Electrochemistry and Corrosion: Study Notes for Engineering Students

Electrochemistry

Introduction to Electrochemistry

Electrochemistry is the branch of chemistry that studies chemical processes which involve the movement of electrons, resulting in the generation of electric current or the use of electricity to drive chemical reactions. It is fundamental in understanding batteries, corrosion, and industrial processes such as electroplating.

• Electrochemical reactions involve oxidation (loss of electrons) and reduction (gain of electrons).

• Applications include batteries, corrosion prevention, and metal refining.

Corrosion

Types of Corrosion

Corrosion is the degradation of metals due to chemical reactions with their environment. It can occur in several forms:

• Uniform Corrosion: Occurs evenly over a large surface area.

• Galvanic Corrosion: Occurs when two different metals are in electrical contact in the presence of an electrolyte.

• Crevice Corrosion: Occurs in confined spaces where the electrolyte becomes stagnant.

Galvanic Series

The Galvanic Series ranks metals by their tendency to corrode in a given environment. Metals higher in the series (anodes) are more susceptible to corrosion than those lower (cathodes).

Anode (Active)	Cathode (Noble)
Magnesium	Gold
Zinc	Platinum
Aluminum	Graphite
Mild Steel	Silver
Cast Iron	Titanium
Lead	Stainless Steel
Brass	Nickel
Copper	Monel

Additional info: The series helps predict which metal will corrode when two are in contact.

Corrosion Mechanisms

• Corrosion involves redox reactions where the metal is oxidized.

• For example, iron forms iron(III) oxide (rust) when exposed to oxygen and water.

• Aluminum forms a protective oxide layer, preventing further corrosion.

Redox Reactions and Galvanic Cells

Oxidation and Reduction

Redox (oxidation-reduction) reactions involve the transfer of electrons between species.

• Oxidation: Loss of electrons (e.g., Cu(s) → Cu2+(aq) + 2e-)

• Reduction: Gain of electrons (e.g., Ag+(aq) + e- → Ag(s))

Half-Reactions

Redox reactions are split into two half-reactions, one for oxidation and one for reduction. The number of electrons lost and gained must balance.

• Example: $Cu(s) ightarrow Cu^{2+}(aq) + 2e^-$ $2Ag^+(aq) + 2e^- ightarrow 2Ag(s)$ Net reaction: $Cu(s) + 2Ag^+(aq) ightarrow Cu^{2+}(aq) + 2Ag(s)$

Galvanic Cell Construction

A galvanic cell generates electrical energy from spontaneous redox reactions.

• Anode: Site of oxidation (negative electrode)

• Cathode: Site of reduction (positive electrode)

• Salt Bridge: Maintains charge neutrality by allowing ion flow between half-cells

• Cell Notation: Anode | Anode electrolyte || Cathode electrolyte | Cathode

Example cell notation: $Cu(s) | Cu^{2+}(aq, 1M) || Ag^+(aq, 1M) | Ag(s)$

Electrode Potentials

Standard Electrode Potentials

Standard electrode potentials ($E^ ext{o}$) measure the tendency of a half-cell to gain electrons (be reduced) compared to the standard hydrogen electrode (SHE), which is assigned $E^ ext{o} = 0$ V.

• Half-cell reactions are written as reductions.

• Positive $E^ ext{o}$: Species is easily reduced (good oxidizing agent).

• Negative $E^ ext{o}$: Species is easily oxidized (good reducing agent).

Calculating Cell Potentials

The standard cell potential is calculated as:

• $E^ ext{o}_ ext{cell} = E^ ext{o}_ ext{cathode} - E^ ext{o}_ ext{anode}$

Example: $Fe^{2+}(aq) + 2e^- ightarrow Fe(s) \, E^ ext{o} = -0.44$ V $Cu^{2+}(aq) + 2e^- ightarrow Cu(s) \, E^ ext{o} = 0.337$ V $E^ ext{o}_ ext{cell} = 0.337 - (-0.44) = 0.78$ V

Nernst Equation

The Nernst equation allows calculation of cell potential under non-standard conditions:

• $E = E^ ext{o} - \frac{RT}{nF} \ln Q$

• Where $R$ is the gas constant, $T$ is temperature (K), $n$ is number of electrons, $F$ is Faraday's constant, and $Q$ is the reaction quotient.

Free Energy and Equilibrium

Relationship to Gibbs Free Energy

The cell potential is related to the Gibbs free energy change ($\Delta G$):

• $\Delta G^ ext{o} = -nFE^ ext{o}_ ext{cell}$

• Negative $\Delta G^ ext{o}$ indicates a spontaneous reaction.

Equilibrium Constant

At equilibrium, the relationship between cell potential and equilibrium constant ($K$) is:

• $\Delta G^ ext{o} = -RT \ln K$

• $E^ ext{o}_ ext{cell} = \frac{0.0592}{n} \log K$ (at 25°C)

Batteries and Fuel Cells

Primary and Secondary Batteries

Batteries are devices that convert chemical energy into electrical energy.

• Primary batteries: Non-rechargeable (e.g., alkaline batteries)

• Secondary batteries: Rechargeable (e.g., lead-acid, nickel-cadmium)

Common Battery Types

• Alkaline battery: Anode: $Zn(s) + 2OH^-(aq) ightarrow Zn(OH)_2(s) + 2e^-$ Cathode: $2MnO_2(s) + H_2O(l) + 2e^- ightarrow Mn_2O_3(s) + 2OH^-(aq)$

• Lead-acid battery: Anode: $Pb(s) + HSO_4^-(aq) ightarrow PbSO_4(s) + H^+(aq) + 2e^-$ Cathode: $PbO_2(s) + 3H^+(aq) + HSO_4^-(aq) + 2e^- ightarrow PbSO_4(s) + 2H_2O(l)$

• Fuel cell: Anode: $H_2 ightarrow 2H^+ + 2e^-$ Cathode: $O_2 + 4H^+ + 4e^- ightarrow 2H_2O$ Overall: $2H_2 + O_2 ightarrow 2H_2O$

Electrolysis and Electroplating

Electrolysis

Electrolysis uses electrical energy to drive non-spontaneous chemical reactions. It is used for metal refining and electroplating.

• Electrodes may be inert or participate in the reaction.

• Example: Hall-Héroult process for aluminum refining uses cryolite to lower melting point.

Electroplating

Electroplating deposits a thin layer of metal onto a surface to prevent corrosion or improve appearance.

• Example: Silver plating Anode: $Ag(s) + 2CN^-(aq) ightarrow Ag(CN)_2^-(aq) + e^-$ Cathode: $Ag(CN)_2^-(aq) + e^- ightarrow Ag(s) + 2CN^-(aq)$

Stoichiometry of Electrolysis

The amount of substance deposited or dissolved during electrolysis can be calculated using Faraday's laws:

• $Q = I imes t$ (Charge in coulombs = current × time)

• Moles of electrons: $n = \frac{Q}{F}$

• Example: To deposit gold, calculate charge and use molar mass to find mass deposited.

Corrosion Prevention

Techniques for Preventing Corrosion

• Protective coatings: Paint or electroplating to isolate metal from environment.

• Cathodic protection: Attach a more easily oxidized metal (e.g., magnesium) as a sacrificial anode.

• Material selection: Use metals that form protective oxide layers (e.g., aluminum).

Cathodic Protection

Connecting a sacrificial anode (e.g., magnesium) to iron forces iron to act as the cathode, preventing its oxidation. The sacrificial anode must be replaced periodically.

References

• Holme, T.A. Chemistry for Engineering Students.

• Beran & Hurley. Chemistry Principles and Reactions.

• Silberberg. Chemistry: The Molecular Nature of Matter and Change.

• Brown et al. Chemistry: The Central Science.

• Silberberg. Principles of General Chemistry.

Additional info:

• Some content inferred and expanded for clarity and completeness.

• Images referenced (e.g., iron and steel processing) are not included in text but relate to industrial applications of electrochemistry and corrosion.