Electrochemistry

Definition

A Redox reaction is a combination of two simultaneous processes: Oxidation and Reduction. Electrochemistry is entirely based on these reactions.

Explanation

• Oxidation: The process of loss of electrons. It increases the oxidation state of the element.
• Reduction: The process of gain of electrons. It decreases the oxidation state of the element.
• Agents: The substance that gets oxidized is the Reducing Agent, while the substance that gets reduced is the Oxidizing Agent.

Examples

• $\text{Zn} \to \text{Zn}^{2+} + 2\text{e}^-$ (Zinc is oxidized)
• $\text{Cu}^{2+} + 2\text{e}^- \to \text{Cu}$ (Copper is reduced)

Real-Life Use

Rusting of Iron: When iron ($\text{Fe}$) is exposed to oxygen and moisture, it loses electrons (oxidation) to oxygen, creating hydrated iron oxide (rust).

Definition

Conductance ($G$) is the ease with which electric current flows through an electrolytic solution. It is the mathematical reciprocal of resistance ($R$).

Explanation

Imagine ions as little boats carrying electricity across a river. More ions (boats) moving faster result in higher conductance.

• Unit: Siemens ($\text{S}$ or $\Omega^{-1}$).
• Conductivity ($\kappa$): The conductance of $1\text{cm}^3$ of a solution.
• Factor: Conductance increases as the concentration of ions increases.

Examples

• Pure distilled water: Very low conductance due to lack of ions.
• Saltwater: High conductance because $\text{Na}^+$ and $\text{Cl}^-$ ions are free to move.

Real-Life Use

Water Quality Testers: Digital TDS (Total Dissolved Solids) meters measure conductance to determine the purity of swimming pool or drinking water.

Definition

Molar Conductivity ($\Lambda_m$): The conducting power of all the ions produced by dissolving one mole of an electrolyte in a solution.

Kohlrausch's Law

At infinite dilution, each ion makes a definite contribution to the total molar conductivity, regardless of the other ion it is paired with.

Examples

• For $\text{NaCl}$: $\Lambda^0_m(\text{NaCl}) = \lambda^0(\text{Na}^+) + \lambda^0(\text{Cl}^-)$
• For $\text{KCl}$: $\Lambda^0_m(\text{KCl}) = \lambda^0(\text{K}^+) + \lambda^0(\text{Cl}^-)$

Real-Life Use

Determining Solubility: Used to find the solubility of "sparingly soluble salts" by calculating molar conductivity in saturated solutions.

Definition

Electrolysis is the process of using a direct current (DC) to drive a non-spontaneous chemical reaction, breaking down a compound into its elements.

Faraday's Laws

• First Law: Mass ($m$) of substance deposited $\propto$ quantity of electricity passed.
  Formula: $m = ZIt$ (where $Z$ is electrochemical equivalent).
• Second Law: When the same amount of electricity is passed through different electrolytes, the mass deposited is proportional to their chemical equivalent weights.

Examples

• Electrolysis of molten $\text{NaCl}$ to produce pure Sodium metal.
• Electroplating a copper spoon with silver.

Real-Life Use

Aluminum Production: The Hall-Héroult Process uses massive electricity to electrolyze alumina ($\text{Al}_2\text{O}_3$) in molten cryolite.

Key Concepts

• Galvanic/Voltaic Cell: Converts chemical energy $\to$ electrical energy.
  • Anode: Site of Oxidation (Negative).
  • Cathode: Site of Reduction (Positive).
• EMF: Maximum potential difference between electrodes when no current flows.
• Standard Electrode Potential ($E^\circ$): Potential measured at $1\text{M}, 298\text{K}, 1\text{atm}$.

The Nernst Equation

Used to calculate electrode potential under non-standard conditions:

Examples

• Daniel Cell: Uses $\text{Zn}$ and $\text{Cu}$ electrodes.
• SHE: Standard Hydrogen Electrode (assigned $0\text{V}$ as reference).

Real-Life Use

Glucose Monitors: Use the relationship between current/voltage and glucose concentration (internally applying Nernst-like principles) to measure blood sugar.

The Thermodynamic Relation

• $\Delta G$: Change in Gibbs Free Energy.
• $n$: Moles of electrons transferred.
• $F$: Faraday's constant ($\approx 96485\text{ C}$).

Explanation

If $E_{\text{cell}}$ is positive, $\Delta G$ is negative. A negative $\Delta G$ indicates a spontaneous reaction, meaning the battery will function on its own.

Real-Life Use

Predicting Feasibility: Scientists use this to determine if new battery chemistries (e.g., Sodium-ion) will actually produce electricity.

Classification

• Primary Batteries: Non-rechargeable (Irreversible). Example: Dry cell (Zinc-Carbon).
• Secondary Batteries: Rechargeable (Reversible). Example: Lithium-ion, Lead-acid.
• Fuel Cells: Produce electricity from a fuel (like $\text{H}_2$) reacting with an oxidant ($\text{O}_2$). Example: $\text{H}_2\text{-O}_2$ fuel cell.

Real-Life Use

Space Missions: NASA used $\text{H}_2\text{-O}_2$ fuel cells in the Apollo program because they are lightweight and the only byproduct is drinkable water.

Definition

Corrosion is the gradual destruction of metals by chemical reaction with the environment (unwanted electrochemistry).

Mechanism

The metal surface acts as a tiny electrochemical cell where some spots act as anodes (metal loss) and some as cathodes.

Real-Life Use

Galvanization: Coating iron with Zinc. Zinc acts as a sacrificial anode, oxidizing first to protect the underlying iron from rust.

Key Formulas

Galvanic vs. Electrolytic Cells

In designing a modern Electric Vehicle (EV) battery, engineers apply every core concept of electrochemistry:

• Selection: Lithium-ion is used because Lithium has a high Standard Electrode Potential ($E^\circ$), providing high voltage and low weight.
• The Reaction: A Redox reaction occurs where $\text{Li}^+$ ions move from the anode (graphite) to the cathode (metal oxide) during discharge.
• Voltage Drop: The Nernst Equation explains why voltage decreases as ion concentration changes, causing the battery to "die" slowly.
• Management: Battery Management Systems (BMS) use Faraday's Laws to control current ($I$) and time ($t$) to prevent overcharging and dendrite formation.
• Efficiency: By optimizing molar conductivity of the electrolyte, internal resistance ($R$) is lowered, reducing heat waste and improving acceleration.

1. Which electrode is where oxidation occurs?
   a) Cathode b) Anode c) Salt bridge d) None
   Ans: b
2. Unit of molar conductivity is:
   a) $\text{S cm}^2\text{mol}^{-1}$ b) $\text{S cm}^{-1}$ c) $\Omega\text{ cm}$ d) $\text{S mol}^{-1}$
   Ans: a
3. The value of one Faraday is approximately:
   a) $96500\text{ C}$ b) $86500\text{ C}$ c) $9650\text{ C}$ d) $1.6 \times 10^{-19}\text{ C}$
   Ans: a
4. In a Galvanic cell, the anode is:
   a) Positive b) Negative c) Neutral d) Variable
   Ans: b
5. Kohlrausch's Law is used to calculate:
   a) EMF b) Molar conductivity at infinite dilution c) Resistance d) Gibbs energy
   Ans: b
6. The relation $\Delta G = -nFE_{\text{cell}}$ implies that for a spontaneous reaction:
   a) $E_{\text{cell}}$ must be negative b) $\Delta G$ must be positive c) $E_{\text{cell}}$ must be positive d) $n$ must be zero
   Ans: c
7. A secondary battery is one that:
   a) Cannot be recharged b) Can be recharged c) Uses only fuel d) Is used in watches
   Ans: b
8. Corrosion of iron is an example of:
   a) Reduction b) Electrochemical process c) Displacement reaction d) Non-spontaneous reaction
   Ans: b
9. The reference electrode used to measure electrode potentials is:
   a) Standard Zinc Electrode b) Standard Hydrogen Electrode c) Copper Electrode d) Lead Electrode
   Ans: b
10. Which of the following is a fuel cell?
    a) Dry cell b) Lead accumulator c) $\text{H}_2\text{-O}_2$ cell d) Mercury cell
    Ans: c

1. Q: Define Molar Conductivity. How does it vary with dilution?
   Answer: Molar conductivity ($\Lambda_m$) is the conducting power of all ions produced by one mole of electrolyte. With dilution, $\Lambda_m$ increases because ions move more freely (decreased inter-ionic attraction) and, in weak electrolytes, the degree of dissociation increases.
2. Q: State Faraday's First Law of Electrolysis.
   Answer: It states that the mass ($m$) of any substance deposited or liberated at any electrode is directly proportional to the quantity of electricity ($Q = It$) passed through the electrolyte. $m = ZIt$.
3. Q: What is the purpose of a salt bridge in a Galvanic cell?
   Answer: A salt bridge completes the electrical circuit and maintains electrical neutrality in both half-cells by allowing the migration of ions, preventing charge accumulation that would stop the reaction.
4. Q: Explain the difference between a primary and a secondary battery.
   Answer: Primary batteries (e.g., Dry cell) are non-rechargeable because the chemical reaction is irreversible. Secondary batteries (e.g., Lead-acid) are rechargeable because the reaction can be reversed by passing current in the opposite direction.
5. Q: Write the Nernst equation for a general electrode reaction $M^{n+} + ne^- \to M$.
   Answer: $E = E^\circ - \frac{0.0591}{n} \log \frac{1}{[M^{n+}]}$ (at $298\text{K}$).
6. Q: Why is Zinc used for galvanizing iron?
   Answer: Zinc is more reactive than iron (higher oxidation potential). It acts as a sacrificial anode, oxidizing itself to protect the iron from rusting even if the coating is scratched.
7. Q: What is a fuel cell? Give one advantage.
   Answer: A fuel cell is a galvanic cell that derives energy from the combustion of a fuel (like $\text{H}_2$) in the presence of an oxidant ($\text{O}_2$). Advantage: They are highly efficient and pollution-free (water is the only byproduct).
8. Q: How is $\Delta G$ related to the cell potential? What does a negative $\Delta G$ signify?
   Answer: $\Delta G = -nFE_{\text{cell}}$. A negative $\Delta G$ signifies that the chemical reaction is thermodynamically spontaneous and the cell will produce electrical energy.
9. Q: What happens to the conductivity of a strong electrolyte on dilution?
   Answer: The conductivity ($\kappa$) decreases because the number of ions per unit volume decreases. However, the molar conductivity ($\Lambda_m$) increases.
10. Q: Distinguish between Conductivity and Molar Conductivity.
    Answer: Conductivity ($\kappa$) is the conductance of a $1\text{cm}^3$ solution. Molar Conductivity ($\Lambda_m$) is the conductance of a solution containing $1$ mole of electrolyte.

• Redox: Oxidation = Loss of $e^-$, Reduction = Gain of $e^-$.
• Galvanic: Chemical $\to$ Electrical (Spontaneous).
• Electrolytic: Electrical $\to$ Chemical (Non-spontaneous).
• $\Delta G$ vs $E$ : $\Delta G$ negative $\iff$ $E$ positive.

Mnemonics:

• OIL RIG: Oxidation Is Loss, Reduction Is Gain.
• LEO the lion says GER: Lose Electrons Oxidation, Gain Electrons Reduction.
• AN OX: Anode = Oxidation.
• CAT RED: Cathode = Reduction.