Electrochemical Cells, Batteries, Fuel Cells & Corrosion – Comprehensive Study Notes

UNIT-III OVERVIEW – ELECTROCHEMICAL CELLS & CORROSION

• Focus: In-depth exploration of the interconversion between chemical and electrical energy, covering fundamental principles, practical electrochemical devices, the mechanisms of metallic degradation (corrosion) and its control, and the application of spectroscopic quantification (Beer–Lambert Law) in analytical chemistry.

• Core components: The unit progresses from foundational redox reaction fundamentals to the thermodynamics of electrochemical cells, then delves into various batteries & fuel cells as energy conversion devices, followed by the degradation of metals (corrosion), and finally, protective strategies against corrosion.

FUNDAMENTAL REDOX CONCEPTS

• Oxidation = loss of electrons (e⁻); Reduction = gain of electrons (e⁻).

Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain).

• Oxidizing agent: The species that causes oxidation by accepting electrons (itself undergoing reduction).

• Reducing agent: The species that causes reduction by donating electrons (itself undergoing oxidation).

• Oxidation and reduction are always interdependent and simultaneous processes; their combination constitutes an overall redox reaction.

• Example 1 (Zn | Cu²⁺ system, common in Daniell cell):

  • Oxidation (at anode): A zinc atom loses two electrons to form a zinc ion: $Zn\;(s) \rightarrow Zn^{2+}\;(aq)+2e^{-}$

  • Reduction (at cathode): A cupric ion gains two electrons to form a copper atom: $Cu^{2+}\;(aq)+2e^{-} \rightarrow Cu\;(s)$

  • Net redox reaction: Overall transfer of electrons from zinc to copper ions, resulting in deposition of copper and dissolution of zinc: $Zn\;(s)+Cu^{2+}\;(aq) \rightarrow Zn^{2+}\;(aq)+Cu\;(s)$

• Example 2 (dichromate / ferrous system, often used in titrations):

  • Reduction (dichromate ion gaining electrons in acidic medium): $(Cr2O7)^{2-}+14H^{+}+6e^{-}\rightarrow2Cr^{3+}+7H_2O$

  • Oxidation (ferrous ion losing electrons): $6Fe^{2+}\rightarrow6Fe^{3+}+6e^{-}$

  • Net redox reaction: Formation of chromic and ferric ions: $(Cr2O7)^{2-}+6Fe^{2+}+14H^{+}\rightarrow2Cr^{3+}+6Fe^{3+}+7H_2O$

ELECTROCHEMICAL CELLS

• Definition: A sophisticated device designed to convert chemical energy stored in a spontaneous redox reaction directly into electrical energy. Conversely, some electrochemical cells (electrolytic cells) use electrical energy to drive non-spontaneous reactions.

Synonyms: Galvanic cell or Voltaic cell (specifically for spontaneous reactions).

• Anode – The electrode where oxidation occurs; it is the source of electrons and is typically negatively charged in a galvanic cell.

• Cathode – The electrode where reduction occurs; it consumes electrons and is typically positively charged in a galvanic cell.

• Cell notation (IUPAC convention): A shorthand representation of an electrochemical cell's components and phase boundaries: $\text{Anode} \vert \text{Anodic electrolyte} \Vert \text{Cathodic electrolyte} \vert \text{Cathode}$

  • A single vertical bar (vert) represents a phase boundary (e.g., between a solid electrode and its solution).

  • A double vertical bar (VertVert) denotes a salt bridge, which connects the two half-cells and facilitates ion flow to maintain electrical neutrality.

DANIELL (Zn–Cu) GALVANIC CELL

• Construction: Consists of a zinc rod immersed in a ZnSO4. These two half-cells are connected externally by a metallic wire and internally by a salt bridge (often a U-tube filled with a gel containing KCl or KNO₃).

• Reactions:

  • Anode (Oxidation): Zinc metal dissolves, releasing electrons: $Zn\;(s) \rightarrow Zn^{2+}+2e^{-}$

  • Cathode (Reduction): Copper ions in solution gain electrons and deposit as copper metal: $Cu^{2+}+2e^{-} \rightarrow Cu\;(s)$

• Electron flow: Electrons travel from the zinc anode (site of oxidation) through the external circuit (wire) to the copper cathode (site of reduction).

• Conventional current: Flows in the opposite direction, from cathode (Cu) to anode (Zn).

• Salt bridge (e.g., KCl / KNO₃ gel): Crucial for maintaining electrical neutrality in both half-cells. As $Zn^{2+}$ ions accumulate at the anode and $Cu^{2+}$ ions are consumed at the cathode, the salt bridge allows anions (e.g., $Cl^{-}$) to migrate towards the anode and cations (e.g., $K^{+}$) towards the cathode, preventing charge imbalance and allowing continuous electron flow.

• Cell EMF (standard electrode potential): For the Daniell cell under standard conditions (1 M concentrations, 298 K), the standard electromotive force is approximately $E^{\circ}_{cell}=1.10\;V$, indicating its strong spontaneity.

ELECTRODE & STANDARD ELECTRODE POTENTIAL

• Electrode potential (E): This is the inherent tendency of an electrode (a metal in contact with its own ions) to either lose electrons (undergo oxidation) or gain electrons (undergo reduction) when it is in contact with its corresponding ions in solution.

  • A highly negative charge on the metal indicates a strong tendency for oxidation (electron release).

  • A highly positive charge indicates a strong tendency for reduction (electron gain).

• Standard Electrode Potential (E⁰): The electrode potential measured under standard conditions:

  • Ion concentration of $1\;M$ for dissolved species.

  • Temperature of $298\;K$ ($25\;^{\circ}C$).

  • Partial pressure of $1\;atm$ ($101.325\;kPa$) for any gases involved (e.g., in gas electrodes like SHE).

• $Relationship: The standard reduction potential ($E{red}$) of a half-reaction is equal in magnitude but opposite in sign to its standard oxidation potential ($ E{ox}$):$E{red}= -E{ox}$. For example, if$Cu^{2+}/Cu$has a reduction potential of$+0.34\;V$, its oxidation potentia$l (for $Cu \rightarrow Cu^{2+}$ $) would be$-0.34\;V$.$

• Measurement vs Standard Hydrogen Electrode (SHE): Electrode potentials are not measured in isolation but always relative to a reference electrode. The Standard Hydrogen Electrode (SHE), assigned a standard potential of $E^{\circ}=0\;V$, is the primary reference. It consists of a platinum electrode immersed in a $1\;M$ $H^{+}$ solution, with hydrogen gas at $1\;atm$ bubbled over it.

  • If the test electrode undergoes reduction when connected to SHE, its potential is assigned a positive value.

  • If the test electrode undergoes oxidation when connected to SHE, its potential is assigned a negative value.

• Other reference electrodes: Besides SHE, a widely used practical alternative is the Calomel electrode (Hg/Hg₂Cl₂ | KCl), which offers portability and ease of use in laboratory settings.

NERNST EQUATION & EMF CALCULATIONS

• General form (at 298 K): The Nernst equation quantitatively describes how the electrode potential (E) of a half-cell or the cell potential ($E_{cell}$) of an electrochemical cell deviates from its standard potential ($E^{\circ}$) under non-standard conditions (i.e., when concentrations or partial pressures are not $1\;M$ or $1\;atm$).

  $E = E^{\circ} - \frac{0.0592}{n}\log\,Q$
  where:

  • $Q$ = reaction quotient, defined as the ratio of product concentrations/pressures to reactant concentrations/pressures, each raised to the power of their stoichiometric coefficients. For a generic reaction $aA + bB \rightleftharpoons cC + dD$, $Q = \frac{[C]^c[D]^d}{[A]^a[B]^b}$. Solids and pure liquids are excluded from Q.

  • $n$ = number of electrons transferred in the balanced half-reaction or overall cell reaction.

• For complete cell: The overall cell potential

• Examples:

  • Daniell cell (non-standard conditions): The Nernst equation can predict the cell potential when $[Zn^{2+}]$ and $[Cu^{2+}]$ are not $1\;M$:
    $E = 1.10\,V - \frac{0.0592}{2}\log\frac{[Zn^{2+}]}{[Cu^{2+}]}$

• Potentiometric measurement: To accurately measure the EMF of a cell without altering the concentrations of the species or drawing significant current (which would change the potential), a high-impedance device like a voltmeter or potentiometer is used. This ensures minimal current is drawn, preventing any significant electrochemical reaction during measurement.

• Relationship to equilibrium: At equilibrium, $E_{cell} = 0$ and $Q=K$ (equilibrium constant). The Nernst equation then becomes: $<br>0 = E^{\circ} - \frac{0.0592}{n}\log\,K \Rightarrow \log K = \frac{nE^{\circ}}{0.0592}$. This equation allows calculation of the equilibrium constant from standard cell potentials.

CONCENTRATION CELLS

• Definition: Electrochemical cells uniquely constructed with identical electrode materials and the same redox couple in both half-cells, but differing either in the ion concentrations in their electrolytes or the gas pressures. These cells generate EMF solely due to the difference in concentration, striving to equalize the concentrations.

• EMF purely concentration-driven: The potential difference arises from the tendency of ions to move from a region of higher concentration to lower concentration, mimicking diffusion but driven by electron flow.

  • For a concentration cell at 298 K:
    $E = -\frac{0.0592}{n}\log\frac{[\text{low conc.}]}{[\text{high conc.}]}$

  • The electrode in the lower concentration solution typically acts as the anode (oxidation occurs to increase ion concentration), and the electrode in the higher concentration solution acts as the cathode (reduction occurs to decrease ion concentration).

• Illustrates direct application of Nernst & principle behind pH glass electrode: The operation of pH meters, particularly using pH glass electrodes, is a practical application of concentration cells. The potential developed across the glass membrane is proportional to the difference in $H^{+}$ ion concentration inside and outside the electrode, allowing direct measurement of pH.

ELECTROCHEMICAL vs GALVANIC SERIES

• Electrochemical series: A theoretical or thermodynamic ordering of elements (primarily metals) based on their standard reduction potentials are stronger reducing agents and are more prone to oxidation/corrosion.

  • Predicts displacement tendencies: A metal higher in the series will displace (reduce) the ions of a metal lower in the series from its solution (e.g., Zn will displace Cu from $CuSO_4$).

• Galvanic series: An empirical ranking of metals and alloys (like stainless steel, brass, bronze) based on their actual measured potentials when immersed in a specific real-world environment (e.g., seawater, tap water, soil). This series is a significantly better predictor of real corrosion behavior because it accounts for practical factors such as:

  • Formation of protective oxide films (passivation).

  • Presence of impurities in alloys.

  • Environmental factors like pH, temperature, and aeration.

  • Metallurgical interactions between different phases within an alloy.

• Deviations: There are notable differences between the two series. For example, passive metals like Aluminum (Al) and Titanium (Ti) appear much nobler (less prone to corrosion) in the galvanic series than their standard electrode potentials (from the electrochemical series) would suggest. This is due to the formation of tenacious, self-healing oxide layers that protect the underlying metal from further attack.

BATTERIES

• Battery: A sophisticated device comprising one or more electrochemical cells connected in series or parallel to provide a continuous source of direct current (DC) at a relatively constant voltage. It is a portable energy storage and conversion system.

• Classification:

  • Primary (non-rechargeable) batteries: These batteries are designed for single use. Their electrochemical reactions are essentially irreversible, meaning once the reactants are consumed or products formed, the battery cannot be effectively recharged by applying an external current. Examples include the Dry cell (Leclanché cell) and the Mercury cell, commonly used in low-drain devices.

  • Secondary (rechargeable) batteries: These batteries are capable of undergoing reversible electrochemical reactions. They can be discharged (providing power) and then recharged numerous times by applying an external DC current, which reverses the chemical reactions and regenerates the reactants. Examples include the Lead-acid battery (automotive), Nickel–Cadmium (Ni–Cd) battery, and Lithium-ion (Li-ion) battery (portable electronics, EVs).

  • Fuel cell / Flow battery: These are distinct from conventional batteries as they are