Chemistry IA for IB Diploma Aim

ChemistryIA for IB Diploma

Aim

Investigatingon the relationship between the duration of time that the currentpasses through an electrolytic cell and copper metal produced at thecathode for the electrolytic cell using CuCl2(aq) as an electrolyte.

ResearchQuestion:

Howdoes the duration of time that a current is passed through anelectrolytic cell made of 0.1 M copper sulphate affect the mass ofcopper deposited on the cathode?

Introduction

Electroplatingis the process of coating a surface that conducts electricity with athin layer of metal using the electrolytic deposition. The process isused in various fields toincrease the value or the quality of a surface.

Duringelectrolysis, the cathode is the surface that is coated with themetal. The anode is the surface that provides with the metal thatcoats the anode, or it can rather be unaffected. The anode and thecathode are both put in an electrolyte that contains ions that willbe reduced to produce metal atoms that coats the cathode. The cathodeis connected to the negative terminal of an external power supply(battery) while the anode is linked to the positive of the powersupply (Ayşe Aygar 4). Reactions that takeplace in electrolysis are shown as follows

Cathode:Cu2+(aq)+ 2e→ Cu(s) (equation1)

Anode:Cl(aq)→ Cl2(g) + e (equation2)

Overall:CuCl2(aq) → Cu(s) + Cl2(g) (equation3)

Theamount of charge that passes through the electrolytic cell can becalculated by using the equation, (equation4).Qrepresents the amount of charge measured in Coulombs: Irepresents the current measured in amperes or coulomb per seconds. T representsthe time in seconds. Equation4is equivalent to Coulombs = Ampere x second. The charge is increasingas the time passed through the circuit increases.

Hypothesisaccording to Faraday’s Law of Electrolysis

Themass of the deposited copper on the cathode in an electrolytic cellreaction will increase with the time of current flowing. As theduration of time, the current flow increases, the mass of copperdeposited on the cathode of the electrolytic cell increases (equation4).According to equation1,number of moles of copper is dependent on the number of moles ofelectrons of current. More moles of electrons will produce more molesof copper. An electrical current is used in an electrolytic cell tosupply cell potential to push the electrons in the oppositedirection. Therefore, the longer that the current is flowing in theelectrolytic cell, the more charges will flow and produce morecopper.

Accordingto equation1,number of moles of copper is dependent on the number of moles ofelectrons of current. More moles of electrons will produce more molesof copper. An electrical current is used in an electrolytic cell tosupply cell potential to push the electrons in the oppositedirection. Therefore, the longer that the current is flowing in theelectrolytic cell, the more charges will flow and produce morecopper.

Controlledvariables:

  • The potential different was controlled by keeping it at 2.76V.

  • The concentration of the solutions of copper chloride throughout the experiment was 0.1M.

  • The volume of the solution is also constant throughout the experiment. Experimenters had to make sure that the volume of the solution was 150 ml throughout the experiment.

  • The resistance of the external circuit was constant throughout the experiment. As they maintain the resistance, experimenters can ensure that the same amount of current flow each trial. Since V (potential difference) = I (current)R (resistance) and the potential difference and resistance was constant, current also did not change. The same wires were used to maintain the resistance.

  • The temperature of the 0.1 M Copper Sulphate solution is constant at room temperature. I used the same electronic balance.

Requirements:

  • 150 ml volumetric flask

  • 25 ml graduated pipette (±0.1 ml)

  • Battery (D.C Power supply)

  • Two electrical wires

  • Stopwatch (±0.2s)

  • 200 ml beaker

  • Two graphic rods

  • Electronic scale (± 0.001g)

  • Sandpaper

  • Colored tape (red and blue)

  • M 15.0 ml CuCl2 (aq)

  • Distilled water 135 ml

Procedure:

Dilutionof the 1.0M Copper Chloride Solution:

  1. 15.0 ml 1.0 M aqueous CuCl2 solution was measured with a graduated pipette.

  2. I calculated the volume of distilled water required to prepare 0.10 M CuCl2 (aq) by using following equation 5.

C1V1= C2V2(equation5)

15.0ml1.0M= X0.10 M

X= 150 ml

Thereforeto prepare 0.1 M, 150 ml CuCl2(aq), I used a pipette measure 15ml of CuCl2andused a beaker to measure 135ml distilled water.

Preparationof an Electrolytic cell:

3.I poured 50 ml of 0.1M copper chloride in a 200ml beaker.

4.I taped the red tape on one graphite electrode and a blue tape on theother graphite so that I could later distinguish between the two.

5.I recorded the mass of the blue taped cathode.

6.I fixed the position of the graphite using a tape so that I placedeach on the opposite side of each other.

7.I connected the positive charge on the blue taped graphite andconnected the negative charge on the red taped graphite, and this waslater helpful to distinguish between the two.

8.I turned the battery and immediately started the stopwatch and letthe reaction happen for 2 minutes with a current of 2.30 A.

9.I removed tape of the blue taped graphite rod and weighed the mass ofthe cathode that copper has deposited.

10.I used a hair dryer to dry off the aqueous solution so thatexperimenters did not record the mass of the graphite rod includingthe aqueous solution.

11.I recorded the final Mass.

12.Using sandpaper, I removed the deposited copper on the graphite rod.

13.I repeated this process was three times. I calculated the averagemass of copper deposited.

14.After repeating, collecting three data for when the current aspassing through 2 minutes, I repeated procedure 1~12 for three timeseach for 4 minutes, 6 minutes, 8 minutes, and 10 minutes.

Figure1.Above shows the general view of the experiment.

DataCollection:

Qualitativeobservation

The0.1M copper chloride aqueous solution had a lighter blue colorcompared to the 1.0 M of chloride solution. When current was notflowing through the graphite rods, nothing was happening inside thecopper chloride solution. As soon as the power cell was turned on andthe current started to flow into the graphite rod gas bubbled fromthe graphite rod that I connected to the positive charge from theanode rod it was assumed that there was production of chlorine gasbecause of the smell. Therefore, the experiment had to be conductedunder a fume hood. After each reaction, the cathode graphite rod hadbrown copper deposited on the rod where the rod was in touch to thecopper chloride solution. The longer the current passed through thegraphite rod, it could be seen that more the copper deposited on theanode graphite rod.

QuantitativeData

FollowingTable 1 shows the initial and final mass of the anode graphite rodfor each trial when the current passed for 120 seconds, 240 seconds,360 seconds, 480 seconds, and 600 seconds. The uncertainty of theinitial and final mass of the graphite rod came from the electricalscale. Since this electrical scale could record to the 0.001, and itis a digital equipment, the uncertainty for each initial and finalmass of the graphite is 0.001g.

Table1.The mass of a cathode before and after the current is flown.

Time that the current passed through the electrolytic cell

( ± 0.2 second)

Initial mass

g ± 0.001g

Final mass

g ± 0.001g

120

Trial1

6.384

6.386

Trial2

6.375

6.379

Trial3

6.374

6.377

240

Trial1

6.362

6.372

Trial2

6.346

6.357

Trial3

6.338

6.344

360

Trial1

6.321

6.336

Trial2

6.315

6.330

Trial3

6.295

6.309

480

Trial1

6.291

6.312

Trial2

6.269

6.291

Trial3

6.261

6.280

600

Trial1

6.252

6.277

Trial2

6.245

6.272

Trial3

6.225

6.251

DataProcessing:

Table2 below shows the mass difference for each trial in the experiment.

Time that the current passed through the electrolytic cell

(± 0.2 seconds)

Mass difference / g ± 0.002g

120

Trial1

0.002

Trial2

0.004

Trial3

0.003

240

Trial1

0.010

Trial2

0.011

Trial3

0.006

360

Trial1

0.015

Trial2

0.015

Trial3

0.014

480

Trial1

0.021

Trial2

0.022

Trial3

0.019

600

Trial1

0.025

Trial2

0.027

Trial3

0.026

Table2.The mass difference of the cathode in grams for the different timeof current flowing

Samplecalculation

TheMass difference of the initial and the final graphite rod for trial1, when the current passed through the electrical cell for 600seconds is:

(6.277± 0.001) – (6.252 ± 0.001) = 0.025 ± 0.002 g

Thereforethe mass difference is:

0.025± 0.002 g

Theaverage mass change is calculated for different time of currentflowing and displayed on table 3.

Time that the current passed through the electrolytic cell / ± 0.2 seconds

Average mass difference / g ± 0.002g

120

0.003

240

0.009

360

0.015

480

0.021

600

0.026

Table3.Average mass difference for different times that passed through theelectrolytic cell

Samplecalculation for average mass difference of the graphite rod for trial1 when the current passed through the electrical cell for 600seconds:

g

Therefore,the average mass difference when current flowed the electrolytic cellfor 600 seconds is 0.026 ± 0.02 g.

Figure1 displays the results. The time that current passed through theelectrolytic cell is set as x-axis (horizontal axis) and the averagemass of copper deposited is set as y-axis (vertical axis).

Figure1.The current passed through the electrolytic cells plotted againstaverage mass deposited.

Thegeneral tendency line is increasing a function that indicates thatthe mass of copper deposited is increasing as the time passes throughthe cell is increasing.

Conclusion:

Theaim of this investigation was to find out the relationship betweenthe time that the current flowed through the electrolytic cell andthe mass of copper deposited. From the results obtained it could besaid that timeis directly proportional to the amount of coated metal. Thismeans that at a particular period, a certain quantity of metal isproduced at the cathode (Ayşe Aygar 9).

Themass of the copper deposited was the dependent variable, and this wasfigured out through subtracting from the initial mass of graphite rodto final mass rod. The independent variable was the time that thecurrent flowed which varied from 120 seconds, 240 seconds, 360seconds, 480 seconds, and 600 seconds.

Themass of copper deposited on the cathode in an electrolytic cellreaction will increase with the increase in time of current flowing.Number of moles of copper is dependent on the number of moles ofelectrons of current. More moles of electrons will produce more molesof copper. An electrical current is used in an electrolytic cell tosupply cell potential to push the electrons in the oppositedirection. Therefore, the longer that the current is flowing in theelectrolytic cell, the more charges will flow and produce morecopper.

Thetrend line agrees with the equations partially. The trend line had anequation of y=510-5x- 0.003. Examining the above equations, if I connected the current tothe electrolytic circuit for 0 seconds there should be no coppermetal deposited. However, the trend line has a y-intercept of -0.003,this shows that there is a source of systematic error. The R2value is 0.99881, and this shows that there is a strong positiverelationship between x-axis (time that current flow) and the y-axis(average mass of copper deposited).

Accordingto Faraday’s 2ndLaw of Electrolysis

Where:

  • n is the number of moles of metal produced

  • I is the magnitude of the current

  • t is the time

  • F is Faraday’s constant

  • z is the oxidation state of metal. (Ayşe Aygar 19)

Oxidationstate for copper is 2

Faraday’sconstant is 96.5KJmol1

Magnitudeof the Current supplied is 2.30 A

Timetaken is 600 s

nx 96.5x 103x 2 = 2.3×600

n=0.00715 Mol

Mass=Molx Molarmass

Molarmass of copper = 63.546

Mass=0.00715×63.546

Theoreticalmass of copper deposited after a period of 600 s = 0.454g

%error = (Theoretical mass –mass of copper deposit weighed /Theoretical mass) * 100

%error= (0.454- 0.026)/0.454×100= 94.27%

Thepercentage error of the mass difference was 94.27%.

Error Type

Description of Error

How to Improve

Random Error

Insufficient lab time that forces us to perform only one trial at each temperature, instead of three for each temperature to be more precise. Hence, we end up with data that is less accurate.

Conducting more trials to produce more reproducible results. After the collection of all data, it should be averaged and graphed.

Random error

The stopwatch has an uncertainty of 0.2s, together with human reaction time while using the stopwatch during the experiment further increasing the uncertainty of the data.

Conducting more trials and practicing timing with the stopwatch may ensure efficient and precise use of the stopwatch and the random error would reduce.

Systematic Error

The copper might be lost to the solution while transferring the graphite rod for measuring the mass.

Use Logger Pro to detect the mass change in order to prevent the loss of copper.

Systematic Error

The mass of copper obtained by electrolysis is too little to cause high value of percent uncertainty, this is because the mass collected depends on the length of time the current runs. Maybe the low change mass was a result of not running the current for long enough. Our time ranges within t=120s and 240s.

Increase the time of the procedure, i.e. add two extra times intervals after 600s.

Oneof the disadvantages of this experiment is that the independentvariable was hard to control. The power cell was turned off accordingto the stopwatch. Although I ensured each trial of the experimentswas stopped within the time, it was hard to control it. It isquestionable whether such variable is significant.

Reference

Aygar,Ayşe. &quotInvestigation of the Factors That Affect the Amount ofMetal Coated in an Electroplating Process.&quot InternationalBaccalaurete Diploma Programme (2009).

Derry,Lanna, Maria Connor and Carol Jordan. Chemistryfor use for the IB Diploma Standard level.Melbourne: Pearson Education, 2008.

Neuss,Geoffrey, IBDiploma Programme Chemistry Course Companion.Oxford: Oxford University Press, 2007.

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