Electroplating is an electrochemical process for depositing a thin layer of metal on a metallic base. Objects can be electroplated in order to avoid corrosion, to obtain a hard surface or attractive finish, for purification purposes of metals (as in the electrorefining of copper), to separate metals for analysis, or, as in electrotyping, to reproduce a form from a mold. Cadmium, chromium, copper, gold, nickel, silver, and tin are the metals used most often. Typical consumer products included silver-plated tableware, chrome-plated automobile accessories, and jewlery.

In the process of electroplating, the object to be coated is placed in a bath of a salt of the coating metal, and is connected to the negative terminal of an external source of electricity. Another conductor, often composed of the coating metal, is connected to the positive terminal of the electric source. A steady direct current of low voltage, is required for the process. When the current is passed through the solution, atoms of the plating metal deposit out of the solution onto the cathode, the negative electrode. These atoms are replaced in the bath by atoms from the anode (positive electrode), if it is composed of the same metal, as with copper and silver. Otherwise they are replaced by periodic additions of the salt to the bath, as with gold and chromium. In either case an equilibrium is kept until the object is plated. Nonconducting materials may be plated by first being covered with a conducting material such as graphite.

A clean object is important to the electroplating process. The object must be cleaned thoroughly by dipping it into an acid or caustic solution. To eliminate irregularity in the depth of the plate, and to ensure that the grain at the surface of the plate is of good quality and polishable, the current density (amperes per square foot of cathode surface) and temperature must be carefully controlled. Colloids are often added to the bath to improve the surface uniformity.


Topic reworded from an entry on http://encarta.msn.com

An Experiment using Electroplating

Introduction

Problem

The outcome of an electrolysis reaction with two metals in a copper sulfate solution is unknown.

Hypothesis

The electrolysis reaction can be determined through visual observation of the electroplating that occurs. One metal will be oxidized, while the other will be reduced in a redox reaction.

Prediction

When an electrical current is applied, the copper in the copper sulfate solution will become copper metal at the anode, or negative part of the reaction.

Explanation Using Chemical Principles

In a metallic bond, and in a solution containing ionic compounds, electrons are able to flow freely between molecules. This is what allows for electric current, a flow of electrons in a circuit.

The electric current applied to an electrolytic cell causes a redox reaction. At the anode, or positive end, the metal is oxidized; electrons are stripped away. At the cathode, or negative end, the metal is reduced; electrons are added. In an electrolysis reaction, this causes the metal in an aqueous solution (in this case, copper in copper sulfate) to be deposited on the cathode.

Method

Materials

Step-by-step Method

  1. Pour CuSO4 solution into the 100 mL beaker
  2. Attach one end of each of the electrical wires to one end of the battery (one should be attached to the positive end, and one should be attached to the negative end)
  3. Attach each of the wires to any single sample of metal (one wire should be attached to one sample, and the other wire should be attached to another sample)
  4. Holding the wires, place both of the samples of metal into the solution so that the metals are submerged or partially submerged, but the alligator clips or wires are not touching the solution
  5. Take note of which metal is attached to the negative and positive ends of the battery
  6. Record the observable results of the reaction: if any metal was deposited on the anode or cathode, and what metal it appears to be. (Also record what metals were used for the anode and cathode.)

Sample Calculations

The half-reaction of the copper plating is: Cu+2 + 2e- --> Cu

This shows that 2 moles of electrons are necessary per mole of copper. The gram equivalent weight] is:

GEW = atomic mass of Cu / 2 = 63.5 / 2 = 31.75 grams / equivalent.

Data

Tables

Note: because everything2 doesn't allow HTML tables, I'll have to come up with something creative here. Watch me.

Table 1. Observable Results of Electrolysis Reactions in CuSO4 Solution

  • Anode: Carbon
    Cathode: Iron
    Results: Copper deposited on cathode
  • Anode: Carbon
    Cathode: Zinc
    Results: Copper deposited on cathode
  • Anode: Zinc
    Cathode: Carbon
    Results: Copper deposited on cathode
  • Anode: Iron
    Cathode: Carbon
    Results: Copper deposited on cathode
  • Anode: Carbon
    Cathode: Carbon
    Results: Copper deposited on cathode
  • Anode: Zinc
    Cathode: Iron
    Results: Copper deposited on cathode
  • Anode: Iron
    Cathode: Zinc
    Results: Copper deposited on cathode
  • Anode: Tin
    Cathode: Tin
    Results: Copper deposited on cathode
  • Anode: Tin
    Cathode: Aluminum
    Results: Copper deposited on cathode
  • Anode: Aluminum
    Cathode: Tin
    Results: Copper deposited on cathode
  • Anode: Tin
    Cathode: Iron
    Results: Copper deposited on cathode
  • Anode: Iron
    Cathode: Tin
    Results: Copper deposited on cathode
  • Anode: Zinc
    Cathode: Tin
    Results: Copper deposited on cathode
  • Anode: Aluminum
    Cathode: Iron
    Results: Copper deposited on cathode
  • Anode: Iron
    Cathode: Aluminum
    Results: Copper deposited on cathode

The same result occurred in each reaction: copper metal was deposited on the cathode.

Discussion

Analysis

The expected results were that copper would plate the metal at the cathode. This is what happened in every trial. The data obtained is therefore correct.

Because the data is not quantitative, no error calculation is possible. However, this is entirely unnecessary, as there is no measurable or immeasurable error in the data.

Evaluation

The prediction was correct. This is due to the reduction potential of the reaction. According to the table of standard reduction potentials, the reduction potential of copper is relatively high, at 0.34 volts . When this reduction potential is higher than that of other possible half-reactions, this is the reaction that occurs. If other metals had been used that had a higher reduction potential, such as silver or gold, those would have instead reacted.

The half reaction of the reduction of copper is Cu+2 + 2e- --> Cu.

Uncertainties Elucidated

In the reaction involving carbon as the anode and zinc as the cathode, a substance other than copper appeared to be coated on the zinc. According to the reduction potential table, this is not possible; only a few metals have a higher reduction potential than copper. None were involved in any of these reactions. Apparently, the metal was in fact copper, and was misidentified at the time of reaction.

The same phenomenon occurred in the reaction involving aluminum as the anode and tin as the cathode, although the same holds true. The metal must in fact be copper.

The copper differed in appearance when coating different metals, due to the way that it was plated on, and the surface of the metals. It appeared redder and lighter on the iron than the average, and appeared darker and more grainy on the zinc and tin.

Reference to Other Reactions/Phenomena/Results

Other similar reactions and their respective calculations can be found in all of the referenced materials.

Conclusion

The evidence and refined data supported the original hypothesis and prediction; the experiment was a success. It can be determined from these results that it is possible to plate copper metal on most metals using the electroplating procedure. It can also be inferred that the following is true:

Abstract

The experiment is to determine the reactions involved in electroplating. The primary half-reaction present in the electrolytic cell, regardless of the metals used, is Cu+2 + 2e- --> Cu. Because this half-reaction has a very high reduction potential, it is more likely to occur than most half-reactions. It is thus very easy to copperplate nearly any metal, as can be seen from the results of the experiment. The same technique can be applied to other metals such as gold or silver, but the reduction potential must be higher than the other half-reactions involved in the redox reaction. The metal in the solution will be plated on the cathode, because it needs to gain electrons to become a metal from an ion.

Bibliography

  1. Electrochemistry and Electric Current
    http://pc65.frontier.osrhe.edu/hs/science/cfaraday.htm
  2. Electrolytic Cells and Faraday's Law of Electrolysis
    http://members.aol.com/logan20/faraday.html
  3. Electrolysis
    http://www.science.uwaterloo.ca/~cchieh/cact/c123/eltlysis.html

E*lec"tro*pla`ting (?), n.

The art or process of depositing a coating (commonly) of silver, gold, or nickel on an inferior metal, by means of electricity.

 

© Webster 1913.

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