Electrochemistry - Theory & Concepts

Electrochemistry

The relationship between chemical energy and electrical energy.

Electrochemistry is the study of the relationship between electricity and chemical reactions. For civil engineers, the most significant application of electrochemistry is understanding and preventing corrosion, particularly the rusting of steel reinforcement in concrete structures, pipelines, and bridges.

Oxidation-Reduction (Redox) Reactions

All electrochemical processes involve redox reactions, where electrons are transferred from one substance to another.

Redox Basics

"LEO says GER" (Lose Electrons = Oxidation; Gain Electrons = Reduction).

Oxidation: Loss of electrons. The oxidation state increases. The substance oxidized is the reducing agent.
Reduction: Gain of electrons. The oxidation state decreases. The substance reduced is the oxidizing agent.

Example: Rusting of Iron

4Fe(s) + 3O_2(g) \rightarrow 2Fe_2O_3(s)

Iron (

Fe

) goes from 0 to +3 (Oxidized). Oxygen (

O_2

) goes from 0 to -2 (Reduced).

Electrochemical Cells

Galvanic (Voltaic) Cells

Spontaneous redox reactions that generate electrical energy (e.g., batteries).

Anode: Electrode where Oxidation occurs (Anode $\rightarrow$ Oxidation). It is negatively charged in a galvanic cell.
Cathode: Electrode where Reduction occurs (Cathode $\rightarrow$ Reduction). It is positively charged.
Salt Bridge: Contains inert ions that flow to balance the charge buildup in the half-cells, completing the circuit.
Cell Potential ( $E_{cell}$ ): The driving force (voltage) of the cell. Calculated using standard reduction potentials ( $E^\circ$ ).

E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}

E^\circ_{cell} > 0

, the reaction is spontaneous.

The Nernst Equation

The Nernst equation allows engineers to calculate the cell potential under non-standard conditions, which is essential since real-world environments (like the water in soil around a pipe) are rarely at

1 \, \text{M}

concentrations.

The Nernst Equation

Calculates cell potential under non-standard conditions.

E_{cell} = E^\circ_{cell} - \frac{RT}{nF} \ln Q

Variables

Symbol	Description	Unit
$E_{cell}$	Cell potential under non-standard conditions	-
$E^\circ_{cell}$	Standard cell potential	-
$R$	Universal gas constant (8.314 J/(mol·K))	-
$T$	Temperature (K)	-
$n$	Number of moles of electrons transferred	-
$F$	Faraday's constant (96,485 C/mol)	-
$Q$	Reaction quotient	-

Galvanic Cell Simulator

Half-Cell 1

Half-Cell 2

Cathode (Reduction +)

Copper (Cu)

Cu^{2+} + 2e^- \rightarrow Cu(s)

E^\circ_{red} = +0.34 \, V

Anode (Oxidation -)

Zinc (Zn)

Zn(s) \rightarrow Zn^{2+} + 2e^-

E^\circ_{red} = -0.76 \, V

E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}

E^\circ_{cell} = (+0.34) - (-0.76)

Standard Cell Potential

+1.10 V

Electrolytic Cells

Non-spontaneous reactions forced to occur by applying an external electrical current.

Examples: Electroplating, extraction of aluminum from bauxite ore, separating water into $H_2$ and $O_2$ .
The anode is positive, and the cathode is negative (opposite of galvanic), but oxidation still occurs at the anode.

Corrosion

Corrosion is essentially an unwanted galvanic cell where a metal is oxidized by its environment. The most economically devastating form in civil engineering is the rusting of iron and steel rebar.

The Mechanism of Rusting

Rusting requires three components: Iron, Oxygen, and Water (which acts as the electrolyte).

Anodic Region (Oxidation): Iron is oxidized to $Fe^{2+}$ .
$Fe(s) \rightarrow Fe^{2+}(aq) + 2e^-$
Cathodic Region (Reduction): Electrons travel through the metal to a region exposed to $O_2$ and $H_2O$ .
$O_2(g) + 4H^+(aq) + 4e^- \rightarrow 2H_2O(l)$
Formation of Rust: $Fe^{2+}$ is further oxidized to $Fe^{3+}$ and forms hydrated iron(III) oxide.
$Fe^{2+} \rightarrow Fe_2O_3 \cdot xH_2O \text{ (Rust)}$

Accelerated Corrosion: The presence of salts (like $NaCl$ from winter road de-icing or chlorides in seawater environments) dramatically accelerates corrosion by increasing the conductivity of the water electrolyte, allowing ions to flow much faster. This is why marine structures degrade quickly.

Corrosion Prevention

Engineers use several methods to interrupt the electrochemical cell responsible for corrosion, extending the life of infrastructure.

Prevention Strategies

Barrier Coatings: Paint, epoxy, or plastic prevent oxygen and water from reaching the metal surface. (e.g., green Epoxy-coated rebar used in bridge decks).
Galvanization: Coating iron with a layer of zinc. Zinc has a more negative reduction potential and is more easily oxidized than iron. Even if the coating is scratched, the zinc acts as a sacrificial anode.
Cathodic Protection (Sacrificial Anode): Attaching a more reactive metal (like Magnesium or Zinc) directly to the protected structure (like an underground steel pipe or ship hull). The active metal oxidizes (sacrifices itself) instead of the steel.
Impressed Current Cathodic Protection (ICCP): Using a DC power source to force electrons into the steel structure, making it the cathode and preventing oxidation entirely. Often used for large pipelines and reinforced concrete bridges.

Key Takeaways

Redox reactions involve the transfer of electrons (Oxidation = Loss of $e^-$ , Reduction = Gain of $e^-$ ).
Galvanic cells produce electricity spontaneously (+V); Electrolytic cells require electricity to drive non-spontaneous reactions (-V).
The Nernst Equation is used to calculate potential under non-standard real-world conditions.
Corrosion is a natural galvanic process requiring metal, oxygen, and an electrolyte (water). Chlorides accelerate this process.
Cathodic protection (sacrificial anodes or impressed current) prevents corrosion by forcing the protected metal to act as a cathode, halting oxidation.

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