Electrochemistry Class 12 Chemistry Chapter 2 Notes

Daniell Cell and Galvanic/Electrolytic Cells

Daniell Cell and Its Redox Reaction
- The Daniell cell is a device that converts chemical energy into electrical energy.
- It operates based on the redox reaction:
  - Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
- This reaction releases energy which is harnessed as electrical energy.
Cell Potential of Daniell Cell
- The Daniell cell has an electrical potential of 1.1 V.
- This potential is achieved when the concentration of Zn2+ and Cu2+ ions is both 1 mol dm–3.
Types of Cells: Galvanic and Voltaic
- The device like the Daniell cell is often referred to as a galvanic or a voltaic cell.
- These cells convert chemical energy to electrical energy spontaneously.
External Potential in Galvanic Cell
- In a galvanic cell, if an external opposite potential is applied gradually, the redox reaction continues.
- It proceeds until the opposing voltage matches the cell potential of 1.1 V.
- When this happens, the reaction ceases, and no current flows through the cell.
Transition to Electrolytic Cell
- If the external potential exceeds 1.1 V, the reaction restarts, but now in the opposite direction.
- The cell transforms into an electrolytic cell.
- Electrolytic cells use electrical energy to drive non-spontaneous chemical reactions.
Importance of Both Cell Types
- Both galvanic and electrolytic cells are significant in various applications.
- Galvanic cells provide a source of electrical energy.
- Electrolytic cells are used to carry out specific chemical processes that require energy input.

Understanding Galvanic Cells and Electrode Potentials

1. Galvanic Cells and Electrical Energy Conversion

Galvanic cells are electrochemical devices that convert the chemical energy from spontaneous redox reactions into electrical energy.
The Gibbs energy of the spontaneous redox reaction is transformed into usable electrical work, which can power various electrical devices.

2. Daniell Cell Redox Reaction

The Daniell cell, an example of a galvanic cell, operates based on the following redox reaction:
- Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
This reaction comprises two half-reactions that combine to give the overall cell reaction.

3. Half-Cells or Redox Couples

Galvanic cells have two portions known as half-cells or redox couples.
The copper electrode hosts the reduction half-reaction, while the zinc electrode hosts the oxidation half-reaction.
These half-cells play distinct roles, with the copper electrode being the reduction half-cell and the zinc electrode being the oxidation half-cell.

4. Construction of Galvanic Cells

Galvanic cells can be constructed by combining different half-cells.
Each half-cell consists of a metallic electrode immersed in an electrolyte.
The two half-cells are connected externally through a metallic wire, a voltmeter, and a switch.
A salt bridge internally connects the electrolytes of the two half-cells. In some cases, both electrodes are placed in the same electrolyte, eliminating the need for a salt bridge.

5. Electrode-Electrolyte Interface

At the interface of each electrode and electrolyte, there is a tendency for metal ions to deposit on the metal electrode, making it positively charged.
Simultaneously, metal atoms from the electrode tend to dissolve as ions into the solution, leaving electrons behind, making the electrode negatively charged.
At equilibrium, a charge separation occurs, resulting in the development of an electrode potential.

6. Standard Electrode Potential

When the concentrations of all species in a half-cell are unity, the electrode potential is termed standard electrode potential.
IUPAC convention refers to standard reduction potentials as standard electrode potentials.

7. Anode and Cathode in Galvanic Cells

In a galvanic cell, the half-cell where oxidation occurs is called the anode. It has a negative potential relative to the solution.
The half-cell where reduction occurs is called the cathode. It has a positive potential relative to the solution.
A potential difference exists between the two electrodes, leading to electron flow from the negative electrode (anode) to the positive electrode (cathode). The direction of current flow is opposite to electron flow.

8. Cell Potential and Electromotive Force (EMF)

The potential difference between the two electrodes in a galvanic cell is called the cell potential, measured in volts.
When no current flows through the cell, the cell potential is termed the cell electromotive force (emf).
Conventionally, the anode is placed on the left, and the cathode is on the right when representing a galvanic cell.
The emf of the cell is positive and can be calculated as: Ecell = Eright – Eleft.

9. Example

Example cell reaction: Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s)
Half-cell reactions:
- Cathode (reduction): 2Ag+(aq) + 2e– → 2Ag(s)
- Anode (oxidation): Cu(s) → Cu2+(aq) + 2e–
Cell representation: Cu(s)|Cu2+(aq)||Ag+(aq)|Ag(s)
Cell potential calculation: Ecell = EAg+˙Ag – ECu2+˙Cu

Electrode Potentials, Standard Hydrogen Electrode, and Their Significance

1. Measuring Half-Cell Potentials

Individual half-cell potentials cannot be directly measured.
Only the difference between the potentials of two half-cells, known as the cell electromotive force (emf), can be determined.

2. Standard Hydrogen Electrode (SHE)

The standard hydrogen electrode (SHE) serves as a reference half-cell with an assigned potential of zero at all temperatures.
SHE comprises a platinum electrode coated with platinum black immersed in an acidic solution, with pure hydrogen gas bubbled through it.
The concentration of both reduced and oxidized hydrogen forms is maintained at unity (1 M H+ and 1 bar H2 pressure at 298 K).

3. Calculating Standard Electrode Potential (E°)

At 298 K, the emf of a cell constructed with SHE as the reference half-cell and the other half-cell gives the reduction potential of the latter.
If the concentrations of species in the right half-cell are unity, the cell potential equals the standard electrode potential (E°) of that half-cell.
E° = E°_R – E°_L, where E°_L for SHE is zero.

4. Example Emf Measurements

Emf measured for Cu2+(aq) + 2e– → Cu(s) half-cell is 0.34 V.
Emf measured for Zn2+(aq) + 2e– → Zn(s) half-cell is -0.76 V.
Positive values indicate easier reduction of Cu2+ compared to H+ ions, while negative values indicate H+ ions can oxidize Zn.

5. Representation of Daniell Cell Half-Reactions

Left electrode (Zn): Zn(s) → Zn2+(aq) + 2e–
Right electrode (Cu): Cu2+(aq) + 2e– → Cu(s)
The overall cell reaction: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Cell potential (E°_cell) = 1.10 V

6. Inert Electrodes

Inert electrodes like platinum (Pt) or gold (Au) are used in some cells.
They do not participate in the reactions but provide a surface for oxidation, reduction, and electron conduction.

7. Importance of Standard Electrode Potentials

Standard electrode potentials are crucial for understanding chemical reactions.
Higher E° values indicate species’ greater tendency to be reduced, making them strong oxidizing agents.
Lower E° values indicate stronger reducing agents.
Electrochemical cells are widely used for various chemical analyses, including pH determination, solubility product determination, equilibrium constant determination, and potentiometric titrations.

8. Relationship between E° and Chemical Properties

As we move from top to bottom in Table 2.1, standard electrode potentials decrease.
Lower E° values correspond to weaker oxidizing agents and stronger reducing agents.
For example, fluorine gas (F2) has the highest E°, signifying its strong oxidizing ability, while lithium (Li) has the lowest E°, indicating its powerful reducing capacity in aqueous solutions.

Nernst Equation and Concentration Dependence of Electrode Potential

1. Introduction to the Nernst Equation

The Nernst equation, developed by Walther Nernst, allows us to determine electrode potentials at non-standard conditions, considering the concentration of species involved in the electrode reaction.
The general electrode reaction is represented as: Mn+(aq) + ne– → M(s)
The electrode potential (E) at a given concentration [Mn+] can be calculated with respect to the standard hydrogen electrode (SHE) using the Nernst equation:
- E = E° – (RT / (nF)) * ln([M]/[Mn+])
Where:
- E is the electrode potential at concentration [Mn+].
- E° is the standard electrode potential.
- R is the gas constant (8.314 J K–1 mol–1).
- T is the temperature in Kelvin.
- F is Faraday’s constant (96487 C mol–1).
- n is the number of electrons transferred in the electrode reaction.
- [M] is the concentration of the reduced form of the species (M).
- [Mn+] is the concentration of Mn+ ions involved in the reaction.

2. Application to the Daniell Cell

For the cathode (Cu electrode) and anode (Zn electrode) in the Daniell cell, the Nernst equation can be applied:
- Cathode (Cu): E(cathode) = E°(Cu/Cu2+) – (RT / (2F)) * ln([Cu2+]/[Cu])
- Anode (Zn): E(anode) = E°(Zn/Zn2+) – (RT / (2F)) * ln([Zn2+]/[Zn])

3. Calculation of Cell Potential for the Daniell Cell

The cell potential (E(cell)) is the difference between the cathode and anode potentials:
- E(cell) = E(cathode) – E(anode)
E(cell) depends on the concentrations of Cu2+ and Zn2+ ions and varies with changes in these concentrations.
An increase in Cu2+ ion concentration or a decrease in Zn2+ ion concentration results in an increase in E(cell).

4. Generalization for a Cell Reaction

For a cell reaction involving species A, B, C, and D as described by the equation: aA + bB + ne– → cC + dD
The Nernst equation becomes: E(cell) = E° – (RT / (nF)) * ln(Q)
Where Q is the reaction quotient: Q = [C]^c [D]^d / [A]^a [B]^b

5. Conversion to Base 10 and Simplified Nernst Equation

By converting the natural logarithm to base 10 and substituting values for R, F, and T (298 K), the Nernst equation can be simplified to:
- E(cell) = E° – (0.0592 / n) * log(Q)
This simplified form is useful for practical calculations.

6. Consistency in Electron Transfer

It is important to ensure that the same number of electrons (n) is used for both electrodes in a cell to calculate the cell potential accurately.

The Nernst equation is a vital tool for calculating electrode potentials under non-standard conditions, considering the concentration of species involved in the electrochemical reaction. It allows for a better understanding of the dependence of electrode potential on concentration changes and plays a crucial role in various electrochemical applications and analyses.

Relationship Between Equilibrium Constants and Standard Cell Potentials

1. Introduction to Equilibrium in the Daniell Cell

When the circuit in the Daniell cell is closed, the redox reaction occurs: Zn(s) + Cu2+(aq) ⇌ Zn2+(aq) + Cu(s)
Over time, Zn2+ concentration increases, while Cu2+ concentration decreases.
The cell voltage, as measured by the voltmeter, gradually decreases.
Eventually, equilibrium is reached when both Zn2+ and Cu2+ concentrations stabilize, and the voltmeter reads zero.

2. Nernst Equation at Equilibrium

At equilibrium, the Nernst equation can be written as: E(cell) = 0 = E° – (2.303 * RT / (2F)) * log([Zn2+]/[Cu2+])
Where:
- E(cell) is the cell voltage at equilibrium, which is zero.
- E° is the standard cell potential (1.1 V for the Daniell cell).
- R is the gas constant.
- T is the temperature (298 K).
- F is Faraday’s constant.
- [Zn2+] and [Cu2+] are the concentrations of Zn2+ and Cu2+ ions at equilibrium.

3. Equilibrium Constant Expression

At equilibrium, the concentrations of Zn2+ and Cu2+ ions can be represented as [Zn2+] and [Cu2+].
The equilibrium constant (Kc) for the reaction Zn(s) + Cu2+(aq) ⇌ Zn2+(aq) + Cu(s) is given by: Kc = [Zn2+] [Cu] / [Zn] [Cu2+]
At T = 298 K, the Nernst equation can be simplified to: 0 = 1.1 V – (0.0592 V) * log(Kc)

4. Calculating the Equilibrium Constant

Solving for Kc: Kc = 2 × 10^37 at 298 K.
The general relationship between the equilibrium constant (Kc) and the standard cell potential (E°) for a cell reaction is given by:
- E(cell) = (2.303 * RT / (nF)) * log(Kc)

5. Significance of Equation. E(cell) = (2.303 * RT / (nF)) * log(Kc)

This equation establishes a direct relationship between the equilibrium constant of a reaction (usually difficult to measure directly) and the standard potential of the cell in which that reaction occurs.
This relationship allows for the calculation of equilibrium constants from experimentally determined standard cell potential values.

Relationship Between Electrical Work, Gibbs Energy, and Standard Cell Potential

1. Electrical Work in a Galvanic Cell

Electrical work done in one second is equal to the electrical potential (voltage) multiplied by the total charge passed through the cell.
To maximize work from a galvanic cell, the charge must be passed reversibly.

2. Reversible Work in a Galvanic Cell

The reversible work done by a galvanic cell is equal to the decrease in its Gibbs energy (ΔrG) during a reaction.
If E is the emf (electromotive force) of the cell, nF is the amount of charge passed, and ΔrG is the Gibbs energy change of the reaction, then the relationship is given by:
- ΔrG = -nFE(cell) (Eq. 2.15)
E(cell) is an intensive parameter, while ΔrG is an extensive thermodynamic property that depends on the value of n.

3. Effect of Reaction Stoichiometry on ΔrG

The value of ΔrG depends on the stoichiometry of the reaction.
For the reaction Zn(s) + Cu2+(aq) ⇌ Zn2+(aq) + Cu(s), ΔrG = -2FE(cell).
When the reaction is written as 2 Zn(s) + 2 Cu2+(aq) ⇌ 2 Zn2+(aq) + 2Cu(s), ΔrG = -4FE(cell).

4. Relationship with Standard Cell Potential

If the concentration of all reacting species is unity, E(cell) = E° (standard cell potential).
In this case, the relationship between standard Gibbs energy change (ΔrG°) and standard cell potential (E°) is given by: ΔrG° = -nFE° (Eq. 2.16).

5. Calculating Equilibrium Constant (K)

From the standard Gibbs energy change (ΔrG°), the equilibrium constant (K) can be calculated using the equation: ΔrG° = -RT ln K.

Electrical Resistance, Resistivity, Conductance, and Conductivity in Electrolytic Solutions

1. Electrical Resistance (R)

Symbol: R
Measured in ohms (Ω), SI units: (kg m²)/(S³ A²).
Resistance is directly proportional to length (l) and inversely proportional to cross-sectional area (A): R ∝ (l / A).
Resistivity (ρ), SI unit: ohm meter (Ω m), is the constant of proportionality.

2. Relationship Between Resistance and Resistivity

R = ρ * (l / A).
Resistivity is the resistance of a substance when it is one meter long with a cross-sectional area of one m².
1 Ω m = 100 Ω cm or 1 Ω cm = 0.01 Ω m.

3. Conductance (G)

Symbol: G
Inverse of resistance: G = 1 / R.
SI unit: siemens (S), equivalent to ohm⁻¹ or mho.
Conductivity (κ) is the inverse of resistivity.

4. Relationship Between Conductance and Conductivity

G = κ * (A / l).
Conductivity is the conductance of a material when it is one meter long with a cross-sectional area of one m².
SI units of conductivity: S m⁻¹, also expressed as S cm⁻¹ (1 S cm⁻¹ = 100 S m⁻¹).

5. Variability of Conductivity

Conductivity varies widely and depends on:
- The nature of the material.
- Temperature and pressure conditions during measurement.
Materials are categorized as conductors, insulators, semiconductors, and superconductors based on their conductivity.

6. Types of Materials Based on Conductivity

Conductors: Materials with high conductivity, including metals and some non-metals like carbon-black and graphite.
Insulators: Materials with very low conductivity, such as glass and ceramics.
Semiconductors: Materials with conductivity between conductors and insulators, e.g., silicon and gallium arsenide.
Superconductors: Materials with zero resistivity or infinite conductivity at low temperatures, including certain metals and ceramic materials.

7. Conductance Mechanisms

Metallic Conductance: Electron movement is responsible for conductance in metals. It depends on the metal’s nature, structure, valence electrons, and temperature.
Ionic or Electrolytic Conductance: Conductance through ions in solution depends on the nature of the electrolyte, ion size, solvent, concentration, and temperature.

8. Factors Affecting Electrolytic Conductance

Nature of the added electrolyte, size of ions, solvent properties, viscosity, concentration, and temperature influence the conductivity of ionic solutions.

9. Changes in Electrolytic Solutions

Prolonged passage of direct current through ionic solutions can lead to changes in composition due to electrochemical reactions.

Measuring Conductivity in Ionic Solutions

Accurate measurement of resistance in an ionic solution is achieved using alternating current (AC) to prevent changes in composition.
Conductivity cells, specially designed vessels, are used for this purpose.

1. Components of a Conductivity Cell

Conductivity cells consist of two platinum electrodes coated with platinum black.
Electrodes have a cross-sectional area (A) and are separated by a distance (l).
The resistance of the solution between the electrodes is represented by R.

2. Calculation of Resistance

Resistance (R) of the solution in the cell: R = ρ * (l / A).
Cell constant (G*): l/A, depends on electrode distance and area.
G* can be determined by measuring the resistance of a cell containing a known-conductivity solution.

3. Determination of Cell Constant (G)*

Commonly used solution for calibration: KCl solutions with known conductivities at various concentrations and temperatures.
G* calculation: G* = l / A = R / κ.

4. Measurement Setup

Measurement of resistance setup includes resistances R3, R4, R1 (variable), and the conductivity cell with unknown resistance R2.
A Wheatstone bridge is balanced using an oscillator (AC power source) and a detector (e.g., headphone).
Balanced bridge condition: No current flows through the detector.
Relationship for R2 determination: R2 = (R4 / R3) * R1.

5. Conductivity Meters

Modern conductivity meters provide direct readings of conductance or resistance in conductivity cells.

6. Calculation of Conductivity (κ)

Once cell constant and cell resistance are known, conductivity can be determined using the equation: κ = 1 / (G* * R).
Conductivity is a measure of a solution’s ability to conduct electricity and depends on ion charge, size, concentration, and mobility.

7. Molar Conductivity (Lm)

Definition: Molar conductivity (Lm) is a physically meaningful quantity.
Relationship to conductivity (κ): Lm = (κ / c), where c is the concentration in mol m⁻3.
Units of Lm: S m² mol⁻1 when κ is in S m⁻1 and c in mol m⁻3, or S cm² mol⁻1 when κ is in S cm⁻1 and c in mol cm⁻3.
Conversion between units: 1 S m² mol⁻1 = 10⁴ S cm² mol⁻1 or 1 S cm² mol⁻1 = 10⁻⁴ S m² mol⁻1.

Conductivity measurements are essential for understanding the electrical behavior of ionic solutions, and the use of conductivity cells and molar conductivity allows for accurate characterization of such solutions.

Variation of Conductivity and Molar Conductivity with Concentration

Conductivity vs. Concentration:

Conductivity (κ) decreases as the concentration of the electrolyte solution decreases.
This trend holds for both strong and weak electrolytes.
The decrease in conductivity with lower concentration is due to a reduced number of ions per unit volume available to carry electrical current.
Conductivity is defined as the conductance of one unit volume of solution placed between electrodes with unit area and unit length.
This relationship can be expressed as: G = kA/l, where A and l are both unity in their appropriate units (m or cm).

Molar Conductivity vs. Concentration:

Molar conductivity (Λm) is a crucial parameter for understanding electrolytic solutions.
Λm is the conductance of a volume (V) of solution containing one mole of the electrolyte placed between electrodes with unit length and an area sufficient to accommodate this volume.
Molar conductivity increases as the concentration of the electrolyte decreases.
This is because the total volume (V) of solution containing one mole of electrolyte increases as concentration decreases.
In essence, Λm at a given concentration can be viewed as the conductance of the electrolytic solution between the electrodes of a conductivity cell with a unit distance but a cross-sectional area large enough to hold one mole of the electrolyte.
As concentration approaches zero, the molar conductivity reaches its maximum value, known as the limiting molar conductivity (Λ°m).
The behavior of Λm with concentration varies between strong and weak electrolytes.

Strong Electrolytes:

For strong electrolytes, Λm increases gradually with dilution.
This trend can be represented by the equation: Λm = Λ°m – A√c.
In this equation, A is a constant that depends on the type of electrolyte, specifically the charges on the cation and anion produced upon dissociation.
Examples of strong electrolytes include NaCl (1-1 electrolyte), CaCl2 (2-1 electrolyte), and MgSO4 (2-2 electrolyte).

Kohlrausch Law of Independent Migration of Ions:

Kohlrausch’s law states that the limiting molar conductivity (Λ°m) of an electrolyte can be represented as the sum of the individual contributions of its cation and anion.
The law allows for the calculation of Λ°m for an electrolyte by considering the limiting molar conductivities (λ°+) of cations and (λ°–) of anions.
The formula for Λ°m is Λ°m = (n+λ°+) + (n–λ°–), where n+ and n– are the numbers of cations and anions, respectively.
Kohlrausch found that the difference in Λ°m between NaX and KX (for any X) is nearly constant, indicating regularities in Λ°m values for strong electrolytes.

Weak Electrolytes:

Weak electrolytes, such as acetic acid, show a significant increase in Λm with dilution, especially at lower concentrations.
Λ°m for weak electrolytes is not estimated through extrapolation to zero concentration; instead, Kohlrausch’s law is used.
The degree of dissociation (α) is approximated as the ratio of Λm at a given concentration to Λ°m: α = (Λm / Λ°m).
This approach is necessary as weak electrolytes do not fully dissociate at low concentrations due to their lower degree of dissociation.

Understanding the behavior of conductivity and molar conductivity concerning concentration is crucial for characterizing the electrical properties of electrolytic solutions, whether strong or weak electrolytes.

Behavior Of Weak Electrolytes

Degree of Dissociation (α): Weak electrolytes like acetic acid do not completely dissociate into ions in solution, unlike strong electrolytes. The degree of dissociation (α) represents the fraction of the solute molecules that ionize in solution. At higher concentrations, weak electrolytes tend to have lower values of α.
Change in Λm with Dilution: For weak electrolytes, the change in molar conductivity (Λm) with dilution is primarily due to an increase in the degree of dissociation (α). As the solution is diluted, the solvent molecules surround and separate the ions, allowing more of the solute molecules to ionize. Consequently, the number of ions in the total volume of the solution containing 1 mol of the electrolyte increases, leading to a steep increase in Λm, especially at lower concentrations.
Extrapolation to L°m: Unlike strong electrolytes, where you can extrapolate Λm to zero concentration to determine the limiting molar conductivity (Λ°m), this approach does not work for weak electrolytes because Λm increases steeply with dilution. Instead, Λ°m for weak electrolytes is obtained using Kohlrausch’s Law of independent migration of ions.
Kohlrausch’s Law: Kohlrausch’s Law relates the degree of dissociation (α) to the molar conductivity at a given concentration (Λm) and the limiting molar conductivity (Λ°m):α = (Λm / Λ°m)This equation allows you to calculate α at different concentrations for weak electrolytes.
Applications: Kohlrausch’s Law is a valuable tool for determining Λ°m for any electrolyte based on the limiting molar conductivities of individual ions. For weak electrolytes like acetic acid, knowing Λ°m and Λm at a specific concentration (c) allows you to determine the dissociation constant (K) using appropriate mathematical relationships.

Principles Of Electrolysis, Faraday’s Laws, And Its Practical Applications

Electrolytic Cells: Electrolytic cells use an external source of voltage to drive a chemical reaction. They are important in laboratory experiments and industrial processes. The example given involves two copper strips immersed in a copper sulfate solution, where copper ions are reduced at the cathode and copper metal is deposited, while copper is oxidized at the anode, dissolving into copper ions.
Purity of Metals: Electrolysis plays a crucial role in producing pure metals from impure ones. Impure copper, for example, can be converted into high-purity copper by using the impure copper as the anode and depositing pure copper at the cathode during electrolysis.
Production of Metals: Many metals, such as sodium, magnesium, and aluminum, are produced on a large scale by the electrochemical reduction of their cations, especially when suitable chemical reducing agents are not available.
Faraday’s Laws of Electrolysis: Michael Faraday formulated two fundamental laws of electrolysis:
- First Law: The amount of chemical reaction at an electrode during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte.
- Second Law: The amounts of different substances liberated during electrolysis are proportional to their chemical equivalent weights.
Quantitative Aspects: Faraday’s laws provide a quantitative understanding of electrolysis. The passage mentions that Faraday’s work was done in a time when constant current sources were not available, so coulometers were used to determine the quantity of electricity based on the deposition or consumption of metal. Modern constant current sources make it easier to quantify electrolysis.
Faraday’s Constant: The charge on one mole of electrons is defined as one Faraday (F), which is approximately equal to 96,487 coulombs per mole. For practical calculations, 1 F is often approximated as 96,500 C/mol.
Examples: The passage provides examples of electrode reactions, such as the reduction of Ag+ ions and the reduction of Mg2+ and Al3+ ions, to illustrate the relationship between the quantity of electricity and the stoichiometry of the reaction.
Industrial Currents: In commercial metal production, extremely high currents, such as 50,000 amperes, are used, resulting in significant quantities of charge being passed through the electrolytic cell in a short time.

Dependence Of The Products Of Electrolysis On Several Factors

Nature of Electrodes: The products of electrolysis can vary depending on the nature of the electrodes used. Inert electrodes, such as platinum or gold, do not participate in the chemical reactions and act only as sources or sinks for electrons. Reactive electrodes, on the other hand, actively participate in the electrode reactions.
Dependence on Species and Potentials: The products of electrolysis depend on the oxidizing and reducing species present in the cell and their standard electrode potentials. The electrode reactions are driven by the difference in standard electrode potentials.
Overpotential: Some electrochemical processes, although thermodynamically feasible, are kinetically slow. In such cases, an extra potential, known as overpotential, must be applied to make the process occur. Overpotential makes these processes more difficult to achieve.
Example: Electrolysis of NaCl: When molten NaCl is electrolyzed, sodium metal (Na) and chlorine gas (Cl2) are produced. In this case, there is competition between reduction reactions at the cathode, where both sodium ions (Na+) and hydrogen ions (H+) are present. The reaction with the higher standard electrode potential is preferred, which in this case is the reduction of H+ to form hydrogen gas (H2). Additionally, H+ ions are produced by the dissociation of water (H2O).
Example: Electrolysis of Sulphuric Acid: In the electrolysis of sulfuric acid (H2SO4), depending on the concentration of the acid, different processes can occur at the anode. At lower concentrations, the preferred reaction is the oxidation of water (2H2O) to form oxygen gas (O2) and protons (H+). At higher concentrations, the oxidation of sulfate ions (SO4^2-) to form persulfate ions (S2O8^2-) becomes preferred.
Nernst Equation: To account for concentration effects, the standard electrode potentials are replaced by electrode potentials calculated using the Nernst equation.

Primary Batteries:

Primary batteries are non-rechargeable batteries in which the chemical reaction occurs only once, rendering the battery dead after use.

Example – Dry Cell (Leclanche Cell):

Components:
- Zinc container (anode)
- Carbon (graphite) rod (cathode)
- Powdered manganese dioxide
- Carbon
- Ammonium chloride (NH4Cl) and zinc chloride (ZnCl2) paste (electrolyte)
Electrode Reactions:
- Anode: Zn(s) → Zn2+ + 2e–
- Cathode: MnO2 + NH4+ + e– → MnO(OH) + NH3
  - Manganese is reduced from +4 oxidation state to +3.
  - Ammonia forms a complex with Zn2+ to give [Zn(NH3)4]2+.
Cell Potential: Approximately 1.5 V

Mercury Cell:

Suitable for low-current devices like hearing aids and watches.
Components:
- Zinc-mercury amalgam (anode)
- Paste of HgO and carbon (cathode)
- Electrolyte paste of KOH and ZnO
Electrode Reactions:
- Anode: Zn(Hg) + 2OH– → ZnO(s) + H2O + 2e–
- Cathode: HgO + H2O + 2e– → Hg(l) + 2OH–
Overall Reaction: Zn(Hg) + HgO(s) → ZnO(s) + Hg(l)
Cell Potential: Approximately 1.35 V
- The cell potential remains constant over its lifetime because the overall reaction doesn’t involve any ions in solution whose concentration changes.

Secondary Cells

Secondary cells are rechargeable batteries that can be reused after discharging by passing current through them in the opposite direction. They are designed to undergo many discharging and charging cycles.

Lead Storage Battery (Lead-Acid Battery):

Commonly used in automobiles and inverters.
Components:
- Lead anode
- Grid of lead packed with lead dioxide (PbO2) as cathode
- 38% sulfuric acid solution as an electrolyte
Cell Reactions (During Discharge):
- Anode: Pb(s) + SO4^2–(aq) → PbSO4(s) + 2e–
- Cathode: PbO2(s) + SO4^2–(aq) + 4H+(aq) + 2e– → PbSO4(s) + 2H2O(l)
Overall Cell Reaction (Discharge):
- Pb(s) + PbO2(s) + 2H2SO4(aq) → 2PbSO4(s) + 2H2O(l)
During Charging, the above reaction is reversed, converting PbSO4 on the anode and cathode back into Pb and PbO2, respectively.

Nickel-Cadmium Cell (NiCd Battery):

Longer life than lead-acid batteries but more expensive.
Detailed electrode reactions during charging and discharging are not provided.
Overall Reaction (Discharge):
- Cd(s) + 2Ni(OH)3(s) → CdO(s) + 2Ni(OH)2(s) + H2O(l)
The nickel-cadmium cell functions by reversible reactions during charge and discharge, involving the interconversion of cadmium (Cd) and nickel hydroxide (Ni(OH)2) compounds.

Fuel Cells

Fuel cells are highly efficient devices that directly convert chemical energy into electricity through electrochemical reactions. They are more efficient and less polluting compared to traditional thermal power plants that use fossil fuels. Here’s an overview of fuel cells:

Working Principle:

Fuel cells utilize the energy of combustion of fuels like hydrogen, methane, methanol, etc., to generate electricity directly through electrochemical reactions.

Components:

Fuel cells consist of porous carbon electrodes immersed in concentrated aqueous sodium hydroxide (NaOH) solution.
Catalysts, often finely divided platinum or palladium metal, are incorporated into the electrodes to accelerate the electrode reactions.

Electrode Reactions:

Cathode: O2(g) + 2H2O(l) + 4e– → 4OH–(aq)
Anode: 2H2(g) + 4OH–(aq) → 4H2O(l) + 4e–
Overall Reaction: 2H2(g) + O2(g) → 2H2O(l)

Operation:

Hydrogen and oxygen are continuously supplied to the porous carbon electrodes.
As long as the reactants are provided, the fuel cell generates electricity continuously.

Efficiency:

Fuel cells are highly efficient, with an efficiency of about 70%.
In contrast, thermal power plants have an efficiency of about 40%, making fuel cells a more efficient and eco-friendly energy conversion technology.

Applications:

Fuel cells have been used in various applications, including providing electrical power in the Apollo space program.
Experimental use in automobiles has also been explored.

Advancements:

Significant progress has been made in developing new electrode materials, catalysts, and electrolytes to increase the efficiency of fuel cells.
Various types of fuel cells have been developed and tested for different applications.

Environmental Benefits:

Fuel cells are pollution-free, emitting only water vapor as a byproduct of the electrochemical reaction.
Due to their environmental friendliness and high efficiency, fuel cells hold promise for future energy needs.

Fuel cells represent a cleaner and more efficient alternative to traditional combustion-based power generation methods, making them a crucial technology in the transition to more sustainable energy sources.

Corrosion of Metals

Corrosion is a natural process that gradually forms oxides or other salts on the surfaces of metallic objects. It is a widespread issue that affects various metals, causing damage to buildings, bridges, ships, and other metal-based objects. Here are some key points about corrosion:

Examples of Corrosion:

Rusting of iron
Tarnishing of silver
Development of a green coating on copper and bronze

Consequences of Corrosion:

Significant economic losses due to repair and replacement of corroded objects.
Potential safety hazards, such as bridge collapses or equipment failure, caused by structural damage due to corrosion.

Chemistry of Corrosion:

Corrosion is primarily an electrochemical phenomenon.
At a specific spot on a metal object (anode), oxidation occurs as electrons are lost to oxygen, forming metal ions.
- Anode Reaction: 2Fe(s) → 2Fe2+ + 4e–
Electrons released at the anodic spot move through the metal to another spot (cathode) where reduction of oxygen occurs in the presence of H+ ions.
- Cathode Reaction: O2(g) + 4H+(aq) + 4e– → 2H2O(l)

Overall Reaction:

2Fe(s) + O2(g) + 4H+(aq) → 2Fe2+(aq) + 2H2O(l)
Cell Potential: E(cell) = 1.67 V

Formation of Rust:

Ferrous ions (Fe2+) are further oxidized by atmospheric oxygen to ferric ions (Fe3+), resulting in the formation of hydrated ferric oxide (Fe2O3.xH2O), which is commonly known as rust. This process also produces hydrogen ions.

Prevention of Corrosion:

Prevention of corrosion is essential to save money and prevent accidents.
Methods for preventing corrosion include:
- Coating: Protecting the metal surface from direct contact with the atmosphere by using paint or chemicals like bisphenol.
- Galvanization: Covering the metal surface with a layer of inert metals like tin (Sn) or zinc (Zn).
- Sacrificial Electrode: Providing a sacrificial electrode made of another metal (e.g., magnesium or zinc) that corrodes itself to protect the object.

Corrosion prevention is critical for preserving the longevity and safety of metal-based structures and objects. It involves a range of strategies to minimize the electrochemical reactions that lead to the deterioration of metals.

The Transition to a Hydrogen Economy

Current Energy Source: Fossil fuels like coal, oil, and natural gas are the primary energy sources driving our economy. They have fueled economic growth and development, with per capita energy consumption often used as an indicator of a country’s progress.
Environmental Concerns: The combustion of fossil fuels produces carbon dioxide (CO2), which contributes to the “Greenhouse Effect.” This phenomenon leads to a rise in the Earth’s surface temperature, resulting in the melting of polar ice and a subsequent rise in sea levels. This poses a significant threat to low-lying coastal areas and island nations like the Maldives.
Hydrogen as an Alternative: Hydrogen is considered an ideal alternative energy source because its combustion produces only water (H2O) as a byproduct. This makes hydrogen a renewable and non-polluting energy option.
Hydrogen Production: Hydrogen production can be achieved through the electrolysis of water, a process that uses electricity, often generated from renewable sources like solar or wind, to split water molecules into hydrogen and oxygen. This is a key aspect of the Hydrogen Economy.
Fuel Cells: In addition to hydrogen production, fuel cells play a crucial role in the Hydrogen Economy. Fuel cells use hydrogen as a fuel and, through electrochemical reactions, generate electricity efficiently and cleanly. They are used to power various applications, including vehicles, backup power systems, and more.
Electrochemical Principles: Both the production of hydrogen through water electrolysis and its use in fuel cells are based on electrochemical principles. These technologies leverage the movement of ions and electrons at electrodes to facilitate chemical reactions.

The Vision of the Hydrogen Economy: The Hydrogen Economy envisions a future where hydrogen serves as a versatile and sustainable energy carrier. It offers several advantages:

Clean Energy: Hydrogen combustion produces only water, making it an environmentally friendly alternative to fossil fuels.
Energy Storage: Hydrogen can store excess renewable energy (e.g., from solar and wind sources) for later use.
Transportation: Hydrogen fuel cells can power vehicles, reducing greenhouse gas emissions in the transportation sector.
Grid Balancing: Hydrogen can help balance the electrical grid by providing a backup power source during peak demand or when renewable energy sources are intermittent.

Challenges and Progress: While the Hydrogen Economy holds promise, there are challenges, including the cost-effective production and distribution of hydrogen, as well as the development of efficient fuel cell technologies. Nonetheless, significant research and development efforts are underway to overcome these challenges and advance the use of hydrogen as a sustainable and clean energy source.