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ExplorationIntroductionElectrochemistry two half-reactions. The reduction half-reaction which occurs in

ExplorationIntroductionElectrochemistry is the study of the chemical process which results in the movement of electrons, also known as electricity. A type of cell which uses chemical energy to produce electricity is called a Galvanic or Voltaic cell. A voltaic cell uses a spontaneous redox reaction to produce electricity. A battery is a common example of a voltaic cell. These batteries can be used in many electronic devices like smartphones or computers.  The essential parts of a voltaic cell are an anode (a negatively charged electrode), a cathode (a positively charged electrode), a salt bridge, two solutions containing cations and anions of the two electrodes, and a piece of wire. The current (the moving electrons) is allowed to flow due to the potential difference (voltage) between the two electrodes. This particular experiment aims to investigate how a change in the concentration of Zinc (II) Sulfate (ZnSO4) affects potential difference in the voltaic cell consisting of Zinc Sulfate (ZnSO4) and Copper Sulfate (CuSO4) solutions. The reduction reaction will take place at the cathode, where copper metal is present. The copper will be reduced during the reaction because it is less reactive than zinc and is located below zinc in the activity series. The oxidation reaction will take place at the anode, where zinc metal is present. Oxidizing and reducing are losing and gaining of electrons respectively. The net redox reaction can be given by the two half-reactions. The reduction half-reaction which occurs in the copper ions is as follows, Cu(aq)2++ 2 e- Cu(s) eq 1The oxidation half-reaction which occurs at the zinc anode is as follows, Zn(s)Zn(aq)2+ +2 e- eq 2 Therefore, the net reaction would be, Zn(s) + Cu(aq)2+Zn(aq) 2+ +Cu(s) eq 3In the reaction, sulfate ions are spectators ions, meaning they do not take part in the reaction and exists in the same form on the product side as the reactant side. This experiment is significant as it will allow in deducing what concentration of the zinc (II) sulfate solution yields the greatest voltage. Electrochemistry has various industrial uses as voltaic cells are an essential part in everyday life. Many researchers are attempting to create a cell, which is effective meaning it produces electricity for long periods of time, but is also cheap. Companies like Tesla are creating effective batteries to help households get off grid by storing solar energy. This has environmental implications such as the reduction of the consumption of electricity produced by coal, a non-renewable and harmful resource. Background TheoryFigure 1 – A Zinc-Copper Voltaic CellIn a voltaic cell, figure 1, the zinc metal is the anode and the copper is the cathode. The copper has a greater tendency to pull electrons (greater electronegativity), thus, when the two electrodes are submerged in sulfate solutions of their respective ions (ZnSO4 and CuSO4) and connected with a wire, the zinc oxidizes and loses its two electrons and those electrons travel through the wire to the copper ions in the CuSO4 and in the process the copper ions in the solution are reduced. The moving electrons is what produces electricity. As the zinc metal oxidizes, it starts to dissolve in the aqueous zinc (II) sulfate solution. On the other hand, the copper ions are reduced, and the mass of the copper metal increases.   A potential difference for a cell can be given by “Ecellwhich is the measure of the potential difference between two half cells in an electrochemical cell.” For the electrons to move from an anode to a cathode, there must be a difference in potential, which can be calculated using the equation, Ecello = Ecathodeo (reduction)+ Eanode (oxidationo) eq 4Alternatively, if only standard reduction potentials are known for the half-reactions, then the equation will transform into, Ecell = Ecathode (reduction)- Eanode (reduction) eq 5The superscript, , indicates “these potentials are correct only when concentrations are 1 M and pressures are 1 bar. A correction called the “Nernst Equation” must be applied if conditions are different.” (Libretexts) Under standard conditions, the cell potential of a zinc sulfate and copper can be calculated using eq 5, Ecell = ECopper (reduction)- EZinc (reduction) Ecell = 0.34 V – (-0.76 V) Ecell = 1.1 VStandard cell potential values referred from Diploma Programme Chemistry Data Booklet. Nernst Equation can be derived using the relationship between Standard Gibbs Free Energy(G) and Standard Cell Potential (Ecell). When Ecellis positive, the redox reaction will be spontaneous, if it is negative, the redox reaction will be non-spontaneous. From Thermodynamics, when Gwas negative, the reaction was spontaneous and when it was positive, the reaction was non-spontaneous. Therefore, G can be related to Ecellby the following equation, G = -nFEcell eq 6.1 Similarly, under nonstandard conditions, eq 6.1 will transform into G = -nFEcell eq 6.2From thermodynamics, eq 6.1 can be related to eq 6.2 via, G =G + RTlnQ eq 6.3 Substituting eq 6.1 and 6.2 into eq 6.3 -nFEcell = -nFEcell + RTlnQ eq 6.4Dividing both sides by -nF, the equation becomes, Ecell = Ecell – RTnFlnQ eq 6.5where , Ecell is the Cell Potential under nonstandard statesEcellis the Cell Potential under standard statesn is the amount of electrons, in moles, transferred in the reaction F is the Faraday’s Constant (96,500 C mol-1)T is the Temperature in Kelvins R is the Gas Constant (8.31 J K-1 mol-1)Q is the reaction quotient for the equation: aA + bB cC + dD Q = CcDdAaBbEq 6.5 is the Nernst Equation, formulated by Walther Nernst. The quotient reaction for eq 3 will become, Q =Zn2+CuCu2+Zn , and since the concentration of Cu and Zn is constant, they can be ignored from the equation. Q =Zn2+Cu2+ According to the equation above, if the concentration of zinc ions increases, the quotient reaction increases as well. As a result, the Ecell would be subtracted by a greater number and the Ecell , thereby, should decrease with an increase in concentration of zinc ions.   Safety, Ethical, or Environmental Considerations As copper (II) sulfate and zinc (II) sulfate are harmful when in contact with eyes, it was necessary to wear safety goggles at all times. Additionally, the chemicals should not be ingested by any means. They were kept at an arm’s length away from the experimenter. It is not safe to dispose of the Copper (II) Sulfate by putting it in the drain because it would react with the metal in the drain. Therefore, it is necessary to add some magnesium or zinc. The Copper (II) Sulfate solution becomes Magnesium Sulfate or Zinc Sulfate, which are relatively safer to dispose. The concentrated Zinc (II) Sulfate must be drained with lots of water as it is toxic to aquatic life. It was made sure that the wires were not around water to avoid any electrical harm. Research Question How does changing the concentration of Zinc (II) Sulfate solution affect the voltage produced in  a Zinc-Copper Voltaic Cell? Variables Independent Concentration of Zinc (II) Sulfate (ZnSO4) How it was manipulated – 5 different concentrations (3M, 2M, 1M, 0.75M, 0.5M) of ZnSO4  solution were created by dissolving dry ZnSO4 in distilled water. Dependent Voltage produced in the Zinc-Copper voltaic cell measured using a voltmeter. Controls Concentration of Copper (II) Sulfate (CuSO4) was kept constant by preparing 1 M of 1 L CuSO4 for each trial at once. Temperature was measured using a temperature probe. It was a control because as temperature is increased, the resistance increases as well. As a result, the voltage output decreases. Over a period of four days, the maximum temperature the probe read was 25.1±0.1oC and the minimum was 24.3±0.1oC. There was a difference of 0.8±0.1oC between the readings Voltmeter used in the experiment was kept the same. This would ensure that the systematic errors in the equipment remained constant and the data was precise. MethodsMaterials Voltmeter Copper and Zinc metalZinc (II) Sulfate Heptahydrate (ZnSO4•7H2O)Copper (II) Sulfate Pentahydrate (CuSO4•5H2O) Alligator WiresPorous CupDistilled WaterPipette Temperature Probe Vernier LabQuest Laptop Logger Pro1000mL Volumetric Flask 100mL Volumetric Flasks400mL Beakers 250mL Beakers 100mL Beaker25mL Graduated Cylinders Preliminary TrialThe first trial was conducted with only two conditions of the IV. 2M and 1M of ZnSO4 were prepared. To find what mass of ZnSO4 was required to prepared the respective solutions, the equation, moles = concentration * volume was used. At first, it was decided to use 50mL of each solution with 3 trials each. However, the volume was reduced to 20mL to avoid wasting. Additionally, as the lowest volume in the volumetric flask was 100mL, it was decided to change the number of trials from 3 to 5. The increase in trials reduced the random error and made the data more precise. Sample Calculations for Solutions 3M ZnSO4·7H2O Molar Mass – 287.54 g mol-1 moles = concentration * volumemoles = 3.000 M * 0.100Lmoles = 0.300 molesUsing stoichiometry to find the grams from moles and the molar mass, we get, grams = moles * molar massgrams = 0.300 moles * 287.54 g mol-1grams = 86.26 g(Similar calculation would be used for 2M, 1M, 0.75M, and 0.5M ZnSO4. Only the value for concentration would change.) Uncertainty in 3M ZnSO4·7H2O %Uncertainty = uncertaintyvalue100The mass of the zinc (II) sulfate heptahydrate was measured using a two point decimal scale. The percent uncertainty in the mass is, %uncert of m = 0.0186.26*100 = 0.01%A volumetric flask with an uncertainty of 0.0001L  was used to create a 0.1 L solution. The percent uncertainty in the volume is, %uncert of V = 0.00010.1*100 = 0.1%Therefore, the total uncertainty in the concentration of ZnSO4 is 0.01 + 0.1 = 0.11%1M CuSO4·5H2O Molar Mass – 249.68 g mol-1 moles = concentration * volumemoles = 1.000 M * 1.000Lmoles = 1.000 molesUsing stoichiometry to find the grams from moles and the molar mass, we get, grams = moles * molar massgrams = 1.000 moles * 249.68 g mol-1grams = 249.68 gUncertainty in 1M CuSO4·5H2O The mass of the copper (II) sulfate pentahydrate was measured using a two point decimal scale. The percent uncertainty in the mass is, %uncert of m = 0.01249.68*100 = 0.004%A volumetric flask with an uncertainty of 0.0001L  was used to create a 1 L solution. The percent uncertainty in the volume is, %uncert of V = 0.00011*100 = 0.01%Therefore, the total uncertainty in the concentration of 1M CuSO4 is 0.004 + 0.01 =0.014% Theoretical Voltage from the Nernst Equation (Eq 6.5) Ecell = Ecell – RTnFlnQThe temperature was kept under a certain range. Thus, it would be convenient to assume the temperature as 25.0oC or 298.15K to make the calculations easier. Since the temperature is a constant, along with the gas constant (R),Faraday’s constant (F), and the number of moles of electrons transferred (n=2) for all of the calculations, the Nernst equation can be rewritten as, RTF=8.314 298.1596500=0.0257 VEcell = Ecell – 0.02572lnQThereby, the theoretical voltage for 3M ZnSO4 solution will be, Ecell = 1.10 – 0.02572ln(31)=1.086VProcedure Preparation of the solutions First, a 1000.0±0.4mL of 1M CuSO4 solution was prepared in a volumetric flask by dissolving 249.68±0.01 g of CuSO4•5H2O in distilled water. Next, five 100mL of ZnSO4•7H2O solutions were prepared with concentrations 3M, 2M, 1M, 0.75M, and 0.5M. Data CollectionA temperature probe was used to measure the room temperature and ensure it remained in an acceptable range throughout the entire data collection process. A small volume of CuSO4 was poured into a 200mL beaker from the 1000mL volumetric flask to make the process of transferring the solution into the 25.00mL graduated cylinder. A 25.00mL graduated cylinder was used to measure out 20.00mL of 1 M CuSO4 into the porous cup.Similarly, a 25.00mL graduated cylinder was used to measure out 20.00mL of 3M ZnSO4 into the 250mL beaker. The porous cup containing the CuSO4 solution was placed into the 250mL beaker containing the ZnSO4. The zinc and copper metal pieces were attached to the alligator wires and the wires were attached into the voltmeter. The metal pieces were submerged into their respective solutions, i.e. zinc metal was placed into the zinc sulfate solution and the copper metal was placed into the copper sulfate solution. The voltage was read off the voltmeter.  After reading off the voltage, the solutions were replaced by new solutions. The old solutions were stored into two different beakers for disposal. Steps 3-11 were repeated for 5 trials for each of the 5 concentrations. DisposalMagnesium metal was added into the copper (II) sulfate solution to form magnesium sulfate. Magnesium sulfate was drained down the sink with water, and the copper remains from the solution were thrown into the garbage. Zinc Sulfate was drained down the sink with water. AnalysisQualitative DataWhen preparing ZnSO4 solution, the temperature of the volumetric flask decreased dramatically indicating that it was an endothermic reaction. Additionally, the zinc and the copper metal used changed their sizes. At the end of the data collection, zinc metal got smaller and the copper metal got bigger. Also, the copper (II) solution was visibly less blue as the reaction progressed. Table 1: Raw Data Concentration of CuSO4 (±0.014%M) Concentration of ZnSO4(±0.11%M) Voltage Produced (±0.001 V)Trials11.0003.000.58421.000 3.000.56631.000 3.000.57041.000 3.000.55551.000 3.000.556Average1.0003.000.566Concentration of CuSO4 (±0.014%M) Concentration of ZnSO4(±0.12%M) Voltage Produced (±0.001 V)Trials11.0002.000.382 21.000 2.000.482 31.000 2.000.488 41.000 2.000.464  51.000 2.000.464 Average1.0002.000.456Concentration of CuSO4 (±0.014%M) Concentration of ZnSO4(±0.13%M) Voltage Produced (±0.001 V)Trials11.0001.000.43421.000 1.000.44831.000 1.000.43041.000 1.000.45251.000 1.000.467Average1.0001.000.446Concentration of CuSO4 (±0.014%M) Concentration of ZnSO4(±0.14%M) Voltage Produced (±0.001 V)Trials11.0000.750.46621.000 0.750.370 31.000 0.750.32141.000 0.750.34551.000 0.750.337Average1.0000.750.368Concentration of CuSO4 (±0.014%M) Concentration of ZnSO4(±0.17%M) Voltage Produced (±0.001 V)Trials11.0000.500.21821.000 0.500.19631.000 0.500.21841.000 0.500.20251.000 0.500.231Average1.0000.500.213Sample CalculationMean x = Sum of all data valuesNumber of data valuesFor 2Mx = 0.382 + 0.482 + 0.488 + 0.464 + 0.4645x = 2.285x = 0.456Same steps for all concentrations Graph 1 – Voltage against Concentration of ZnSO4The graph above should have a negative and proportional slope according to the Nernst equation. However, the collected data does not follow the trend, indicating several errors. Despite this, the error bars on the concentration of ZnSO4 are quite small to the point where they are not visible. A cubic curve fit was chosen because it is the only fit which goes through all the collected data points. The graph shows a positive correlation between the two variables. The graph refutes the hypothesis that with an increase in the concentration of ZnSO4, the voltage would decrease. Additionally, the voltage produced hugely differs from the voltage produced under standard conditions, or the theoretical voltage which the cell should be producing with the given concentrations. A graph of the theoretical voltages is graphed below.  Table 2: Theoretical Voltage Produced Concentration of ZnSO4Voltage (V)3.001.0862.001.0911.001.1000.751.1030.501.109Graph 2: Theoretical voltage against Concentration of ZnSO4The graph above exhibits a strong, negative linear correlation between voltage produced and a change in concentration of ZnSO4. The correlation constant is -0.9715, which shows an almost perfect negative relationship between the two variables. EvaluationConclusion From the results of the experiment, it was evident that the graph of the raw data does not follow the trend explained by the Nernst equation. The graph is a cubic fit, which was not expected, thus refuting the hypothesis. This indicates the collected data is not reliable, despite having really small uncertainties. A cubic graph is the only fit which manages to stay close to the data point and within the error bars. The trend which the graph follows cannot be explained. It is also unclear as to why the voltage produced was only half of the standard cell potential, even with 1M of the solution. A comparison between the graph from the collected data and the graph of theoretical values shows the presence of several procedural errors. There is 47.88% error in the collected data. This is a relatively large percent error. Another huge difference between the two graphs is the voltage production when the concentration is less than 1M. The voltage should increase with a smaller concentration (less than 1M) as the value of Q from the Nernst equation will become negative, thus a small value is added to the standard cell potential. This is the case in the graph with theoretical voltages, however, Graph 1 shows the opposite and the smaller concentrations have a smaller production of voltage.   Evaluation Seeing that the graph from the experiment differs drastically from the graph with theoretical values, it can be deduced that there were several errors in the procedure of the experiment. One major weakness of the experiment was a lack of specific timing in reading the voltmeter. The readings on the voltmeter were not stable and were fluctuating. This hindered the ability to get an accurate reading of the voltage. Additionally, the values were read off the voltmeter too soon and reading time should have been delayed for a some time to allow the reaction to take place and get a stable reading for voltage. This is a systematic error, as the values were consistently lower than the expected. This can be improved by setting a specific time (at least a minute after the reaction had started) at which the values should be recorded. Another weakness can be attributed to the lack of control of the length of the zinc and copper metals (the anode and the cathode respectively). A longer piece of metal would have a greater surface area, thus the reaction would take place faster, i.e. more of the zinc ions would be produced with a longer piece of metal. This would be classified as a systematic error as well. This can easily be avoided by measuring out small and equal pieces of the two metals for each trail. It is also possible that the failure to calibrate the voltmeter to zero before taking the readings caused some errors in the values. However, there could be some errors which can be attributed to other factors such as the environment or random human error.  Finally, a major improvement in the experiment could be the modification of the research question. Instead of changing the concentration of zinc, which results in a graph with negative, linear slope, it could be modified to changing the concentration of copper, which should result in a graph with a positive, linear slope. The procedure would be identical, except the solutions would be switched. There was one strength in the procedure of the experiment, this is, all of the solutions used for each trial were prepared in one batch. This ensures minimal relative error in the solutions for each set. This experiment can be expanded upon by researching how temperature or the activity of the metal can affect the voltage produced. Additionally, one can also investigate if and how the surface area of the anode and cathode affects the voltage produced in a specific time period. Works CitedBrown, Catrin, and Mike Ford. Higher Level Chemistry. 2nd ed., Pearson Education, 2014.”Diploma Programme Chemistry Data Booklet”. Online curriculum centre. Cardiff: International Baccalaureate Organization, June. 2014. Print.Libretexts. “Nernst Equation.” Chemistry LibreTexts, Libretexts, 6 July 2017, chem.libretexts.org/Core/Analytical_Chemistry/Electrochemistry/Nernst_Equation.Libretexts. “Voltaic Cells.” Chemistry LibreTexts, Libretexts, 21 Mar. 2017, chem.libretexts.org/Core/Analytical_Chemistry/Electrochemistry/Voltaic_Cells.


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