dm-3

6.5 Preparing the Electrodes

6.5.1 Preparation 1. Cut a small piece of sand paper, approximately 2 cm by 2 cm, and use it to polish the Copper plate. 2. Keep rubbing the plate with sand paper evenly, until the surface of the plate becomes shiny and golden in colour. 3. Measure and mark a point around 8 centimetres from the bottom on both sides, and draw a line across, using a marker. This will be a reference line for the electrode to be immersed in electrolyte. 4. Repeat the steps 1 to 3 for the Zinc electrode.

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7 Procedure

7.1 Setting up the Apparatus 1. Take two stands and keep them touched together so that the long side of the base of both stands touch, 2. Place one 300 cm3 beaker on the each of the base of the stands. 3. Using the clamp, fix the holders to the stands such that the gripping part of the holders rest exactly above the corresponding beakers. 4. On one of the holders, fix the copper electrode, and on the other, the zinc electrode. For now, position the holders at a height such that the electrodes do not enter the inside of the beaker. 5. Take a salt bridge and make sure the distance between the beakers is enough to immerse both ends of the bridge into the beakers, without touching either electrodes. 6. Attach a crocodile clip wire on the top of both electrodes so that it is completely in contact with the corresponding electrode. 7. Attach the other end of the wires to the wires of a Digital multimeter, such that the wire connected to the Zinc electrode connects to the negative terminal wire, and the wire connected to the Copper electrode connects to the positive terminal wire of the Digital multimeter. 8. Set the multimeter to the 2000mV setting.

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7.2 Investigation Procedure

Table 7.2.1 – Concentrations of the electrolytes to be used in the trials. Small variations in the volume do not affect the overall EMF generated. Concentration of Concentration of Volume of Trial Volume of Zinc Copper sulphate Zinc sulphate Copper sulphate number sulphate (cm3) (mol dm-3) (mol dm-3) (±0.4%) (cm3) 1 1.0000 ± 0.40% 1 225 225 2 0.1000 ± 0.56% 1 225 225 3 0.0100 ± 0.72% 1 225 225 4 0.0010 ± 0.88% 1 225 225 5 0.0001 ± 1.04% 1 225 225

1. Pour 225 cm3 of Copper sulphate (see concentration from Table 7.2.1 above) in the beaker above which the copper electrode is positioned, and pour the same amount of 1 mol dm-3 Zinc sulphate in the other beaker. 2. Lower the holders into the beaker by loosening the clamps, until the electrodes are immersed up to the mark which was made on them. 3. Make sure the Digital Multimeter is on, and set in the correct mode, which is 2000mV 4. Place the temperature probe inside one of the cells, and connect it to the Pasco Data logger. 5. Create a new experiment in the Data logger, making a table with Temperature and Time as the two variables. Set the time interval to be 1 second. 6. Introduce the salt bridge into the beakers. The current will start flowing. 7. Start the stopwatch and the Data logger logging at the same time. 8. Record the voltage readings from the multimeter for every second, till 10 seconds. If necessary, record a video of the multimeter screen for reference later. 9. Stop the stopwatch and data collection after 10 seconds. 10. Clean and dry the beakers and the electrodes properly for the next run. 11. Repeat the Steps 1 to 9 for the same concentration of Copper sulphate, four more times, which will give five sets of readings for each trial. Do this for all five different values of concentration. After every five readings, use another salt bridge. 12. Record the Data for a particular trial in a table similar to the one shown on the next page (Table 7.2.1). In the end, you should have five such tables, one for each trial.

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Table 7.2.2 – Blank table for observations of any one Trial with a particular concentration of Copper sulphate.

Concentration Reading 1 Reading 2 Reading 3 Reading 4 Reading 5 Trial of Copper Time number Sulphate Temp Voltage Temp Voltage Temp Voltage Temp Voltage Temp Voltage (°C) (V) (°C) (V) (°C) (V) (°C) (V) (°C) (V) 0 1 2 3 4 5 6 7 8 9 10

7.3 Safety Considerations  Wear Gloves to prevent chemicals from coming in contact with skin.  Wear a lab coat to prevent chemical spillage on clothes or skin.  Be sure not to use any rusted wires for the circuit.

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8 Data Collection and Processing

8.1 Qualitative Data During the investigation, there was no observable physical change.

8.2 Raw Data Collection The following tables show the data collected for the five trials. Each trial had a different concentration of Copper sulphate, as expressed, and five runs for each trial were conducted. For the tables, T is temperature and V is Voltage.

Table 8.2.1 – Readings for Trial 1, with Copper sulphate concentration of 1 mol dm-3

Trial [CuSO4] Reading 1 Reading 2 Reading 3 Reading 4 Reading 5 Time number (mol dm-3) T (°C) V (mV) T (°C) V (mV) T (°C) V (mV) T (°C) V (mV) T (°C) V (mV) 0 26.9 1094 27.1 1093 27.0 1095 26.9 1094 27.0 1092 1 27.0 1095 27.2 1093 27.1 1096 26.9 1095 27.0 1093 2 27.0 1095 27.2 1094 27.1 1096 26.9 1095 26.9 1092 3 27.0 1095 27.1 1094 27.0 1095 26.9 1094 26.9 1091 4 27.1 1093 27.1 1093 27.0 1096 26.9 1095 27.0 1092 1 1 5 27.1 1094 27.1 1095 27.1 1097 27.0 1095 27.0 1092 6 27.1 1095 27.2 1095 27.2 1095 26.9 1094 27.0 1092 7 27.0 1095 27.2 1094 27.2 1096 26.9 1095 27.0 1090 8 27.1 1094 27.2 1094 27.3 1095 27.0 1095 26.9 1092 9 27.1 1095 27.2 1093 27.2 1095 27.0 1096 26.9 1092 10 27.1 1094 27.1 1094 27.2 1095 27.0 1096 26.9 1091

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Table 8.2.2 - Readings for Trial 2, with Copper sulphate concentration of 0.1 mol dm-3

Trial [CuSO4] Reading 1 Reading 2 Reading 3 Reading 4 Reading 5 Time number (mol dm-3) T (°C) V (mV) T (°C) V (mV) T (°C) V (mV) T (°C) V (mV) T (°C) V (mV) 0 27.1 1065 27.2 1064 27.2 1067 27.0 1064 27.2 1066 1 27.1 1066 27.2 1065 27.2 1066 27.0 1065 27.2 1066

2 27.1 1065 27.2 1065 27.2 1066 27.0 1064 27.2 1064 3 27.1 1065 27.1 1064 27.2 1067 27.1 1066 27.2 1065 4 27.2 1067 27.2 1064 27.3 1067 27.1 1065 27.3 1064

2 0.1 5 27.2 1066 27.2 1065 27.3 1068 27.1 1065 27.3 1064 6 27.2 1066 27.2 1063 27.3 1066 27.1 1066 27.3 1065 7 27.1 1065 27.2 1064 27.2 1066 27.2 1066 27.3 1066

8 27.1 1064 27.1 1064 27.2 1067 27.2 1066 27.3 1066 9 27.1 1065 27.1 1065 27.2 1066 27.2 1065 27.2 1065 10 27.2 1066 27.1 1065 27.2 1066 27.2 1066 27.2 1066

Table 8.2.3 - Readings for Trial 3, with Copper sulphate concentration of 0.01 mol dm-3

[CuSO4] Trial Reading 1 Reading 2 Reading 3 Reading 4 Reading 5 (mol Time number dm-3) T (°C) V (mV) T (°C) V (mV) T (°C) Voltage T (°C) V (mV) T (°C) V (mV) 0 26.8 1032 27.1 1033 27.3 1035 27.1 1034 27.2 1033 1 26.8 1033 27.1 1034 27.3 1034 27.2 1033 27.3 1034 2 26.9 1034 27.2 1034 27.3 1034 27.2 1033 27.3 1034 3 26.9 1034 27.2 1035 27.4 1034 27.1 1034 27.3 1034 4 26.8 1033 27.2 1035 27.4 1035 27.2 1032 27.3 1034 3 0.01 5 26.8 1033 27.3 1035 27.4 1036 27.3 1034 27.2 1035 6 26.8 1032 27.3 1033 27.4 1035 27.3 1033 27.2 1033 7 26.9 1034 27.2 1034 27.3 1034 27.3 1033 27.3 1033 8 26.9 1033 27.2 1034 27.3 1034 27.2 1035 27.2 1034 9 27.0 1034 27.1 1035 27.3 1036 27.2 1034 27.2 1035 10 27.0 1033 27.1 1035 27.3 1035 27.3 1034 27.3 1035

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Table 8.2.4 - Readings for Trial 4, with Copper sulphate concentration of 0.001 mol dm-3

[CuSO4] Trial Reading 1 Reading 2 Reading 3 Reading 4 Reading 5 (mol Time number dm-3) T (°C) V (mV) T (°C) V (mV) T (°C) V (mV) T (°C) V (mV) T (°C) V (mV) 0 27.3 1006 27.3 1004 27.4 1005 27.2 1004 27.3 1006 1 27.3 1006 27.3 1004 27.3 1005 27.2 1004 27.3 1006 2 27.3 1006 27.3 1003 27.3 1006 27.2 1005 27.3 1005 3 27.4 1005 27.3 1004 27.3 1004 27.3 1003 27.3 1007 4 27.4 1006 27.3 1005 27.3 1005 27.3 1003 27.4 1006 4 0.001 5 27.5 1007 27.4 1006 27.4 1005 27.4 1004 27.4 1005 6 27.4 1005 27.4 1005 27.4 1006 27.4 1005 27.3 1004 7 27.4 1006 27.4 1005 27.4 1005 27.4 1005 27.3 1005 8 27.5 1006 27.4 1006 27.5 1005 27.3 1006 27.3 1005 9 27.5 1007 27.3 1005 27.5 1004 27.3 1005 27.3 1006 10 27.5 1007 27.3 1006 27.5 1005 27.4 1004 27.4 1005

Table 8.2.5 – Readings for Trial 5, with Copper sulphate concentration of 0.0001 mol dm-3

[CuSO4] Trial Reading 1 Reading 2 Reading 3 Reading 4 Reading 5 (mol Time number dm-3) T (°C) V (mV) T (°C) V (mV) T (°C) V (mV) T (°C) V (mV) T (°C) V (mV) 0 27.5 976 27.3 974 27.6 975 27.2 973 27.4 974 1 27.5 975 27.3 973 27.6 975 27.2 974 27.4 973 2 27.5 975 27.3 974 27.5 976 27.2 973 27.3 974 3 27.4 974 27.4 974 27.5 976 27.2 973 27.3 973 4 27.4 975 27.4 975 27.6 975 27.3 973 27.3 974 5 0.0001 5 27.5 976 27.3 974 27.7 976 27.2 972 27.3 974 6 27.5 975 27.3 974 27.6 976 27.3 972 27.4 975 7 27.6 976 27.4 975 27.5 975 27.4 973 27.4 975 8 27.5 976 27.5 975 27.5 974 27.4 974 27.4 974 9 27.5 976 27.5 975 27.6 975 27.3 973 27.5 975 10 27.5 975 27.4 976 27.5 975 27.4 974 27.5 975 20

8.3 Data Processing Firstly, all the voltage values and temperature values for a particular reading in a particular trial, were averaged.

Sample Calculation for Trial 1 Reading 1:

The values are obtained of the reading from Table 8.2.1.

Firstly, the values of Temperature were averaged out, to obtain a fixed value, which can be assumed constant throughout the reading. 26.9 + 27.0 + 27.0 + 27.0 + 27.1 + 27.1 + 27.1 + 27.0 + 27.1 + 27.1 + 27.1 ∴ ≈ 27.0 ℃ 11 Uncertainty:

The uncertainty in the average is calculated by the Standard Deviation. The formula for standard deviation of a set of values is11:

푛 1 휎 = √ ∑(푥 − 휇)2 푛 𝑖 𝑖=1

11 1 휎 = √ ∑(푥 − 27.0)2 = 0.0797724 ≈ 0.08℃ 11 𝑖 𝑖=1

Next, the values of Voltage were averaged out, to obtain the average EMF generated in the reading. 1094 + 1095 + 1095 + 1095 + 1093 + 1094 + 1095 + 1095 + 1094 + 1095 + 1094 ∴ 11 = 1094.5 ≈ 1095 푚푉 Uncertainty:

The Uncertainty because of standard deviation is as follows:

11 1 휎 = √ ∑(푥 − 1095)2 = 0.8528 ≈ 0.9 푚푉 11 𝑖 𝑖=1

11 https://www.mathsisfun.com/data/standard-deviation-formulas.html

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Similarly, the average values of Temperature and Voltage were calculated for all the readings, of all the trials. The results are shown in the tables below.

Table 8.3.2 – Average values of Temperature and Voltage of each reading of all the trials. The uncertainties in the readings have been derived from Standard Deviation values of each set. Trial 1 -3 Reading no. [CuSO4] (mol dm ) Temperature (°C) Voltage (mV) 1 27.0 ± 0.08 1095 ± 0.9 2 27.2 ± 0.07 1094 ± 0.7 3 1 ± 0.40% 27.1 ± 0.10 1096 ± 0.8 4 26.9 ± 0.06 1095 ± 0.7 5 27.0 ± 0.07 1092 ± 0.8 Trial 2 -3 Reading no. [CuSO4] (mol dm ) Temperature (°C) Voltage (mV) 1 27.1 ± 0.06 1066 ± 1.0 2 27.2 ± 0.06 1064 ± 0.7 3 0.1 ± 0.56% 27.2 ± 0.05 1067 ± 0.8 4 27.1 ± 0.08 1065 ± 0.8 5 27.2 ± 0.07 1065 ± 0.9 Trial 3 -3 Reading no. [CuSO4] (mol dm ) Temperature (°C) Voltage (mV) 1 26.9 ± 0.08 1033 ± 0.7 2 27.2 ± 0.07 1034 ± 0.8 3 0.01 ± 0.72% 27.3 ± 0.06 1035 ± 0.8 4 27.2 ± 0.07 1034 ± 0.9 5 27.3 ± 0.07 1034 ± 0.7 Trial 4 -3 Reading no. [CuSO4] (mol dm ) Temperature (°C) Voltage (mV) 1 27.4 ± 0.08 1006 ± 0.7 2 27.3 ± 0.06 1005 ± 1.0 3 0.001 ± 0.88% 27.4 ± 0.08 1005 ± 0.6 4 27.3 ± 0.08 1004 ± 1.0 5 27.3 ± 0.05 1006 ± 1.0 Trial 5 -3 Reading no. [CuSO4] mol dm Temperature (°C) Voltage (mV) 1 27.5 ± 0.05 975 ± 0.7 2 27.4 ± 0.08 975 ± 0.9 3 0.0001 ± 1.04% 27.6 ± 0.07 975 ± 0.7 4 27.3 ± 0.09 973 ± 0.7 5 27.4 ± 0.07 974 ± 0.7

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Next, the values of Temperature and Voltage obtained from each Reading set across a trial were averaged, in order to obtain the final Voltage and Temperature values for a particular concentration of Copper sulphate.

Sample Calculation for Trial 1:

The results of Trial 1 from Table 8.3.2 are taken in the calculations.

The calculation of the average Temperature value for Trial 1 is as follows: 27.0 + 27.2 + 27.1 + 26.9 + 27.0 = 27.04 ≈ 27.0℃ 5 Uncertainty:

The uncertainties of each of these values get added as the average is taken because the individual uncertainties come from the Standard Deviation. Hence, the final absolute uncertainty in the value of Temperature in Trial 1 is:

±(0.08 + 0.07 + 0.10 + 0.06 + 0.07)℃ = ±0.33℃

Therefore, the final value of Temperature for Trial 1 is 27.0 ± 0.33˚C.

Similarly, the average value of Voltage for Trial 1 is calculated: 1095 + 1094 + 1096 + 1095 + 1092 = 1094.4 ≈ 1094 푚푉 5 Uncertainty:

The uncertainties, like in the value of Temperature, get added. Hence, the final absolute uncertainty in the value of Voltage in Trial 1 is:

±(0.9 + 0.7 + 0.8 + 0.7 + 0.8)푚푉 = ±3.9 푚푉 Therefore, the final value of Voltage for Trial 1 is 1094 ± 3.9 mV.

This process was repeated for all the Trials to obtain final values of Temperature and Voltage for each concentration of Copper sulphate. The results are shown in the table below.

Table 8.3.4 – The values of Voltage and Temperature of each Trial with uncertainties. -3 Trial No. [CuSO4] (mol dm ) Temperature (˚C) Voltage (mV) 1 1.0000 ± 0.40% 27.0 ± 0.33 1094 ± 3.9 2 0.1000 ± 0.56% 27.2 ± 0.33 1065 ± 4.2 3 0.0100 ± 0.72% 27.2 ± 0.35 1034 ± 3.9 4 0.0010 ± 0.88% 27.3 ± 0.35 1005 ± 4.3 5 0.0001 ± 1.04% 27.4 ± 0.36 974 ± 3.7

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The graph of the concentration of the electrolyte against the voltage produced is shown below.

Concentration vs Voltage Produced 1120

1100

1080

1060

1040

1020

Voltage Produced (mV) Produced Voltage 1000

980

960 -0.1 0.1 0.3 0.5 0.7 0.9 1.1 [CuSO4] (mol dm-3)

Figure 8.3.1 – The graph of Concentration of Copper sulphate against the Voltage produced.

Next, for calculations, we convert the temperature readings from Celsius to the Kelvin scale. This is done by adding 273 K to the Celsius scale value.

Sample Calculation for the Temperature reading in Trial 1:

27.0 + 273 = 300퐾 Uncertainty:

The percentage uncertainty in the reading remains the same. Hence, the absolute uncertainty in the Kelvin reading is calculated as follows: 퐴푏푠표푙푢푡푒 푢푛푐푒푟푡푎푖푛푡푦 푖푛 푉푎푙푢푒 0.33 × 100 = × 100 = 1.22222 ≈ 1.22% 푉푎푙푢푒 27.0

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The same was repeated for the other trials. The results are shown in the table below.

Table 8.3.5 – The final values of Temperature (in Kelvin) and Voltage of each Trial, with uncertainties. -3 Trial No. [CuSO4] (mol dm ) Temperature (K) Voltage (mV) 1 1.0000 ± 0.40% 300.0 ± 1.22% 1094 ± 3.9 2 0.1000 ± 0.56% 300.2 ± 1.21% 1065 ± 4.2 3 0.0100 ± 0.72% 300.2 ± 1.29% 1034 ± 3.9 4 0.0010 ± 0.88% 300.3 ± 1.28% 1005 ± 4.3 5 0.0001 ± 1.04% 300.4 ± 1.31% 974 ± 3.7

To plot the graph of the Nernst plot, we use the equation given below, that was derived in Section 1. The equation is: 푅푇 푅푇 퐸 = ln([퐶푢2+]) + (퐸0 − ln([푍푛2+])) 푐푒푙푙 푛퐹 푐푒푙푙 푛퐹 Sample Calculation for Trial 1:

For Trial 1, the ideal value of the Ecell or the Voltage, can be calculated as follows: 푅푇 푅푇 퐸 = ln([퐶푢2+]) + (퐸0 − ln([푍푛2+])) 푐푒푙푙 푛퐹 푐푒푙푙 푛퐹 8.314 × 300.0 8.314 × 300.0 퐸 = ln(1) + (1.10 − ( ln(1))) 푐푒푙푙 2 × 96500 2 × 96500

퐸푐푒푙푙 = 0 + (1.10 − (0))

퐸푐푒푙푙 = 1.10 푉 = 1100 푚푉 Uncertainty:

It is important to realise that there is some uncertainty in the value which is calculated, because the parameters – Temperature, Concentration of Copper sulphate, and Concentration of Zinc sulphate, each have their own uncertainties. Thus, the uncertainty in the ideal value of Ecell for Trial 1 can be calculated by adding the percentage uncertainties of the measured parameters. The uncertainty for the value of Trial 1 was calculated as follows:

Percentage Uncertainty in Temperature = 1.22%

Percentage Uncertainty in [CuSO4] (From Pre-Lab Preparations, Section 6.3.2) = 0.40%

Percentage Uncertainty in [ZnSO4] (From Pre-Lab Preparations, Section 6.2.2) = 0.40%

푇표푡푎푙 푃푒푟푐푒푛푡푎푔푒 푈푛푐푒푟푡푎푖푛푡푦 = (1.22 + 0.40 + 0.40)% = 2.02%

Therefore, the Absolute uncertainty in the Voltage is: 푃푒푟푐푒푛푡푎푔푒 푈푛푐푒푟푡푎푖푛푡푦 2.02 × 푉푎푙푢푒 = × 1100 = 22.22 ≈ 22.2 푚푉 100 100 Therefore, the final value of voltage is 1100 ± 22.2 mV.

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Similarly, the ideal values of Voltage for the other trials were determined. The results are displayed below.

Table 8.3.6 – Ideal Voltage values along with the Experimental Voltage values for all Trials. [CuSO4] Ideal Voltage Experimental Trial no. Temperature (K) (mol dm-3) Value (mV) Voltage Value (mV) 1 1.0000 ± 0.40% 1100 ± 22.2 1094 ± 3.9 300.0 ± 1.22% 2 0.1000 ± 0.56% 1070 ± 23.2 1065 ± 4.2 300.2 ± 1.21% 3 0.0100 ± 0.72% 1040 ± 25.1 1034 ± 3.9 300.2 ± 1.29% 4 0.0010 ± 0.88% 1011 ± 25.9 1005 ± 4.3 300.3 ± 1.28% 5 0.0001 ± 1.04% 981 ± 27.0 974 ± 3.7 300.4 ± 1.31%

To plot the Nernst Plot, we have to use the logarithmic function of the value of Concentration of Copper sulphate, so that the graph can be analysed with greater ease. For simplicity, the Logarithmic function to the base 10 was selected.

Sample Calculation for Trial 1:

For the Trial 1, the log of concentration of Copper sulphate is:

log10[퐶푢푆푂4] = log10 1 = 0 Uncertainty:

Similarly, the other values were calculated for the respective trials. The percentage uncertainty in the initial value, and the value obtained after the log function, remains the same. The log of the concentration of Copper sulphate for the other trials is shown in the table below.

Table 8.3.7 – Values of the logarithmic function of the concentration of Copper Sulphate. Concentration of Uncertainty in Trial Uncertainty in Copper sulphate Concentration of log10([CuSO4]) number -3 -3 log10([CuSO4]) (mol dm ) CuSO4 (mol dm ) 1 1.0000 ± 0.40% 0 ± 0.40% 2 0.1000 ± 0.56% -1 ± 0.56% 3 0.0100 ± 0.72% -2 ± 0.72% 4 0.0010 ± 0.88% -3 ± 0.88% 5 0.0001 ± 1.04% -4 ± 1.04%

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Now, we can plot the Nernst plot, which is the Ecell values against the logarithmic value of the concentration of the electrolyte, which in this case is the concentration of Copper sulphate. The graph is shown in below.

Nernst Plot of Ideal Voltage values at different concentrations 1140 1120 1100 1080 1060

1040 (mV)

cell 1020 E 1000 980 960 940 -4.5 -4 -3.5 -3 -2.5 -2 -1.5 -1 -0.5 0

log10([CuSO4])

Figure 8.3.1 – Nernst plot of the Ideal Voltage produced at different concentrations of Copper sulphate, based on the calculations done previously.

The graph below shows the Nernst plot for the Experimental Voltage values, obtained through the investigation, against the logarithmic function of the concentration of Copper sulphate.

Nernst Plot of Experimental Voltage values at different concentrations 1120

1100

1080

1060

(mV) 1040 cell E 1020

1000

980

960 -4.5 -4 -3.5 -3 -2.5 -2 -1.5 -1 -0.5 0

log10([CuSO4])

Figure 8.3.2 - Nernst plot of the Experimental Voltage values at different concentrations of Copper sulphate, based on the readings taken in the investigation.

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The following graph shows a comparison between the two graphs plotted earlier – The Nernst plot of the Ideal voltage values and the Nernst plot of the Experimental voltage values. This will give us a clear insight and help us compare the trends, concluding if the Nernst equation is actually followed.

Comparision of the Nernst plots of Ideal and Experimental Voltage Values. 1140 1120 1100

y = 29.7x + 1099.8 1080 1060 1040 (mV) y = 30x + 1094.4

1020 cell E 1000 980 960 940 -4.5 -4 -3.5 -3 -2.5 -2 -1.5 -1 -0.5 0

log10([CuSO4])

Figure 8.3.3 – The comparison of the Nernst plots of the Ideal voltage value (blue) with the Experimental voltage values (red).

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9 Conclusion and Evaluation

9.1 Conclusion The experimental values of Voltage are quite close to the ideal voltage values, and fall in the range of the uncertainty of each value. The results from the investigation as well as the ideal voltage values calculated, are shown in Table 8.3.6. Clearly, the experimental values are close to the ideal values, but in all cases, slightly lesser than the original value. This trend is seen in Figure 8.3.3 where the trend line for Experimental voltage values (red) is below that of the Ideal voltage values, with roughly a similar slope – 30 in the case of the experimental values and 29.7 in the case of the ideal values. This tells us that the values obtained by experimentation, follow the trend of the ideal values, proving the Nernst Equation to be correct. The small reduction in experimental values could be due to the internal resistance of the circuit, which reduces the potential difference in the circuit. This resistance could be there due to the resistance in the salt bridge made, or the wires which are externally connecting the cell to the voltmeter.

The hypothesis made during the beginning of the investigation, was that the EMF or Voltage produced by the cell would increase, with the increase in the concentration of the electrolyte, Copper sulphate. By Figure 8.3.1, the graph of the concentration of the electrolyte, against the Voltage produced, this hypothesis is clearly proven correct. Furthermore, it is found that there is actually a logarithmic relationship between the concentration and the EMF generated, which indicates that initially, when the concentration is increased by a lesser degree, the change in the EMF produced is greater.

Thus, as we can see in Figure 8.3.2 and 8.3.3, the logarithmic function of the Concentration against the EMF generated, follows a linear trend, telling us that increasing the concentration would increase the Voltage generated.

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9.2 Evaluation Source of Error Type of Error Effect on the Investigation Possible Improvements Internal resistance of Random Increased internal resistance The resistance cannot be the circuit can reduce the potential in the completely eliminated, cell by a small amount, and but to make it as less as inconsistent internal resistance possible, the salt bridge in the cell could spoil the trend can be made with greater of the results obtained care. There should be no air bubbles remaining in the tube in order to ensure a continuous flow of ions. Small changes in the Random The temperature of the cell A mechanism which temperature of the directly affects the EMF keeps the temperature reaction generated, according to the constant in the cell can be Nernst Equation. An increase in used. This mechanism temperature increases the could simply be pasting a Voltage produced. In the wet filter paper around experiment, there was small, the beakers which but measurable variation in the contain the electrolyte, to temperature of the electrolyte. minimise heat exchange with the surroundings. Alternatively, the electrolytes, after they are poured in the beaker, can be kept aside so that they are equal to the room temperature completely, thus, minimizing heat transfer during the experiment. Presence of water in Systematic While making solutions of To avoid this, the beakers the volumetric different concentrations or should be washed with flasks/beakers conducting the reaction, the acetone so that after the containers may have some wash, the liquid remaining water drops in them evaporates, leaving the due to a previous wash. This flask dry. way, there might be variations in the concentration which can be very significant, when talking about concentrations as small as 0.0001 mol dm-3, or presence of impurities in the reaction which could have an effect on the overall voltage produced.

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10 Further Investigation

This investigation can be taken further by proving the Nernst Equation for other metals and their electrolytes at different concentrations. The trend can be observed in different cases and the values of EMF can be compared. Also, from the slope of the graph, the values of the Faraday constant and the Universal Gas constant can be proven. The effect of changing the chemical in the electrolyte on the Voltage produced can be investigated further, by using chemicals like Potassium chloride, Ammonium chloride, and etcetera.

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11 Bibliography

University of Colorado Colorado Springs http://www.uccs.edu/Documents/chemistry/nsf/106%20Expt9V-GalvanicCell.pdf

ChemWiki UCDavis http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Nernst_Equation http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrochemistry_2%3A_Galvan ic_cells_and_Electrodes

Chem1 http://www.chem1.com/acad/webtext/elchem/ec2.html

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