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Electrochemistry describes the interaction between electrical energy and chemical processes. Electricity continues to intrigue us, as it has since people first observed the sky shattered by bolts of lightning. Electrochemistry is of great practical value to contemporary living. Consider the number of batteries used for powering the many portable items of pleasure and need – everything from cassette recorders to hearing aids, from calculators to pacemakers. Pure metals are produced from natural ores, inorganic and organic compounds are synthesized, metal surfaces are plated with other metals or coated with paint to enhance their value and utility – all through electrochemistry.

Electricity is a moving stream of electrical charges. This flow, or electric current, can occur as electrons moving through a wire or as ions flowing through an aqueous solution. If the electrons lost and gained in a spontaneous reaction can flow through a wire on their pathway from the substance oxidized to the substance reduced, the energy of the reaction is released as electrical energy. Conversely, a non-spontaneous redox reaction can be driven forward by the introduction into the system of electrical energy from another source. Any device in which either process can occur is called an electrochemical cell.

There are two types of electrochemical cells. The first type generates electrical energy from a spontaneous redox reaction. These are called voltaic or galvanic cells, common household batteries are classic examples. An Italian physicist, Allesandro Volta in 1800 explained that electricity is generated by the connection of two dissimilar metals separated by any moist body (not necessarily organic). A simple voltaic cell, similar to that made by Volta, can be assembled using twelve pennies and twelve nickels (construct a column of alternating pennies and nickels with each coin separated by disk-size pieces of wet filter paper soaked in salt water).

In the second type of electrochemical cell, called an electrolytic cell, a non-spontaneous redox reaction is caused by the addition of electrical energy from a direct current source such as a generator or a storage battery. The process of generating a non-spontaneous redox reaction by means of electrical energy is called electrolysis.

Electrolysis can be used for purifying a metal through the electrolytic dissolution of an impure anode and the subsequent re-crystallization of the pure metal on the cathode. The impurities are left behind in solution. Copper is refined commercially by this electrolytic technique.

Electrolysis is often used for electroplating a metal to another material acting as the cathode. The other material must also be electrically conducting. Non-conducting materials, such as leaves, can also be plated by first being painted with a metallic conductive paint. Silver plating can be done with a silver anode and the object to be plated as the cathode.

Electrolytic reduction (cathodic reduction) has developed into a useful technique for the restoration of artifacts such as corroded nails and encrusted silver. In the case of silver, the degradation is usually due to the surface formation of insoluble (black) silver sulfide ( Ag2SAg2S size 12{"Ag" rSub { size 8{2} } S} {}). The artifact (a silver coin, for example) is attached to the negative electrode of the electrolysis cell. The Ag+Ag+ size 12{"Ag" rSup { size 8{+{}} } } {} ions of the silver sulfide pick up electrons and are converted back to metallic silver:

 The sulfide ions are swept away by the water and the surface of the object is restored.

In this experiment, you will electroplate copper quantitatively to a copper cathode (the anode is also composed of copper). The current is measured over an interval of approximately one hour so that the amount of charge passing through the cell is known. The molar mass of copper is calculated from its equivalent mass using Faraday’s second law. In the second part of the experiment, you will use turn copper into “gold”!

In the 1830s, Michael Faraday published his experiments using the recently discovered voltaic column to decompose substances through the use of electric current. Electrolysis is an oxidation-reduction process involving a conversion of electrical energy to chemical energy. The electrolytic cell is a galvanic cell operating in reverse. The automobile battery is acts as a collection of galvanic cells when delivering electric current, but acts as a collection of electrolytic cells when being recharged.

Faraday first described the quantitative relationships between the amount of electric charge (number of electrons) that has passed through an electrolytic cell and the amount of materials that have formed at the electrodes. These are summarized as Faraday’s Laws of Electrolysis: 

Through exhaustive experimentation, the charge of a single electron has been determined to be 1.602×10−191.602×10−19 size 12{1 "." "602" times "10" rSup { size 8{ - "19"} } } {}coulombs (C). The coulomb charge unit – defined as useful for much larger charged objects – is inconvenient for expressing such a small charge, so other electrical charge units are commonly used. One mole of electrons has a total charge calculated to be 96,485 C; this quality is defined as faraday (F):

  1 F=96,485 C/mol e−1 F=96,485 C/mol e− size 12{"1 F"="96,485 C/mol e" rSup { size 8{ - {}} } } {}

 Electric currents (I) are measured in amperes (A), amps for short, and defined in terms

I = Q/t

1 A = 1 C/s

 For example, a constant current of .600 A (milliamperes) over a period of 2.00×1022.00×102 size 12{2 "." "00" times "10" rSup { size 8{2} } } {} seconds represents

a movement of 120 coulombs. The number of moles of electrons (n) transported during the time interval is

Figure 1

Time intervals measured in minutes and hours must be converted to seconds in such calculations.

Experimental Procedure

Figure 2

CAUTION WEAR EYE PROTECTION!

CAUTION: The 6 M nitric acid used in the next step will burn and stain the skin as well as damage clothing. In case of skin or clothing contact, wash the area immediately with large amounts of water.

1. Obtain a piece of copper mesh (about 5 cm×8 cm5 cm×8 cm size 12{"5 cm" times "8 cm"} {}) and remove any loose pieces of copper. Clean and rinse. Place the copper mesh on a watch glass in the drying oven. Be careful not to touch the cleaned surfaces. This is the cathode.
2. Obtain 1 piece of copper foil (about 2 cm×8 cm2 cm×8 cm size 12{"2 cm" times "8 cm"} {}). Holding the foil with tweezers or tongs, dip it into 6 M nitric acid several times until its surface is bright and shiny. Do not allow tweezers or tongs to touch the acid solution. Rinse the foil in de-ionized water and set it aside. This is the anode. Set the nitric acid aside to use in the electroplating exercise.
3. Add 350 mL 1.0 M KNO3KNO3 size 12{"KNO" rSub { size 8{3} } } {} solution to a 400 mL beaker.

 CAUTION: The copper sulfate used in the next step is toxic. Avoid skin contact!

1. To this solution, add about 5 mL of 1 M H2SO4H2SO4 size 12{H rSub { size 8{2} } "SO" rSub { size 8{4} } } {} and 10 g of CuSO4⋅5H2OCuSO4⋅5H2O size 12{"CuSO" rSub { size 8{4} } cdot "5H" rSub { size 8{2} } O} {}. Stir until the copper sulfate pentahydrate is fully dissolved.
2. Assemble the apparatus shown in Figure 1, but leave the copper mesh electrode in the oven. Add a magnetic stirring bar to the beaker. If necessary, add additional 1.0 M KNO3KNO3 size 12{"KNO" rSub { size 8{3} } } {} to bring solution level in the beaker within 2 cm of the rim. You will either measure the electricity directly with an ammeter in series with the electrolytic cell or you will measure the current indirectly by measuring the voltage across a resistor of known value (about 10 ohms).
3. Remove the copper mesh electrode from the oven, let it cool, and determine its mass to the nearest milligram.
4. Attach the copper mesh electrode to the negative terminal of your power supply using an alligator clip. Turn on the magnetic stirrer.
5. Turn on the low voltage power supply and adjust the current until about 140 mA are flowing through the cell. Record the time and current.
6. Record the time and current every five minutes for an hour.
7. After the last reading, gently remove the cathode from the solution while leaving it attached to the power supply. After the copper mesh has cleared the solution, remove the wire and turn off the power supply.
8. Gently dip-rinse the copper mesh electrode several times in a beaker of deionized water, and place it on a watch glass in the drying oven. Be careful not to rinse the mesh too harshly because you don’t want any copper that has deposited on it to come off.
9. When dry, remove the electrode from the oven and let it cool. Reweigh the mesh electrode.
10. Remove the magnetic stirring bar from your beaker and dispose of the solution in the sink.

Figure 3

Figure 2

The bottom of the beaker is half-covered with 30-mesh zinc metal. (A) Two pennies are lying on top of the 30-mesh zinc.

(B) Two pennies are lying on the bottom of the beaker but not in contact with the 30-mesh zinc. (C) A penny soldered to copper wire is immersed in solution, the other end of the copper wire being soldered to a strip of zinc metal in contact with 30-mesh zinc on the bottom of the beaker. (D) A penny soldered to copper wire is immersed in solution. The solution in the beaker is 1 M NaOH.

 Review the General Soldering Instructions in Part I. Place the freshly tinned tip of the penny next to the wire angling the iron to get good thermal contact. Don't dab at the joint with the tip of the iron while soldering.

 Example of the calculation of Molar Mass:

1. Graph the electric current (in amps) on the y-axis against time (in seconds) on the x-axis. The total charge that passed through the electrolysis cell is given by the area beneath this curve. If the current is constant, this area is:

Q = area = I × t Q = area = I × t size 12{Q="area"=I times t} {}

(4)

Calculate this charge in coulombs.

1. Convert the coulombs of charge to mol electrons:  

N =

Figure 4

  The equation for the reduction half-reaction responsible for the plating at the cathode is

Cu 2 + ( aq ) + 2e − → Cu ( s ) Cu 2 + ( aq ) + 2e − → Cu ( s ) size 12{"Cu" rSup { size 8{2+{}} } \( "aq" \) +"2e" rSup { size 8{ - {}} } rightarrow "Cu" \( s \) } {}

(5)

 Use the mol ratios of the preceding balanced equation to calculate the number of moles of Cu plated out:

Figure 5

n(Cu)=n/2n(Cu)=n/2 size 12{n \( "Cu" \) ="n/2"} {} 

1. Use the initial and final masses of the copper mesh electrode to calculate the mass of copper plated out:

m(Cu) = m(final) – m(initial)

1. Calculate the molar mass (M) of copper:  

M =

Figure 6