How duration affects the rate of electrolysis in a Voltaic Cell Essay Example for Free

How duration affects the rate of electrolysis in a Voltaic Cell EssayDesign and Conduct an prove to investigate the effect of ONE divisor on redox replys.Introduction-The dickens main comp onents of redox reactions are decline and oxidation. Reduction is a gain in electrons and the decrease in oxidation number whereas oxidation is the loss of electrons and the increase in oxidation number. Voltaic cells, also known as galvanic cells generate their own electricity. The redox reaction in a Voltaic cell is a spontaneous reaction. For this reason, voltaic cells are commonly used as batteries. Voltaic cell reactions supply heftiness which is used to perform work.The energy is harnessed by situating the oxidation and reduction reactions in separate containers, joined by an apparatus (known as the common brininess nosepiece which primarily completes a circuit and maintains galvanic neutrality) that allows electrons to flow. The functions of a voltaic cell are quite simple. Th ere happens to be an anode and a cathode. The positive ions go the veto electrode (anode) whereas the negative ions go to the positive electrode (cathode). Electrons always flow from the anode (where oxidation takes bit) to the cathode (where reduction takes place). Electrons flow across wires whereas ions flow across the electrolyte and the salt bridge.Aim-The objective of this experiment is to see how the metre affects the majority of the surface electrode (anode) and the tomentum electrode (cathode) in a voltaic cell.Variables-VariableType of variableHow it will be controlledTime (s) self-employed person (The one you throw)Values from 5 to 35 minutes will be usedMass of anode cathode (g)Dependent (The one you measure)Electrodes will be measured later each time intervalCurrent (A)ControlledMeasure the on-line(prenominal) with the help on an ammeter sign loudness of cathode and anode (g)ControlledWeigh out the electrodes using top travel balance from the beginning of t he experimentCharge on ionControlledUse the same solution for all the trials. The charge on the blur ion should be 2+ since the slob 2+ is being converted to blur metal. The charge on the zinc ion should be 0 because Zn is being converted to Zn 2+ assimilation of electrolyteControlledUse the same solution for all the trials. The solution primarily should be 1 mol dm-3 (just like standard conditions)Area of electrodes (cm2)ControlledMeasure the electrodes to ensure they defend the same dimensions (92.5cm). Use the same electrodes for all the trials.Volume of electrolyte (cm3)ControlledUse a measuring cylinder to measure out the electrolytes volumeAtmosphere which we are working underControlledPrimarily we are working under standard room temperature of 298 KApparatus-* 122.5cm2 copper electrode* 122.5cm2 zinc electrode* 100cm3 1mol dm-3 Zinc sulphate solution* 100cm3 1mol dm-3 copper (II) sulphate solution* Filter paper (required to create a salt bridge)* 100cm3 of potassium nitr ate solution (the spectator ion which I will require for creating the salt bridge which will complete the circuit and maintain electrical neutrality)* 2x two hundredcm3 beakers* Stopwatch (0.01s)* 1x100cm3 measuring cylinder (1.0cm3)* Voltmeter* 2 connecting wires* Top travel balance (0.01g)Method-1) present up the voltaic cell. Use a measuring cylinder to measure out 100cm3 of copper sulphate solution. Pour it into the 200 cm beaker.2) Next do the same for zinc sulphate. Use a measuring cylinder to help measure out 100cm3 of zinc sulphate solution. Pour it into a different 200 cm beaker.3) Weigh the aggregatedes of the electrodes separately using a top pan balance. Record the sign portiones.4) Connect the wires to the outlets in the zinc and copper electrode. smear them in the corresponding outlets of the voltmeter.5) After that we cut out some filter paper and dip that into our spectator ion (potassium nitrate) in order to build a salt bridge. The salt bridge will primarily complete the circuit, allow flow of ions and maintain electrical neutrality. The salt bridge will be placed in such a way that the ends of the salt bridge will be touching separate solutions of zinc sulphate and copper sulphate. The overall circuit should check the diagram in Figure.1.6) Place the zinc electrode into the beaker with the zinc sulphate solution and the copper electrode into the beaker with the copper sulphate solution and at the same time, start the stopwatch. bear on the stopwatch running until 200 seconds elapse. *Note- we will be recording the time every 5 minutes because 1 or 2 minutes barely isnt enough for the change to take place7) Take the cathode out of the solution and measure its mass (remember, before doing so, shake it a couple of times in order to remove any moisture). Record the mass. Do the same for the zinc electrode8) Place the electrodes into their respective solutions once again and start timing. Repeat steps 5 to 69) Repeat the same steps until we get mass readings for up to 60 minutes of experimenting.Data Collection and ProcessingRaw data- Initial mass of anode (zinc electrode) 31.29 0.01g Initial mass of cathode (copper electrode) 32.05 0.01gTable 1 Mass of anode and cathode obtained from different time intervalsDuration of electrolysis (0.21s)Mass of anode (zinc electrode) (0.01g)Mass of cathode (copper electrode) (0.01g)300.00 (5 minutes)31.2732.08600.00 (10 minutes)31.1432.16900.00 (15 minutes)31.0832.271200.00 (20 minutes)31.0032.421500.00 (25 minutes)30.8332.491800.00 (30 minutes)30.6132.802100.00 (35 minutes)30.2533.08Qualitative observations- We whoremaster see that the copper is deposited at the cathode where the cathode begins to get more than pink/ brownish colour. Blue colour of copper sulphate solution begins to get paler. Zinc electrode begins to corrode a bit. most corrosion can be observed at 35 minutes time interval.Note* UncertaintiesThe average reaction time was 0.5s even though it did falsify fro m interval to interval. Note that there is also a 0.01s time incredulity in the stopwatch itself. The uncertainty for mass is inscribed on the top pan balance as well.Data ProcessingWe must now calculate the mass changes which have taken place due to experimenting with different time intervals. (Different time intervals would result in a different mass change)This can be calculated simply by doing the followingMass change = final examination mass initial massDue note however that this formula can only be used for calculating the mass change taking place at the cathode (copper electrode where reduction takes place). This is because copper 2+ is being converted to copper metal and is being deposited at the cathode. Obviously this would result in a mass gain at the cathode. indeed, it would be better for us to use the formula Mass change = final mass initial mass so that it gives us a positive protect for the mass change taking place at the cathode.Example 1Mass change = final ma ss initial mass= 32.08 32.05= 0.03gExample 2Now to calculate the mass change taking place at the anode (zinc electrode), we use the following formula, Mass change = initial mass- final mass. In this case we use this formula because we know that the zinc is being oxidized to zinc 2+ leading the zinc electrode to corrode. This and then results in a decrease in mass of the anode (zinc electrode). Thus, it would be better for us to use the formula Mass change = initial mass final mass so that it gives us a positive apprise for the mass change taking place at the anode.Mass change = initial mass final mass= 31.29 31.27= 0.02Table 2 -Mass changes of anode and cathode for each time intervalTime (0.21s)Mass change of Anode (Zinc electrode)(0.01g)Mass change of cathode (copper electrode) (0.01g)300.00 (5 minutes)0.020.03600.00 (10 minutes)0.150.11900.00 (15 minutes)0.210.221200.00 (20 minutes)0.290.371500.00 (25 minutes)0.460.441800.00 (30 minutes)0.680.752100.00 (35 minutes)1.041.0 3Graph 1-Graph 2-To derive the equivalence for the two separate reactions, the number of electrons gained or lost during the process has to be deduced.The mass change per minute can be deduced from the gradient. Therefore we root calculate the gradient of graph 1 (mass changes for zinc electrode). For calculating the gradient, find two points which perfectly fits in the grid. In this case, the points (0.04. 100) and (0.08, 200)Gradient= (Y2 Y1) (X2 X1)= (0.08- 0.04) (200 100)= (0.04) (100)= 0.0004Therefore, the gradient of the start graph is 0.0002. So the mass change per minute for the anode is 0.0004.Next, we calculate the gradient of graph 2 (mass changes for copper electrode). To find the gradient, we work with the points (0.20. 500) and (0.24, 700)Gradient= (Y2 Y1) (X2 X1)= (700 500) (0.24- 0.20)= (200) (0.04)= 0.0002Therefore, the gradient of the first graph is 0.0002. So the mass change per minute for the cathode is 0.0002.The uncertainties also need to be prop agated through the summation of the fractional uncertainties.Uncertainties regarding zinc electrode-Fractional uncertainty of mass = absolute uncertainty actual value= 0.01 0.02= 0.500Fractional uncertainty of time = absolute uncertainty actual value= 0.21 300= 0.0007 = 0.001Total uncertainty = 0.001 + 0.500 = 0.501 to 3 decimal placesTherefore the rate of change is 0.004 0.501 g/sTable 3 yard of change for each time interval for anode (zinc electrode)Time (0.21s)Rate of change of anode (zinc electrode) (g/s)60.000.0040.501120.000.0040.067180.000.0040.048240.000.0040.035300.000.0040.022360.000.0040.015420.000.0040.001To calculate the number of electrons in zinc electrode, the following equation may be used-Number of electrons = molar mass mass of electrode (mass of one of the samples)= 65.37 31.27= 2.09Therefore, this would be the half-equation which would occur at the cathodeZn Zn2.09+ + 2.09e-Due to the loss in a bit more electrons compared to the theoretical formula, it would be a soakeder reducing instrument therefore the electrode potential would be swallow (more negative) than that of the original value. Nevertheless, the electrode potential cannot be determined.Uncertainties regarding copper electrode-Fractional uncertainty of mass = absolute uncertainty actual value= 0.01 0.03= 0.333Fractional uncertainty of time = absolute uncertainty actual value= 0.21 300= 0.0007 = 0.001Total uncertainty = 0.001 + 0.333= 0.334 to 3 decimal placesTherefore the rate of change is 0.002 0.334 g/sTable 3 Rate of change for each time interval for cathode (copper electrode)Time (0.21s)Rate of change of cathode (copper electrode) (g/s)60.000.0020.334120.000.0020.091180.000.0020.046240.000.0020.027300.000.0020.023360.000.0020.013420.000.0020.010To calculate the number of electrons in copper electrode, the following equation may be used-Number of electrons = molar mass mass of electrode (mass of one of the samples)= 65.50 32.08= 2.04Therefore, this would be the half-equation which would occur at the cathodeCu2.04+ + 2.04e- CuDue to the gain of a bit more electrons compared to the theoretical formula, it would be a slightly weaker oxidizing agent therefore the electrode potential would be slightly lower than that of the original value. Nevertheless, the electrode potential cannot be determined.ConclusionMy results show that as the duration/ time intervals increase, the mass of the anode (zinc electrode) decreases and the mass of the cathode (copper electrode) increases. We can see that there is a strong positive correlation between the time it takes for both electrodes to change in masses. If the duration is longer, then more electrons flow from the zinc electrode to the copper electrode (anode to cathode) through the electrical wires, while ions flow through the salt bridge to complete.As we know, in a voltaic cell/ galvanic cell, oxidation occurs at the anode (negative electrode) where as reduction occurs at the cathode (positive electrode). Primarily, zinc is oxidized at the anode and converted to zinc 2+. This causes corrosion at the zinc electrode due to the metal being converted to ions therefore the mass of the zinc electrode (anode) decreases. On the other hand, copper undergoes reduction at the cathode and the copper 2+ ions get converted to copper metal. This causes the copper metal to be deposited at the cathode thus leading to the copper electrode (cathode) to increase in mass as the duration is increased. The following anodic reaction takes place at the zinc electrode (this is the theoretical equation)-Zn (s) Zn2+ (aq) + 2e-However the equation we found experimentally is-Zn Zn2.09+ + 2.09e-Hence, this suggests that since the former zinc sample has more electrons to lose, it is an even stronger oxidizing agent compared to the theoretical equation and is slightly higher in the electrochemical series than the latter zinc samples. harmonise to the results that have been gathered, there is a positi ve correlation between the time it takes to electrolyse an aqueous solution and the rate of electrolysis. The rate of electrolysis was measured using the mass of cathode. If the duration of electrolysis is longer, then more electrons will flow through the circuit and more ions will flow from the anode to the cathode. Oxidation occurs at the anode whereas reduction occurs at the cathode. The cathode gains electrons therefore the mass decreases. The following reaction has taken place (although this is the theoretical equation)Cu2+ (aq) + 2e- Cu (s)However, the experimental equation isCu1.75+ + 1.75e- CuTherefore this implies that since the former copper sample has more electrons to gain, it is a stronger oxidizing agent and it is lower in the electrochemical series than the latter copper sample.The value of the electrode potential hasnt been calculated, however, the number of electrons is 25% off there that shows that there is a great leaving between the literature value and the ex perimental value. According to the graph in the previous page, there is a very strong positive correlation between the mass change and duration of electrolysis as can be deduced from the high R squared value. The change in mass over a certain period of time is very gradual because of the size of it of the electrons. Although a lot of electrons are able to flow through the electrolyte, there is not such a drastic change. By looking at the graph, intimately all the error bars for the points touch the line of best fit which means the data is fairly accurate.The theoretical mass of a copper electrode would be 31.75g. From the results that have been tabulated, the mass of a copper electrode is 36.21g.The percentage error can be calculated using the following formulaPercentage error = difference x 100theoretical value= 4.46 x 10031.75= 14.04%This shows that although there is not such a big difference between the theoretical value and the experimental value.EvaluationLimitationType of er rorImprovementThe mass of the anode was not measured therefore the rate of electron transfer between the two electrodes could not be determined. This could have increased or decreased the mass of the cathode.RandomMeasure the mass of the anodeThe power pack has internal resistance therefore not all the current was emitted. This could have decreased the current, thus decreasing the number of electrons produced.RandomUse a resistor to accurately measure the currentThe top pan balance had a zero offset error. This could have increased the mass of the cathode.SystematicUse the top pan balance with the 0.001 uncertainty to obtain more accurate values.a