Electrochemistry in Electroplating

This Exemplar will teach the following concept(s) from the ACS Examinations Institute General Chemistry ACCM:

V. D. 2. a. Oxidation–reduction reactions are important to chemistry and are sometimes referred to as “redox” reactions.

V. D. 2. b. Being able to identify an oxidation–reduction reaction often involves assigning oxidation numbers.

V. D. 2. c. The use of oxidation–reduction reactions as a category includes the concepts of oxidizing and reducing agents.

VI. C. 3. a. Oxidation–reduction reactions can be considered in two halves, and can be physically separated to form electrochemical systems.

VI. C. 3. b. Electrical energy can be obtained from electrochemical systems operating as galvanic cells.

Electroplating Can be Seen Everywhere

If you have ever used a metal spoon, used a coin, or worn jewelry, then it is likely that you have come across electroplating.²  Electroplating is the process by which a thin layer of metal can be applied to an object. This is useful for decorative purposes or for protecting the material underneath the electroplated metal from corrosion.²

So, how is it even possible to get a smooth, even coating of metal on an object? To know the answer to this question, we must first talk about oxidation-reduction reactions and electrolysis. 

Oxidation-Reduction Reactions

Oxidation-reduction reactions (or redox reactions) are reactions which involve a transfer of electrons. In these reactions, we can assign atoms oxidation numbers to help us visualize the movement of electrons. Oxidation numbers can be determined based on the following principles: the most electronegative atom in a compound has a negative oxidation number equal in magnitude to it’s ionic charge, oxygens and hydrogens in compounds respectively have an oxidation number of -2 and +1, the sum of every atom’s oxidation number in a compound should equal the compound’s net charge, and pure elements have oxidation numbers of zero. Oxidized atoms are those whose oxidation numbers increase, and reduced atoms are those whose oxidation numbers decrease.

\[2Na + Cl_2 \rightarrow 2NaCl\]

In the example above, we see that Na is oxidized and Cl is reduced (Na loses an electron and Cl gains an electron).³ Redox reactions can also be split into half-reactions that show only the reaction’s oxidation or reduction component:

\[Cl_2 + 2e^- \rightarrow 2Cl-\] (reduction)

\[2Na \rightarrow 2Na^+ + 2e^-\] (oxidation)

Example \(\PageIndex{1}\): Oxidation of Iron

A rusty iron spoon will be electroplated with silver to prevent the spoon from further rusting. The process by which the spoon rusts is shown below:

\[Fe + O_2 \rightarrow Fe_2O_3\]

Which atom is reduced? Which atom is oxidized? What are the oxidation numbers of each atom before and after the reaction?

Solution

Both Ba and Cl₂ start off with oxidation numbers of 0 since they are elements. The O in Fe₂O₃ has an oxidation number of -2 (remember, "O" in compounds usually has an oxidation number of -2!). Therefore, Fe must have an oxidation number of +3 since Fe₂O₃ is neutral. Thus, O is reduced and Fe is oxidized.

 

Electrolysis and Electroplating

Examine the digram below:

Figure \(\PageIndex{1}\): Galvanic Cell involving Cd and Zu⁵

Here, we have two solution filled beakers: one with a solid plate of Cd and aqueous Cd²⁺ cation, and the other with a solid plate of Cu and Cu²⁺ ion. The electrons produced from the oxidation of Cd travel through a wire to the plate of Cu, which then causes the Cu²⁺ cations in solution to reduce to Cu. The two beakers are connected by a salt bridge containing cations and anions that can flow into each beaker, which ensures neutral charge in both beakers (the flow of electrons would stop if the beakers’ charge neutrality was not maintained due to the Law of Conservation of Charge). This apparatus is known as a galvanic cell, and is the general principle behind batteries.

However, what if we wanted to reverse the reaction? To do this, we require electrolysis: the use of an applied current to cause a nonspontaneous redox reaction to occur.

Figure \(\PageIndex{2}\): Electrolytic Cell involving Cd and Zu⁵

In this diagram, we insert an external power supply into the apparatus. The power supply forces the current to flow in the opposite direction, and the reactions are reversed.

We can take advantage of electrolysis to perform electroplating:

Figure \(\PageIndex{3}\): Electroplating copper on a metal object⁶

\[Cu \rightarrow Cu^{2+} + 2e^-\]

\[Cu^{2+} + 2e^- \rightarrow Cu\]

Here, an anode made of copper and a metal plate are submerged in a solution containing aqueous copper ions. The copper anode is attached to the positive terminal of the battery, which pulls electrons from the copper atoms. This oxidizes atoms in the anode, turning them into aqueous copper ions (top equation). The electrons produced from this process travel from the anode to the metal plate. The copper ions are attracted to these free-flowing electrons. Copper ions surrounding the metal plate are reduced to solid copper by these electrons (bottom equation), and deposit on the surface of the plate, thus coating the metal plate in copper. This process can be done with almost any metal, so long as it will stick to the surface. That means that metal parts on cars can be plated with chromium to avoid rusting, that spoons can be plated in silver so as to appear fancier, and that tools can be plated in extremely hard metals to improve their lifetime! Electroplating is all around us.

Example \(\PageIndex{2}\): Identifying Reactions

Examine the electroplating process below. Write the redox half reactions for the process.

Figure \(\PageIndex{4}\): Electroplating of silver on a spoon⁴

Solution

Ag⁺ is being reduced at the spoon (cathode) and Ag is being oxidized at the Ag plate (anode). The redox reactions are:

On the cathode side:
\[Ag^{+} + e^- \rightarrow Ag\]

On the anode side:

\[Ag \rightarrow Ag^{+} + e^-\]

Example \(\PageIndex{3}\): Calculating Time

A watch is being electroplated with Pt. If we want 100g of Pt on the watch, how many electrons are transferred through the wire during the electroplating process? If the current is 5 A, how long will it take for the 100g of Pt to be plated?

Solution

First, we should write the half-reactions:

\[Pt \rightarrow Pt^{2+} + 2e^-\]

\[Pt^{2+} + 2e^- \rightarrow Pt\]

Next, convert the 100g of Pt to moles:

\[100\ g \ \left(\frac{1\ mol}{195.08g}\right) \approx\ 0.513\ mol\ Pt\]

Using the mole ratio of Pt to e^-, we get:
\[0.513\ mol\ Pt\ \left(\frac{2\ mol\ e^-}{1\ mol\ Pt}\right)\left(\frac{6.02\textrm{x}10^{23}}{1\ mol\ e^-}\right) \approx 6.177\textrm{x}10^{23}\ e^-\]

Since 1 ampere is 6.28 x 10¹⁸ electrons per second, then 5 amperes is 3.14 x 10¹⁹ electrons per second. Thus, amount of time it takes to plate the watch is:

\[\left(\frac{1\ second}{3.14\textrm{x}10^{19}\ e^-}\right)\ \left(6.177\textrm{x}10^{23}\ e^-\right) \approx 19671.97\ seconds \approx 328\ minutes\]

Example \(\PageIndex{4}\): Comparing Electroplating Systems

A metal spoon is being electroplated with silver. A second metal spoon is being electroplated with copper. After 30 minutes, which spoon will have more moles of the desired plated metal on it's surface? Assume that the current in both electroplating systems is the same.

Solution

The oxidation and reduction half-reactions involved in silver electroplating are:

\[Ag \rightarrow Ag^+ + e^-\]

\[Ag^+ + e^- \rightarrow Ag\]

And the oxidation and reduction half-reactions involved in copper electroplating are:

\[Cu \rightarrow Cu^{2+} + 2e^-\]

\[Cu^{2+} + 2e^- \rightarrow Cu\]

The electroplating of silver requires the transfer of less electrons per atom than the electroplating of copper does. Thus, for a given current, more silver will be plated per second than copper. After 30 minutes, the spoon being plated with silver will have of its desired plated metal than the spoon being plated with copper.

 

References

1. Admin. What is electroplating - definition, working principle & its uses. https://byjus.com/chemistry/electroplating-process/. (accessed Nov 10, 2022). 

2. McMullen, J. 7 daily life examples of electroplating: Dorsetware. https://www.dorsetware.com/7-daily-l...t%20tarnishing. (accessed Nov 10, 2022). 

3. Zumdahl, S. S.; DeCoste, D. J. Chemical principles; Cengage Learning: Boston, MA, 2017. 

4. https://chem.libretexts.org/Bookshel...Electroplating 

5. https://chem.libretexts.org/Courses/Grand_Rapids_Community_College/CHM_110%3A_Chemistry_of_the_Modern_World/10%3A_Electrochemistry/10.2_Batteries_and_Electrolytic_Cells

6. https://en.wikipedia.org/wiki/Electroplating#/media/File:Copper_electroplating_principle_(multilingual).svg