DAT General Chemistry: Electroplating, Batteries, Fuel Cells, and Corrosion

DAT High-Yield Concept

Oxidation Always at Anode, Reduction Always at Cathode: This fundamental rule applies to both galvanic (spontaneous, $E^\circ > \pu{0 V}$ ) and electrolytic (nonspontaneous, $E^\circ < \pu{0 V}$ ) cells.
Electroplating: An electrolytic process using external electrical energy to deposit metal cations onto a conductive cathode.
Batteries vs. Fuel Cells: Batteries store a self-contained supply of reactants. Fuel cells require continuous external feeds of fuel ( $\ce{H2}$ ) and oxidant ( $\ce{O2}$ ).
Corrosion: Spontaneous oxidation of metal by ambient oxygen and moisture. Iron rusts through an electrochemical micro-cell mechanism.
Cathodic Protection: Connecting a more active metal (more negative reduction potential) as a sacrificial anode forces the protected steel structure to act as the cathode.

Mnemonic: Remember AN OX (Anode = Oxidation) and RED CAT (Reduction = Cathode). Electrons always flow through the external wire from Anode to Cathode.

1. Electroplating

Electroplating uses an electrolytic cell to deposit a thin metal layer onto an object. An external power source drives a nonspontaneous redox reaction ( $\Delta G > 0$ , $E^\circ < \pu{0 V}$ ).

The object to be coated is connected as the cathode (the negative terminal in an electrolytic cell). Metal ions in the electrolyte bath migrate to the cathode and undergo reduction to solid metal:

\ce{Cu^{2+}(aq) + 2e^- -> Cu(s)}

The anode consists of the coating metal (e.g., pure copper). It oxidizes and dissolves to replenish cations in solution:

\ce{Cu(s) -> Cu^{2+}(aq) + 2e^-}

Worked Example: Electroplating Setup
Problem: To plate a nickel key with copper using a

\ce{CuSO4(aq)}

electrolyte, which electrode should the key be connected to, and what reaction occurs at its surface?
Solution: To coat the key with solid copper,

\ce{Cu^{2+}(aq)}

ions must gain electrons (reduction). Reduction always occurs at the cathode, so the key must be connected to the cathode where

\ce{Cu^{2+}(aq) + 2e^- -> Cu(s)}

takes place.
DAT Exam Trap: Students often mistakenly connect the object to the anode because the anode is positively charged in electrolytic cells. Remember: metal deposition always occurs at the cathode where reduction takes place!

2. Practical Galvanic Cells: Batteries and Fuel Cells

Galvanic (voltaic) cells convert chemical energy into electrical energy via spontaneous redox reactions ( $\Delta G < 0$ , $E^\circ > \pu{0 V}$ ).

Batteries

Batteries are sealed electrochemical power units. Primary batteries (e.g., alkaline dry cells) use non-reversible reactions and cannot be recharged. Secondary batteries (e.g., lead-acid, lithium-ion) can be recharged by applying an external voltage that reverses the spontaneous reaction.

Fuel Cells

A fuel cell is a galvanic cell operating as an open system. Rather than storing a finite supply of chemical reactants internally, fuel cells continuously consume fuel (e.g., $\ce{H2(g)}$ ) at the anode and an oxidant (e.g., $\ce{O2(g)}$ ) at the cathode.

In a acidic Proton Exchange Membrane (PEM) fuel cell, the half-reactions are:

Anode (oxidation): $\ce{2H2(g) -> 4H^+(aq) + 4e^-}$

Cathode (reduction): $\ce{O2(g) + 4H^+(aq) + 4e^- -> 2H2O(l)}$

Overall reaction: $\ce{2H2(g) + O2(g) -> 2H2O(l)}$

Worked Example: Fuel Cell Operation
Problem: Identify the chemical product emitted by a hydrogen-oxygen PEM fuel cell.
Solution: Hydrogen is oxidized to

\ce{H^+}

ions at the anode, which migrate across the membrane to reduce oxygen gas at the cathode, yielding only pure liquid water (

\ce{H2O}

).
DAT Exam Trap: Do not confuse fuel cells with standard batteries. Batteries eventually run out of internal reactants, while fuel cells operate continuously as long as fuel is supplied.

3. Corrosion as Spontaneous Oxidation

Corrosion is the unwanted, spontaneous oxidation of a metal by environmental agents. The rusting of iron is an electrochemical process requiring oxygen and moisture.

At an anodic site on the iron surface, elemental iron oxidizes:

\ce{Fe(s) -> Fe^{2+}(aq) + 2e^-}

The released electrons flow through the metal to a cathodic site, where oxygen in moisture is reduced:

\ce{O2(g) + 2H2O(l) + 4e^- -> 4OH^-(aq)}

The resulting $\ce{Fe^{2+}}$ and $\ce{OH^-}$ ions precipitate as $\ce{Fe(OH)2(s)}$ , which undergoes further oxidation and dehydration to form hydrated iron(III) oxide (rust, $\ce{Fe2O3.H2O}$ ).

4. Cathodic Protection with a Sacrificial Anode

To prevent corrosion of steel structures (like buried pipelines or ship hulls), cathodic protection is applied. A more active metal (having a more negative standard reduction potential) is attached to the steel structure.

Because the sacrificial metal (e.g., $\ce{Mg}$ or $\ce{Zn}$ ) oxidizes more readily than iron, it acts as the sacrificial anode:

\ce{Zn(s) -> Zn^{2+}(aq) + 2e^-}

The generated electrons flow into the iron, forcing the entire steel structure to function as a protected cathode. The sacrificial anode corrodes away and is replaced periodically, while the iron remains intact.

Worked Example: Sacrificial Anode Selection
Problem: Given standard reduction potentials:

E^\circ(\ce{Fe^{2+}/Fe}) = \pu{-0.44 V}

E^\circ(\ce{Zn^{2+}/Zn}) = \pu{-0.76 V}

, and

E^\circ(\ce{Cu^{2+}/Cu}) = \pu{+0.34 V}

. Which metal will protect an iron pipe?
Solution: To act as a sacrificial anode, the metal must have a lower (more negative) reduction potential than iron. Zinc (

\pu{-0.76 V}

) is more easily oxidized than iron (

\pu{-0.44 V}

) and will protect the pipe.
DAT Exam Trap: Selecting copper (

\pu{+0.34 V}

) would make copper the cathode and force iron to act as the anode, drastically accelerating iron corrosion!

Summary Comparison of Electrochemical Systems

System Type	Cell Nature	$\Delta G$ & Potential	Anode Reaction / Sign	Cathode Reaction / Sign
Electroplating Cell	Electrolytic	$\Delta G > 0$ , $E^\circ < \pu{0 V}$	Oxidation / Positive (+)	Reduction / Negative (-)
Battery	Galvanic (Closed)	$\Delta G < 0$ , $E^\circ > \pu{0 V}$	Oxidation / Negative (-)	Reduction / Positive (+)
Fuel Cell	Galvanic (Continuous)	$\Delta G < 0$ , $E^\circ > \pu{0 V}$	Oxidation / Negative (-)	Reduction / Positive (+)
Sacrificial Cathodic Protection	Galvanic (Corrosion Control)	$\Delta G < 0$ , $E^\circ > \pu{0 V}$	Sacrificial metal oxidizes / (-)	Protected metal reduces oxygen / (+)

DAT Critical Thinking Questions

Why does exposing a steel ship hull to ocean saltwater accelerate rusting compared to exposure to pure freshwater?
If a galvanized steel sheet (coated with zinc) receives a deep scratch exposing the underlying iron, why does the iron still resist immediate rusting?
Compare a secondary lithium-ion battery to a hydrogen fuel cell. What thermodynamic and practical operational factors make fuel cells advantageous for heavy vehicle applications?

References

Sources returned by Google Search grounding during lesson verification.

Batteries and Fuel Cells - EdTech Books — edtechbooks.org
8.3: Electrochemistry- Cells and Batteries - Chemistry LibreTexts — chem.libretexts.org

Electroplating, Batteries, Fuel Cells, and Corrosion

Summary

Practice Questions

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