Electrochemistry Laboratory Electroplating of Metals By Thomas …jhwang43/CHM116Labs... · 2020. 8. 8. · Electrochemistry Laboratory ... The electroplating that you are conducting

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Name: ______________________________________________ Date:____________________

Lab Partners: ___________________________________________________________________

Electrochemistry Laboratory – Electroplating of Metals

By Thomas Cahill, Arizona State University, New College of Interdisciplinary Arts and Sciences.

Background:

Electroplating is a common means to deposit a thin layer of some expensive metal like gold or

silver onto an item made from a cheaper metal such as steel. The advantage is to impart the desired

characteristics of the surface metal, such as aesthetics and corrosion resistance, without making the

entire item from the expensive metal. Electroplating is commonly conducted on silverware where a

stainless steel base is coated with a thin layer of silver to give it a desired appearance. Electroplating is

also commonly used for electrical connections where gold is plated onto the connectors to prevent

corrosion and to improve electrical conduction.

Electrochemistry is abundant in metallurgy, but that use of electrochemistry is not often

appreciated. Many of the common elements are produced from electrolysis of salts and solutions.

Aluminum metal is probably the greatest use of electrolysis, and the production of aluminum metal used

to account for about 6 to 8% of the total electricity used in the United States. Electroplating is also

commonly used as a purification mechanism for some of the more exotic metals such as gold. In this

case, the impure gold mixture is dissolved in a cyanide solution that is selective for dissolving gold. The

solution is then transferred to an electrochemical cell to electroplate the pure metal out of solution.

Your laboratory exercise is basically identical to this last application of electrochemistry. You

will recover pure metals from salt solutions. Most of your salts will be metal nitrate or metal acetate

salts such as silver nitrate & copper acetate. You will not get the ideal thin, uniform layers of metals that

are the typical image of electroplating, but you will get metal crystals and aggregates that you will

collect, dry and quantify.

Processes involved:

Electrochemistry involves the use of direct current (DC) electricity, such as that from a battery,

to precipitate a pure metal. Normal power outlets in your house are alternating current (AC) and cannot

perform this type of process. The basic process is that two electrodes are placed in a salt solution

containing metal cations. The electrodes are attached to a DC power supply that makes one of the

electrodes negatively charged and the other electrode positively charged. The metal cations are attracted

to the negatively-charged electrode (called the cathode). When the ions reach the cathode, they acquire

enough electrons to make them a neutral atom. Once the metal has a neutral charge, it will precipitate

on the cathode, which gives the pure metal that we want.

Reactions:

The electroplating that you are conducting in lab is the result of two redox reactions. The metal

ion is being reduced at the cathode to form metal, so the metal ion is gaining electrons. However, to

have a complete circuit, we need to oxidize something at the anode to give up electrons. Since we have

a pure water + metal salt system, water is oxidized at the anode to form oxygen. Therefore, the two

half-reactions are as follows (using silver nitrate example):

Cathode: Ag+(aq) + e− → Ag(s) Ered = +0.799 V

Anode: 2H2O(l) → O2(g) + 4H+(aq) + 4 e− Ered = +1.23 V

Ered = −0.43 V

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Since the net energy is negative, this reduction is not spontaneous, thus a voltage must be applied

to make this reaction proceed. This voltage is supplied by the battery. The end result of the reactions is

Ag(s) and H+(aq), which means that you end up with a rather acidic solution at the end of the reaction,

so handle it with care.

Battery capacity:

The easiest means to determine the number of electrons that pass through the electroplating cell

is to use the rated charge capacity of the battery and then allow the battery to completely discharge

before collecting the metal. The charge capacity of a battery is generally given as milli-ampere hours

(mAh). The units of milli-ampere hours tells you that the battery has a charge equivalent to the stated

current (in milli-amperes) for one hour. For example, AA batteries typically have a capacity of 1800

mAh, which means they have a charge equivalent to 1800 mA for a single hour. Since an ampere is a

coulomb per second (A= C/s), the coulombs in the battery can be determined (C = A·s) by converting

the milli-amperes into amperes by dividing by 1000 and by multiplying by 3600 sec/hr.

To calculate the number of electrons that passed through the cell, we divide the coulombs by Faraday’s

constant (1 F = 96,485 C/mol e−)

Mole e− = (mAh)(1A/1000mA)(3600s/1h)(1/F)

For example, the charge capacity of an 1800 mAh battery is:

(1800 mAh)(1A/1000mA)(3600s/1h)(1/96485C/mol e−) = 0.0671 mole electrons

You will use this calculation in your experiment since your batteries will be completely discharged at

the end of the experiment.

The rated battery capacity is the total charge the battery can give up. Unfortunately, only about

85% of this charge is actually useable in our experiments. Therefore, you will need to multiply the rated

battery capacity by 0.85 to find the useable charge in this experiment. As the battery gets weaker, the

voltage drops to the point where the battery cannot drive the reaction, so the reaction stops. This is a

combination of both the battery getting weaker and the “back voltage” of the electroplating cell

increasing as you reduce the metal salts in solution. You can measure this back voltage at the electrodes

when the battery has been disconnected.

Multiple Batteries in Series:

You will use multiple batteries in series (two or more batteries end to end) in some of your

experiments. When batteries of the same type are put in series, the total voltage will double but the total

electron capacity will remain the same. You will have the same number of electrons, but each electron

will have twice the energy. Therefore, you will use the capacity (mAh) from a single battery to calculate

the moles of electrons.

Moles of Metal Condensed:

Once the moles of electrons that passed through the cell are known, then it is a simple matter to

determine the moles of the metal condensed. The moles of metal condensed:

Moles metal = (moles electrons)

(charge of metal ion)

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Therefore, it takes more electrons to condense a +2 ion element (e.g. Cu2+) compared to a +1 element

(e.g. Ag+). The mass of metal condensed (in grams) is the moles of metal condensed multiplied by the

molar mass of the element.

For example, the amount of Cu2+ condensed by an 1800 mAh AA battery is:

Coulombs: 1800mAh × (A/1000mA) × (3600 sec/hr) = 6480 C

Mole e− = Coulombs/F = (6480C)/(96,485C/mol e−) = 0.0672 mol e−

Grams Cu = 0.0672 mole e− × 63.5 g/mol Cu = 2.13 g copper metal

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Part A: Laboratory Experiment

Quantitative recovery of pure metals from metal solutions.

You will be working in groups of 4 or 5 students for this part of the lab.

Materials: (per single setup)

AAA rechargeable batteries (number varies depending on solution)

Two stainless steel spatulas as electrodes

Short (20 to 30 cm) length of flexible wire with bare ends

Wooden electrode holder

Battery holder (you can use tape, but it is not preferred because it leaves a sticky residue)

Volt meter (with alligator clip connections if possible)

One 250 mL beaker

200 mL of a metal salt solution (silver or copper)

Graduated cylinder (to measure the volume of the solution)

Three small glass vial for the final product (two for copper, one for silver)

Sticky labels or tape to identify your electroplating setups.

Diagram of Experimental Apparatus

(using the silver nitrate solution as an example)

Battery+ −

250 mL beaker

with 200 mL

solution

wooden

electrode

holder

wirewire

NO3−

e−

e−

e−

e−

Ag+ + e-

Ag

+

e−

e−

+

+

+

+

++

e−

+

cathode

electrode

anode

electrode

~1.3 V

H2O

2H+ + ½O2 + 2e-

Battery+ −Battery+ −

250 mL beaker

with 200 mL

solution

wooden

electrode

holder

wirewire

NO3−

e−

e−

e−

e−

Ag+ + e-

Ag

Ag+ + e-

Ag

+

e−

e−

+

+

+

+

++

e−

+

cathode

electrode

anode

electrode

~1.3 V

H2O

2H+ + ½O2 + 2e-

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One online computer simulator, and one excel file, and some materials you need to gather at home (see

Part B)

The excel files are locked. To unlocked it use the password: unprotect

You will be testing three solutions:

#1 - 200 mL of 0.080 M copper(II) acetate with two AA batteries in series.

#2 - 200 mL of 0.175 M copper acetate with three AAA batteries in series.

#3 - 200 mL of 0.05 M silver nitrate with one AAA battery.

PROCEDURE:

For solution #1: 200 mL of 0.080 M copper(II) acetate, Cu(C2H3O2)2(aq)

Copper Experiment I 1) In the excel file, add 200 mL of 0.080 M copper acetate, Cu(C2H3O2)2(aq), in the 250 mL

“beaker”, by enter 200 in the “ ” spot. Record the volume of the solution (200 mL) and the

molarity of the solution (0.080 M) on your data sheet.

2) On the data sheet, convert the 200 mL to liters. Multiply the liters by 0.080 M to get the moles

of Cu.

3) On the data sheet, multiply the moles of Cu by the molar mass of Cu, to get the “Grams of

copper available for electroplating”.

4) In the excel file, use the drop down arrow, choose the number and type of batteries by clicking

on “two AA” in the “ ” spot.

5) In the excel file, read the “Battery Charge Capacity, mAh” and record it on your data sheets as

“Reported battery charge”. This value is expressed as milli-ampere hours (mAh). Then divide

mAh by 1000 to convert to Ah. The batteries used in this experiment have a usable charge of

about 85% before the voltage drops below the voltage necessary to drive the electroplating.

Therefore the “Useable battery charge” should be the Ah multiplied by 0.85.

6) In the excel file, enter the starting date and time in the “ ” spot, example 4/3/2020 1:21 pm.

Record the starting date and time on data sheet, though, it is not used in any calculations.

7) In the excel file, read the “Starting battery voltage, V”, and record it on your data sheet as

“Starting voltage across the electrodes”. Check to make sure you are getting the correct voltage

reading for each experiment based on the following ranges:

1.3 to 1.5 V for one battery

2.5 to 2.95 V for two batteries in series

3.8 to 4.5 V for three batteries in series

8) The voltage measurements are not used in any calculations, but they simply test that all the

connections are good and the proper voltage is reaching the electrodes.

9) The cations (positively charged ions) will be attracted to the negatively-charged cathode

electrode (which is the electrode attached to negative end of the battery). When the cations reach

the negatively charged cathode, they will gain an electron and condense on the electrode as the

metal. In real experiment, the formation of copper metal occurs slowly, shiny copper crystals

will form over the course of 36 to 48 hours.

10) In real experiment, if everything is working properly, you should see tiny (<0.5 mm) bubbles

forming on the anode electrode (the one attached to the + end of the battery) in a couple of

minutes. The bubbles should break off and rise to the surface. If you do not see any bubbles

rising to the surface, then check all your electrical connections. You will not see the bubbles in

the excel file simulation.

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11) In the excel file, enter the end date and time in the “ ” spot, example 4/10/2020 1:21 pm.

Make sure the ending data and time is one week later than the starting data and time.

Record the ending date and time on data sheet, though, it is not used in any calculations.

12) In the excel file, read the “Ending battery voltage, V”, and record it on your data sheet. The

voltage measurements are not used in any calculations, but they simply test that all the


13) On your data sheet, calculate the “Total coulombs of charge passed through electrodes”,

using C = Ah x 3600. Record the coulombs, C, on your data sheet.

14) On your data sheet, calculate the “Moles of electrons passed through electrodes”, mole e– = C/F,

F is the Faraday constant, F = 96485 C per mol of e–. Record moles of electrons on your data

sheet.

15) On your data sheet, convert mole of e– to mole of metal. For Cu2+ , the reduction reaction

equation is Cu2+(aq) + 2 e– → Cu(s), so the mole ratio between Cu2+ ion and e– is 1:2. Thus,

moles of Cu2+ = (mole of e– ) / 2 . Record the moles of Cu2+ on your data sheet.

16) On your data sheet, multiply the moles of Cu from step-15 by the molar mass of Cu, to get the

“Grams of copper that could have been electroplated”

17) On your data sheet, re-write the grams from step-3 as “Grams of copper available for

electroplating (from line 4):”

18) In the excel file, read the “Actual grams of metal (copper or silver) recovered from solution (dry

wt.)”, and record it on your data sheet. This is the actual yield.

19) On your data sheet, compare the “Grams of copper that could have been electroplated” with the

“Grams of copper available for electroplating”, to find out the “What limited the amount of

copper condensed? (circle one)”. If the “Grams of copper that could have been electroplated” is

more than the “Grams of copper available for electroplating”, then “available copper” is the

limiting amount. Oppositely, the “electric current” is the limit amount. Circle the limit amount,

either “available copper” or “electric current”. The limit amount is the theoretical yield.

20) On your data sheet, calculate the percent yield.

Percent Yield (Actual Yield/Theoretical Yield)100

Actual yield is from step-18. Theoretical yield is from step-19

For solution #2: 200 mL of 0.175 M copper acetate with three AAA batteries in

series.

Copper Experiment II

1) In the excel file, add 200 mL of 0.175 M copper acetate, Cu(C2H3O2)2(aq), in the 250 mL

“beaker”, by enter 200 in the “ ” spot. Record the volume of the solution (200 mL) and the

molarity of the solution (0.175) on your data sheet.


of Cu.

3) On the data sheet, multiply the moles of Cu by the molar mass of Cu, to get the “Grams of

copper available for electroplating”


on “three AAA” in the “ ” spot.



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metal. In real experiment, the formation of copper metal occurs slowly, shiny copper crystals

will form over the course of 36 to 48 hours.









12) In the excel file, read the “Ending battery voltage, V” from the excel file and record it on your

data sheet. The voltage measurements are not used in any calculations, but they simply test that

all the connections are good and the proper voltage is reaching the electrodes.





sheet.

15) On your data sheet, convert mole of e– to mole of metal. For Cu2+ , the reduction reaction

equation is Cu2+(aq) + 2 e– → Cu(s), so the mole ratio between Cu2+ ion and e– is 1:2. Thus,

moles of Cu2+ = (mole of e– ) / 2 . Record the moles of Cu2+ on your data sheet.

16) On your data sheet, multiply the moles of Cu from step-15 by the molar mass of Cu, to get the

“Grams of copper that could have been electroplated”

17) On your data sheet, re-write the grams from step-3 as “Grams of copper available for



wt.)” , and record it on your data sheet. This is the actual yield.

19) On your data sheet, compare the “Grams of copper that could have been electroplated” with the

“Grams of copper available for electroplating”, to find out the “What limited the amount of

copper condensed? (circle one)”. If the “Grams of copper that could have been electroplated” is

more than the “Grams of copper available for electroplating”, then “available copper” is the

limiting amount. Oppositely, the “electric current” is the limit amount. Circle the limit amount,

either “available copper” or “electric current”. The limit amount is the theoretical yield.

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For solution #3: 200 mL of 0.05 M silver nitrate with one AAA battery.

Silver Experiment

1) In the excel file, add 200 mL of 0.05 M silver nitrate, AgNO3(aq), in the 250 mL “beaker”, by

enter 200 in the “ ” spot. Record the volume of the solution (200 mL), and the molarity of

the solution (0.05 M) on your data sheet.


of Ag.

3) On the data sheet, multiply the moles of Ag by the molar mass of Ag, to get the “Grams of silver

available for electroplating”


on “one AAA” in the “ ” spot.



















metal. In real experiment, the formation of silver metal occurs slowly, shiny silver crystals will

form over the course of 36 to 48 hours.









12) In the excel file, read the “Ending battery voltage, V” from the excel file and record it on your

data sheet. The voltage measurements are not used in any calculations, but they simply test that

all the connections are good and the proper voltage is reaching the electrodes.

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sheet.

15) On your data sheet, convert mole of e– to mole of metal. For Ag+ , the reduction reaction

equation is Ag+(aq) + e– → Ag(s), so the mole ratio between Ag+ ion and e– is 1:1. Thus,

moles of Ag+ = (mole of e– ) . Record the moles of Ag+ on your data sheet.

16) On your data sheet, multiply the moles of Ag from step-15 by the molar mass of Ag, to get the

“Grams of silver that could have been electroplated”

17) On your data sheet, re-write the grams from step-3 as “Grams of silver available for



wt.)”, and record it on your data sheet. This is the actual yield.

19) On your data sheet, compare the “Grams of silver that could have been electroplated” with the

“Grams of silver available for electroplating”, to find out the “What limited the amount of silver

condensed? (circle one)”. If the “Grams of silver that could have been electroplated” is more

than the “Grams of silver available for electroplating”, then “available silver” is the limiting

amount. Oppositely, the “electric current” is the limit amount. Circle the limit amount, either

“available silver” or “electric current”. The limit amount is the theoretical yield.




Part B: Home Experiment

Copper coating a quarter

This experiment needs to be conducted by each individual student at home so you can witness

the reaction as it proceeds. This experiment demonstrates a different method of electroplating. In the

laboratory experiment, the metal ions were already in solution due to a salt solution. However,

electroplating is rarely conducted in this fashion. Most often, the source of the metal is an anode made

from the pure metal that you want to coat the cathode with. Therefore, metal is dissolved into solution

at the anode and condensed onto the cathode (connected to the “−” side of the battery). This basically

transfers the metal from one electrode to the other. The solution generally contains some salts to help

with the electricity conduction, but the salts that are present often are just spectator ions. Acid is

sometimes added to control the form of metal condensed and to help dissolve the metal from the anode

(the end connected to the + end of the battery).

Materials:

1 copper penny (pre 1980 are the best since they were still pure copper)

1 quarter

1 water container (plastic cup, jar, mug, etc.) that is NOT metal or conductive. I would not use anything

that you care too much about. Do not use a paper or waxed cup.

1 battery (AAA or AA). I would use rechargeable if available to avoid wasting batteries. You can use

other batteries types (camera, power tools, etc.) so long as the voltage is greater than 1.3V.

Tape

Two short lengths of wire with bare ends

Two binder clips with the paint scraped off of the contact point.

Deionized or distilled water (Phoenix tap water works, but it does not look as good.)

2 teaspoons vinegar (which is dilute acetic acid)

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Battery+ −

wire

wire

e−

e−

e−

e−

+

e−

e−

+

+

+

+

++e−

+

Quarter coin

(to be plated with copper)

Penny

(ideally pre-1980)

+

e−

e−

e−

e−e−

e−

Plastic cup

Deionized water

+ vinegar

(acetic acid)

+

+

+

+

+

+

+ e−e−

e−

1.3 V

Battery+ −Battery+ −

wire

wire

e−

e−

e−

e−

+

e−

e−

+

+

+

+

++e−

+

Quarter coin

(to be plated with copper)

Penny

(ideally pre-1980)

+

e−

e−

e−

e−e−

e−

Plastic cup

Deionized water

+ vinegar

(acetic acid)

+

+

+

+

+

+

+ e−e−

e−

1.3 V

Procedure:

1) A demonstration setup will be in the lab for you to look at. It is very simple but a little hard to

explain, so look at the demonstration setup (and the figure above) to get a good idea of how

everything is set up. I strongly suggest that you start this experiment early in case you need to

repeat it due to bad batteries or poor electrical contacts.

2) Get 2 binder clips and scrape the paint off of the contact point (where the clip grips) to expose

the bare metal to allow for a good electrical contact.

3) Wrap the bare end of a wire around the handle of one of the clips. Connect the second length of

wire to the second clip.

4) Clip the clips to the water container (hereafter “cup”). Position the clips on the opposite side of

the cup. Also make sure that they are about at the same level so that they will both be out of

water when the coins are clipped into place and the water is added.

5) Put a penny in one clip and put a quarter in the other clip. Make sure that the tops or the coins

are about the same level in the cup so that they will both be mostly submerged when the water is

added to the cup.

6) Add deionized or distilled water to the cup so that both coins are mostly submerged but the

binder clips and wires must remain dry and out of the water. Phoenix tap water will also work,

but the quarter will not look as good. It will be more of a blackish color rather than a copper

color. We will not take off any points if you do not use DI water. You can get DI water in the

lab, so I suggest that you bring an empty water bottle to fill with DI water.

7) Add 2 teaspoons vinegar (which is acetic acid) to the water. The measurement does not need to

be very accurate. The vinegar helps to increase the conductivity of the water and make sure the

solution is slightly acidic so the copper deposit to look better.

8) Connect the wire from the penny clip to the + end of the battery and the wire from the quarter

clip to the “−” end of the battery. Securely tape the wire to the ends of the batteries. Getting a

good connection at the battery seems to be one of the most problematic aspects of this

experiment. Coiling the wire helps to make the contact because the thicker wire coil is easier to

tape down.

9) Wait 24 hours. There will be no visible activity for the first several hours. After 24 hours,

copper deposits should be clearly visible. You can continue the electroplating for another 24

hours to continue adding more copper to the quarter.

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a. If you used DI or distilled water and a 1.3V battery as suggested, then deposit will look

coppery like the quarter on the left in the below picture.

b. If you used Phoenix tap water or a higher battery voltage, then the copper deposits will

look dark rusty brown like the quarter on the right below.

10) Remove the quarter from the clip and rinse it with water. Don’t rub the quarter too much since

the copper may be loosely on the quarter and it will rub off easily. This is particularly true for

the dark brown plated quarters.

11) Answer the questions on the work sheet.

12) TAPE YOUR COPPER-COATED QUARTER TO YOUR WORKSHEET. WE WILL GRADE

YOU ON THE COIN TURNED IN SINCE THIS PROVES YOU CONDUCTED THE

EXPERIMENT. Take picture of it together with your lab report. No quarter, no credit for

this experiment! Attempts to fake the experiment (e.g. dipping the quarter in copper colored

paint) is considered academic dishonesty and will result in an automatic zero for the whole

laboratory exercise and other potential disciplinary action.

Part C: Computer simulator of Electroplating.

1. Open the simulator using the link (preferred using Google Crown Web browser):

https://media.pearsoncmg.com/bc/bc_0media_chem/chem_sim/html5/Electro/Electro.php Acknowledge “Pearson Education”. The opening page of the simulator appears as:

https://media.pearsoncmg.com/bc/bc_0media_chem/chem_sim/html5/Electro/Electro.php

11

2. Click on the “Experiment” tag, then click on the “Run experiment” tag, the simulator appears as:

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The + and – signs on the electrodes refer to the terminals of the DC battery. So, the + terminal is

connected to the anode of the electrolysis cell, where the oxidation reaction happens. The – terminal is

connected to the cathode of the electrolysis cell, where the reduction reaction happens.

3. Click on the tag to choose “Aluminum (Al)” as the metal on the “+”

electrode. Record it on your data sheet

4. Click on the tag to choose “Aluminum (Al)” as the metal on the “–” electrode.

Record it on your data sheet

5. Click on the tag to choose Al(NO3)3(aq) as the solution. Record it on your data sheet 6. Record the mass of the two metals on you data sheet. Those are the initial masses.

7. Set your DC current by sliding the bar on the tag. You should use a current between

10.00 to 30.00 AMPS. Record it on your data sheet 8. Set your time duration for the electrolysis. You should use a time between 10.00 to 30.00

minutes. Record it on your data sheet. Those times are not real time. For example, 20 minutes

in the simulator is about 2 minute in real clock. But in your calculation, you use the numbers you

chosen on the simulator as real minutes 9. Click on the “Off ON” tag to run the electroplating, with the default voltage set at “+ 6.00

VOLTS”.

10. When the electrolysis finish, read the final mass of the metals on the electrodes, and record on

your data sheet. Take a screen shot of the simulator when finished run. Attach the picture

to your lab report.

11. For the cathode metal (the electrode with a “–” sign), the final mass is larger than the initial

mass. Subtract the final mass by the initial mass to get the grams of the metal (Al) deposited on

the cathode. Record this as the “Actual yield of the metal reduced”.

12. On your data sheet, write the oxidation reaction happened on the anode.

13. On your data sheet, write the reduction reaction happened on the cathode.

14. On your data sheet, use the AMPS from step-7 and the time duration from step-8, to calculate

the Coulombs (C) passed through the electrodes.

C = AMPS x minutes x 60

15. On your data sheet, convert the coulombs to moles of electrons

Mole of electrons = C / F

Use the C you got from step-15, and the Faraday constant, F = 96485 C per mole of electron

16. On your data sheet, use the mole ratio between the Al(s) and the e– to convert moles of electrons

to moles of Al(s).

17. On your data sheet, convert the moles of Al(s) to grams of Al(s), using the molar mass of Al =

26.98 g/mol. This is the theoretical yield.

18. On your data sheet, calculate the percent yield.

Percent yield = actual yield/theoretical yield x 100

Actual yield is from step-11, theoretical yield is from step-17.

If you need to do the experiment again, click on the reset tag in the simulator.

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Data Sheet for Part A Copper Experiment I

Use 200 ml of 0.080 M copper(II) acetate solution with 2 AA batteries

Volume of copper acetate solution: ______________ mL

Molarity of copper acetate: ______________ M

Moles of copper available for electroplating: ______________ moles Cu

Grams of copper available for electroplating: ______________ g Cu

Reported battery charge ______________ mA·h = ______________ A·h

The batteries you are using can give only 85% of their charge above the voltage needed to drive this

electroplating. Therefore, the useable battery charge is the reported (A·h)(0.85)

Useable battery charge ______________ A·h (= reported charge(0.85))

Start time of electroplating ______________ (time and date)

Starting voltage across the electrodes ______________ V

End time of electroplating ______________ (time and date)

Ending voltage across the electrodes ______________ V

(with the batteries still attached)

Total coulombs of charge passed through electrodes: ______________ C

This should equal (useable battery charge in A·h)(60min/1 h)(60sec/1min).

Moles of electrons passed through electrodes: ______________ mol e−(=C/F)

Moles of copper that could have been electroplated: ______________ mol Cu

Grams of copper that could have been electroplated: ______________ g Cu

Grams of copper available for electroplating (from line 4): ______________ g Cu

Actual grams of copper recovered from solution (dry wt.) ______________ g Cu

What limited the amount of copper condensed? (circle one)

available copper OR electric current

Product Yield (Actual Yield/Theoretical Yield)100 ______________ %

This equals the actual g copper recovered / the theoretical grams recovered by the limiting reagent

(which is either electric current of amount of copper available).

______ of 3 pts

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Data Sheet for Part A Copper Experiment II Use 200 ml of 0.175 M copper(II) acetate solution with 3 AAA batteries

Volume of copper acetate solution: ______________ mL

Molarity of copper acetate: ______________ M

Moles of copper available for electroplating: ______________ moles Cu

Grams of copper available for electroplating: ______________ g Cu













Moles of copper that could have been electroplated: ______________ mol Cu

Grams of copper that could have been electroplated: ______________ g Cu

Grams of copper available for electroplating (from line 4): ______________ g Cu

Actual grams of copper recovered from solution (dry wt.) ______________ g Cu

What limited the amount of copper condensed? (circle one)

available copper OR electric current


This equals the actual g copper recovered / the theoretical grams recovered by the limiting reagent

(which is either electric current of amount of copper available).

______ of 3 pts

15

Data Sheet for Part A Silver Experiment: Use 200 ml of 0.05 M silver nitrate solution with 1 AAA battery

Volume of silver nitrate: ______________ mL

Molarity of silver nitrate: ______________ M

Moles of silver available for electroplating: ______________ moles Ag

Grams of silver available for electroplating: ______________ g Ag













Moles of silver that could have been condensed: ______________ mol Ag

Grams of silver that could have been condensed: ______________ g Ag

Grams of silver available for electroplating (from line 4): ______________ g Ag

Actual grams of silver recovered from solution (dry wt.) ______________ g Ag

Which limited the amount of silver condensed? (circle one)

available silver OR electric current


This equals the actual g silver recovered / the theoretical grams recovered by the limiting reagent (which

is either electric current of amount of silver available)

______ of 3 pts

16

Part B HOME EXPERIMENT

How many batteries did you use and what voltage were they? ______________

What was the final color of the penny in solution? ______________

What was the final color of the electroplated quarter in solution? ______________

Write the redox reactions that are occurring at the anode and cathode:

1st half redox reaction at the anode: __________________________________________

2nd half redox reaction at the cathode: __________________________________________

TAPE YOUR ELECTROPLATED QUARTER FROM THE EXPERIMENT HERE

(no quarter, no credit!)

______ of 5 pts

17

Part C Computer Simulator Experiment

Identity of metal used on the “+” electrode: ______________

Initial mass of metal used on the “+” electrode, (grams): ______________

Identify of metal used on the “–” electrode: ______________

Initial mass of metal used on the “–” electrode, (grams): ______________

Solution used: ______________

Current used (AMPS): ______________

Duration of time of electroplating, minutes ______________

Final mass of metal on the “+” electrode, (grams): ______________

Final mass of metal on the “–” electrode, (grams): ______________

Actual yield of the metal reduced ______________

Oxidation reaction happened on the anode ______________

Reduction reaction happened on the cathode ______________

Coulombs (C) passed through the electrodes ______________

C = AMPS x minutes x 60

Moles of electrons (Mole of electrons = C / F) ______________

Moles of Al(s) deposited on the cathode ______________

Grams of Al(s) deposited on the cathode (theoretical yield) ______________

Percent yield of Al(s) deposited on the cathode. ______________

Attach the picture of the screen shot when the simulator finished run.

______ of 4 pts

18

Post Lab Questions: 1) Suppose a similar electroplating experiment as the ones conducted in this lab is conducted to purify

copper metal from 200 mL of a 1.5 M copper sulfate solution (CuSO4). You used a battery whose

energy capacity is listed as 2500 mAh (milli-ampere hours) and you can assume that the entire battery

charge is useable for electroplating. You completely drain the battery during the experiment. How

much copper (g) was pulled out of solution? Note that copper ions have a +2 charge and thus need two

electrons to make it a neutral metal atom. (Show your work for full credit)

2) Gold is frequently purified by electroplating. Typically the gold is dissolved in a cyanide solution

(which is fairly selective for gold) and then it is pulled out of solution by electroplating. If the

electroplating apparatus used a current of 3.5 amperes, then how long will it take (h) to purify one ounce

of gold from a gold(III) cyanide solution? (Show your work for full credit)

3) You conduct an experiment where you use a single battery with a charge capacity of 900 mAh to

precipitate an unknown metal from a salt solution that was leaking from an old battery. You use a vast

excess of the salt, so the amount of metal condensed is limited by the battery. You obtain 3.368 g of the

metal. You can assume that the metal was the only substance being reduced in the electrochemical cell

and 100% of the listed battery charge is useable. If you were able to determine that the metal ion was a

+2 ion, then which metal is it? This metal was used in batteries for a long time (even in “alkaline”

batteries) due to its good electrochemical properties, but it is now being phased out. Why? (Show your

work for full credit)

4) You conduct another experiment on silver condensation. You add 8.0 g of silver nitrate to 200 mL of

pure water. You attach a single 850 mAh AAA battery (at 1.3 V) to the electrochemical cell. You can

assume that the entire battery charge is useable for electroplating. You get 3.24 g of silver metal at the

end of the experiment.

a) Was your silver recovery limited by the battery capacity or the available silver in solution?

b) What was your product yield (%) using the limiting resource?

______ of 4 pts

Total ______ of 25 pts

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Electrochemistry Laboratory Electroplating of Metals By Thomas …jhwang43/CHM116Labs... · 2020. 8. 8. · Electrochemistry Laboratory ... The electroplating that you are conducting