Electrochemistry | General Chemistry 3

SI units used in electrochemistry:

• Joule (J): Unit of energy.
• Coulomb (C): Unit of electric charge.
• Volt (V): Unit of electric potential, where 1 V = 1 J.C^-1.
• Ampere (A): Unit of electric current,where 1 A = 1 C.s^-1.

Electrical Energy (U):

The energy associated with electric charge is given by:
 

Electrical Charge (Q):

The total electric charge passing through a circuit is calculated as:
 

Electrochemical reactions:

Chemical reactions can occur as a result of the passage of an electric current through a solution, which is the basis of electrochemical cells and processes like electrolysis.

Oxidation-reduction reactions:

Oxidation-reduction (Redox) reactions involve the transfer of electrons from one substance to another. In a redox reaction:

• Oxidation occurs when a substance loses electrons.
• Reduction occurs when a substance gains electrons.

These electron transfers can be harnessed to produce an electric current, making redox reactions the foundation for galvanic cells.

Galvanic cells:

Galvanic cells are electrochemical cells that convert chemical energy into electrical energy through spontaneous redox reactions. These are commonly used in batteries.

A typical galvanic cell consists of two half-cells, each containing an electrode immersed in an electrolyte solution, connected by an external circuit and a salt bridge.

Structure of a galvanic cell:

• Anode: The electrode where oxidation occurs. The anode has a negative sign in a galvanic cell.
• Cathode: The electrode where reduction occurs. The cathode has a positive sign in a galvanic cell.
• Salt bridge (or porous barrier): Allows ion flow between the half-cells to maintain electrical neutrality.
• External circuit: Provides a pathway for electron flow from the anode to the cathode.

Zinc-copper galvanic cell (Daniell cell):

• Electrodes: Copper and Zinc
• Anode (Zinc): Zn (s) → Zn²⁺ (aq) + 2 e^- [oxidation]
• Cathode (Copper): Cu²⁺ (aq) + 2 e^- → Cu (s) [reduction]
• Salt bridge: Contains ions (e.g., Na⁺ and SO₄^2-) to balance charge as the reaction proceeds.

• Electron flow: Electrons flow from the anode to the cathode through an external circuit.
• Ion flow: Cations move toward the cathode, and anions move toward the anode through the salt bridge to balance the charges in each half-cell.

Electrochemical reactions in a galvanic cell:

Each half-cell contains a redox reaction, and the overall cell reaction is the sum of the two half-reactions.
 

Overall cell reaction in a zinc-copper galvanic cell:

• Anode (Zinc): Zn (s) → Zn²⁺ (aq) + 2 e⁻
• Cathode (Copper): Cu²⁺ (aq) + 2 e⁻ → Cu (s)
• Overall cell reaction: Zn(s) + Cu²⁺ (aq) → Zn²⁺ (aq) + Cu (s)

Cell diagrams:

Galvanic cells are often represented using line notation known as a cell diagram, which shows the components of each half-cell in a standardized format. By convention:

• Metal electrodes are written at the ends of the diagram.
• Left-hand electrode: Represents the oxidation reaction (anode).
• Right-hand electrode: Represents the reduction reaction (cathode).
• Single vertical bars (|): Indicate phase boundaries where different phases are in contact.
• Double vertical bars (||): Represent the salt bridge, which separates the two half-cells.

Cell diagram of a zinc-copper galvanic cell:

• The spontaneous chemical reaction is: Zn (s) + Cu²⁺ (aq) → Zn²⁺ (aq) + Cu (s)
• The cell diagram is Zn (s) | ZnSO₄ (aq) || CuSO₄ (aq) | Cu (s)

Standard reduction potential (E^o):

The standard reduction potential is a measure of the tendency of a species to gain electrons (undergo reduction) under standard conditions, which include:

• 1 M concentration for solutions,
• 1 atm pressure for gases,
• Pure solids or liquids for other reactants or products,
• Temperature at 25^oC or 298K.

Standard reduction potentials are organized into an electrochemical series, ranking substances based on their ability to be reduced:

•  Positive E^o values indicate a strong tendency to gain electrons.
•  Negative E^o values suggest a weaker tendency to be reduced.

Standard Hydrogen Electrode (SHE):

It's not possible to measure the potential of a single electrode directly; only the potential difference between 2 electrodes can be measured.

In chemistry, the Standard Hydrogen Electrode is chosen as the reference electrode, with an assigned potential of 0.00 V. The SHE consists of a platinum electrode in contact with H₂ gas and aqueous H⁺ ions under standard-state conditions: 1 atm H₂ (g), 1 M H⁺ (aq), and 25^oC.

The half-reaction for SHE is:
 

H⁺ (aq, 1 M) + e^- → $\frac{1}{2}$ H₂ (g, 1 atm)     [E^o = 0 V]

Standard cell potential (E^o_cell):

The standard cell potential for a galvanic cell reflects the voltage the cell can produce under standard conditions. It is calculated using the standard reduction potentials of the cathode and anode:
 

For a particular half-reaction, the standard half-cell potential for oxidation at the anode is the negative of the reduction potential at the cathode:
 

Calculating E^o_cell​ for a zinc-copper cell:

Half reactions:

• Anode (Oxidation): Zn (s) → Zn²⁺ (aq) + 2 e⁻     [E^o_ox = -E^o_red = -0.76 V]
• Cathode (Copper): Cu²⁺ (aq) + 2 e⁻ → Cu (s)     [E^o_red = +0.34 V]

Standard cell potential: E^o_cell = E^o_cathode + E^o_anode = 0.34 + (-0.76) = 1.10 V

Spontaneity and free energy change:

The spontaneity of a redox reaction is determined by the Gibbs free energy change (ΔG). A reaction is spontaneous if ΔG < 0.

The relationship between Gibbs free energy change and cell potential (E_cell) is:
 

Relationship between E^o_cell and the standard free energy change ΔG^o:

The standard free energy change for a redox reaction can be calculated using the standard cell potential:
 

Interpretation:

• If E^o_cell > 0 ⇒ ΔG^o < 0: The reaction is spontaneous under standard conditions.
• E^o_cell < 0 ⇒ ΔG^o > 0: The reaction is non-spontaneous under standard conditions.

Determining spontaneity for a Zinc-Copper cell:

• n = 2 for the Zn/Cu reaction
• ΔG^o = - nFE^o_cell = - 2 x 96,485 x 1.10 = - 212.3 kJ.mol^-1

Since ΔG^o < 0, the reaction is spontaneous under standard conditions.

Relationship between E^o_cell and the equilibrium constant K:

The standard cell potential also relates to the equilibrium constant K for the reaction. Given that: ΔG^o = - RT ln K and ΔG^o = - nFE^o_cell, we can rearrange to find:
 

At 298 K, this equation simplifies to:
 

Purpose of the Nerst equation:

The Nernst Equation allows us to calculate the cell potential (E_cell​) for a redox reaction under non-standard conditions, taking into account the concentrations (or partial pressures) of reactants and products.
 

Nernst equation:

The Nerst equation is derived from the relationship between the free energy change and the reaction quotient, ΔG = ΔG^o + RT ln Q.

By dividing this equation by - nF, the general form of the Nernst equation is obtained:
 

At 298 K, this equation simplifies to:
 

Batteries:

Batteries are devices that convert chemical energy into electrical energy through spontaneous redox reactions, providing portable energy sources for a wide range of applications, from small electronics to vehicles.

• Primary batteries: Non-rechargeable, used until reactants are depleted (e.g., alkaline batteries).
• Secondary batteries: Rechargeable, designed for multiple uses by reversing the chemical reactions during recharging (e.g., lithium-ion batteries).

Dry cells and alkaline batteries:

A dry cell battery has no fluid component and consists of a zinc anode and a manganese dioxide (MnO₂) cathode, with an electrolyte paste made of ammonium chloride and zinc chloride in water. This type of battery is used in flashlights, clocks, and remote controls.

Reactions:

• Anode: Zn (s) → Zn²⁺ (aq) + 2 e⁻
• Cathode: 2 MnO₂​ (s) + 2 NH₄⁺​ (aq) + 2 e⁻ → Mn₂​O₃ ​(s) + 2 NH₃​ (aq) + H₂​O (l)
• Overall: Zn (s) + 2 MnO₂ ​(s) + 2 NH₄⁺​ (aq) → Zn²⁺ (aq) + Mn₂​O₃ ​(s) + 2 NH₃ ​(aq) + H_2​O (l)

An alkaline battery is similar to dry cells but uses an alkaline electrolyte (KOH), which provides a longer shelf life and higher energy density.

Reactions:

• Anode: Zn (s) + 2 OH^- (aq) → Zn(OH)₂ (s) + 2 e⁻
• Cathode: 2 MnO₂​ (s) + 2 H₂O​ (l) + 2 e⁻ → 2 MnO(OH) (s) + 2 OH^- (aq)
• Overall: Zn (s) + 2 MnO₂ ​(s) + 2 H₂O​ (l) → Zn(OH)₂ (s) + 2 MnO(OH) (s)

Lead storage batteries:

A lead storage battery consists of lead (Pb) anodes and lead dioxide (PbO₂​) cathodes in a sulfuric acid (H₂SO₄) electrolyte. It is used in automobiles for starting and lighting, due to high power output and reliability.

Reactions:

• Anode: Pb (s) + SO₄²⁻ ​(aq) → PbSO₄​ (s) + 2 e⁻
• Cathode: PbO₂​(s) + 4 H⁺ (aq) + SO₄²⁻ ​(aq) + 2 e⁻ → PbSO₄​ (s) + 2 H₂​O (l)
• Overall: Pb (s) + PbO₂ (s) + 4 H⁺ (aq) + 2 SO₄^2- (aq) → 2 PbSO₄ (s) + 2 H₂O (l)

Unlike dry cells and alkaline batteries, lead storage batteries can be recharged by reversing the redox reaction.

Lithium-ion batteries:

A lithium-ion battery contains a lithium cobalt oxide (LiCoO₂​) cathode and a graphite anode, with a lithium salt in an organic solvent as the electrolyte. It is widely used in portable electronics, such as smartphones, laptops, and electric vehicles.

Reactions:

• Anode: Li (s) → Li⁺ + e^-
• Cathode: Li⁺ + CoO₂ (s) + e^- → LiCoO₂ (s)
• Overall: Li (s) + CoO₂ (s) + → LiCoO₂ (s)

Advantages: High energy density, lightweight, and long cycle life. However, these batteries are sensitive to high temperatures and may pose safety risks if damaged.

Fuel cells:

A fuel cell is a galvanic cell that requires a continuous supply of reactants to produce electricity. The hydrogen fuel cell is an example of a fuel cell that combines hydrogen fuel with oxygen from the air, producing only water as a byproduct. This type of cell is used in electric vehicles, space exploration, and backup power systems.

Reactions:

• Anode: 2 H₂ ​(g) + 4 OH^- (aq) → 4 H₂O (l) + 4 e^-
• Cathode: O₂ ​(g) + 2 H₂O (l) + 4 e⁻ → 4 OH^- (aq)
• Overall: 2 H₂ ​(g) + O₂ ​(g) → 2 H₂​O (l)

Electrolysis:

Electrolysis is a process that uses electrical energy to drive a non-spontaneous chemical reaction. This occurs in an electrolytic cell, where an external power source supplies the necessary energy.

Overvoltage is the additional voltage required beyond the theoretical cell potential to drive a reaction at an electrode, reducing the efficiency of electrolysis by requiring more energy than the theoretical minimum.

Electrolysis of molten sodium chloride:

This process is a major industrial source of pure sodium metal and chlorine gas:

• Anode: 2 Cl⁻ (l) → Cl₂ (g) + 2 e⁻
• Cathode: 2 Na⁺ (l) + 2 e⁻ → 2 Na (l)
• Overall: 2 Cl⁻ (l) + 2 Na⁺ (l) → 2 Na (l) + Cl₂ (g)

Electrolysis of water:

This process produces hydrogen and oxygen gases:

• Anode: 2 H_2​O (l) → O₂​ (g) + 4 H⁺ (aq) + 4 e⁻
• Cathode: 4 H⁺ (aq) + 4 e⁻ → 2 H₂ (g)
• Overall: 2 H_2​O (l) → 2 H₂ ​(g) + O₂ ​(g)

Electrolysis of aqueous sodium chloride:

This process produces hydrogen gas, chlorine gas, and sodium hydroxide solution:

• Anode: 2 Cl⁻ (aq) → Cl₂ (g) + 2 e⁻
• Cathode: 2 H_2​O (l) + 2 e⁻ → H₂ ​(g) + 2 OH⁻ (aq)
• Overall: 2 H_2​O (l) + 2 Cl⁻ (aq) → H₂ ​(g) + 2 OH⁻ (aq) + Cl₂ (g)

Faraday's laws of electrolysis:

First Law: The amount of substance produced at each electrode is directly proportional to the total charge passed through the electrolyte.
 

Second Law: For the same amount of charge, the mass of a substance deposited or dissolved at each electrode is proportional to its equivalent weight.

Corrosion:

Corrosion is a natural electrochemical process where metals deteriorate due to reactions with environmental agents, typically involving oxygen and moisture. The most common example is the rusting of iron, which forms iron oxides, gradually weakening the metal structure.

Electrochemical process of corrosion:

Corrosion is an electrochemical redox reaction:

• Anode: The metal loses electrons and is oxidized.
• Cathode: Environmental agents, usually oxygen, gain electrons and are reduced.

Corrosion of iron: 

• Anode: 2 Fe (s) → 2 Fe²⁺ (aq) + 4 e⁻
• Cathode: O₂ (g) + 4 H⁺ (aq) + 4 e⁻ → 2 H₂O (l)
• Overall: 2 Fe (s) + O₂ ​(g) + 4 H⁺ (aq) → 2 Fe²⁺ (aq) + 2 H₂​O (l)

Further oxidation can lead to iron (III) oxides, commonly known as rust, forming hydrated iron (III) oxide (Fe₂O₃.nH₂O).

Factors affecting corrosion:

• Presence of water and oxygen: Moisture accelerates corrosion, as water acts as an electrolyte, and oxygen is required for the cathodic reaction.
• pH of the environment: Acidic conditions increase corrosion rates by providing more H⁺ ions for reduction.
• Salt and other electrolytes: The presence of salts, such as NaCl, increases conductivity, which facilitates electron transfer and accelerates corrosion.
• Temperature: Higher temperatures generally increase the rate of corrosion due to faster reaction kinetics.

Methods of corrosion prevention:

• Galvanizing: Coating iron or steel with a layer of zinc, which corrodes preferentially to protect the underlying metal.
• Cathodic protection: Attaching a more reactive metal, like magnesium or zinc, to act as a sacrificial anode, which corrodes in place of the protected metal.
• Alloying: Adding elements such as chromium or nickel to form corrosion-resistant alloys (e.g., stainless steel).