Batteries and corrosion Knowledge-Based Study

Study notes

1. A battery is a portable, self-contained electrochemical source consisting of one or more galvanic cells connected in series that convert chemical energy directly into electrical energy through spontaneous redox reactions.

   Classification of Batteries

   Commercial batteries are broadly categorized based on their ability to be recharged:

   • Primary Batteries: These are non-rechargeable cells. The electrochemical redox reactions taking place within the cell are irreversible. Once the cell runs down because the reactants are largely converted to products, the battery is discarded. Classic examples include the zinc-carbon (Leclanché) cell and the alkaline battery. In an alkaline battery, zinc powder serves as the anode, manganese dioxide (MnO2) acts as the cathode, and a basic electrolyte (potassium hydroxide, KOH) is used, offering longer operational life and stable voltage.
   • Secondary Batteries: These are rechargeable cells. The cell reactions can be reversed by passing an external electric current in the opposite direction, regenerating the original active materials. Common examples include the lead-acid storage battery, nickel-cadmium (Ni-Cd), and lithium-ion (Li-ion) batteries.

   Secondary batteries are critical for automotive, industrial, and consumer electronics applications due to their high energy density and rechargeability.

   The Lead-Acid Storage Battery

   Commonly found in automobiles, a typical 12 V car battery consists of six 2 V galvanic cells connected in series.

   • Anode: Spongy lead (Pb)
   • Cathode: Lead(IV) oxide (PbO2)
   • Electrolyte: Aqueous sulfuric acid solution (H2SO4, approximately 38% by mass)
   • Discharge Reaction: Pb(s) + PbO2(s) + 2H2SO4(aq) → 2PbSO4(s) + 2H2O(l)

   During discharge, both electrodes are converted into insoluble lead(II) sulfate (PbSO4), and sulfuric acid is consumed. During recharging via an external power supply, these processes are reversed.

   Lithium-Ion Batteries

   Lithium-ion cells power modern portable electronics and electric vehicles. They operate via the intercalation and de-intercalation of lithium ions (Li+) shuttling between a graphite anode and a transition metal oxide cathode (such as LiCoO2). Because lithium is the lightest metal and possesses a very negative standard reduction potential, lithium-ion batteries provide exceptionally high energy densities.

   A fuel cell is an electrochemical cell that converts the chemical energy of a continuously supplied fuel (such as hydrogen, methane, or methanol) and an oxidant directly into electrical energy.

   Characteristics of Fuel Cells

   • Unlike conventional batteries, fuel cells do not store chemical energy; they operate continuously as long as fuel and oxidant are continuously fed into the system.
   • They do not burn fuel directly, allowing them to bypass the thermodynamic efficiency limits of Carnot heat engines.
   • The most widely utilized system is the hydrogen-oxygen fuel cell.

   Hydrogen-Oxygen Fuel Cell Chemistry

   In an alkaline hydrogen-oxygen fuel cell:

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

   Because the only chemical byproduct is water, hydrogen fuel cells represent a remarkably clean, high-efficiency energy technology.

   Corrosion is the spontaneous deterioration of a metal resulting from an electrochemical reaction with its environment. The most familiar example is the rusting of iron in the presence of oxygen and moisture.

   Electrochemical Steps of Rusting

   Rusting operates through micro-galvanic cells established on the surface of iron:

   1. Anodic Region: A region of the iron surface with mechanical stress or impurities acts as an anode, where iron undergoes oxidation: Fe(s) → Fe2+(aq) + 2e−
   2. Cathodic Region: The electrons released migrate through the metal to another site exposed to atmospheric oxygen and water, where oxygen reduction occurs: O2(g) + 4H+(aq) + 4e− → 2H2O(l)
   3. Precipitation and Rust Formation: The generated Fe2+ ions migrate through the electrolyte film, undergo further oxidation by dissolved oxygen to Fe3+, and precipitate as hydrated iron(III) oxide: 4Fe2+(aq) + O2(g) + 4H2O(l) + 2xH2O(l) → 2Fe2O3·xH2O(s) + 8H+(aq)

   Unlike aluminum oxide, which forms a dense, passivating layer that halts further attack, rust is porous and flakes off, continually exposing fresh iron to corrosion.

   Preventing corrosion is essential for preserving industrial infrastructure, pipelines, vessels, and transportation systems.

   Methods of Corrosion Control

   • Protective Surface Barriers: Applying paint, grease, plastic, or impermeable coatings prevents oxygen and moisture from reaching the metal substrate.
   • Galvanization: Coating steel or iron with a thin layer of zinc. Zinc serves two roles: it creates a protective barrier, and if scratched, zinc acts as a sacrificial anode because its standard reduction potential is more negative than that of iron.
   • Cathodic Protection: A technique used to protect buried pipelines, storage tanks, and ship hulls. The structure to be protected is made the cathode of an electrochemical cell. This can be achieved by:
     • Sacrificial Anode (Galvanic Protection): Connecting the iron structure to a more active, easily oxidized metal such as magnesium (Mg) or zinc (Zn). The sacrificial metal oxidizes preferentially, providing electrons to the iron and keeping it reduced.
     • Impressed Current Cathodic Protection: Supplying electrons continuously from an external direct current (DC) power source to force the protected metal to remain strictly cathodic.