Electrolysis is a process in which an external electrical source forces a nonspontaneous redox reaction to occur. It is the reverse of a galvanic cell. Applications include:

• Electroplating: Deposition of a thin metal layer onto a substrate (another metal or nonmetal). For example, silver plating (silvering) a copper surface: a copper object is made the cathode in a silver salt solution (e.g., silver nitrate). When current passes, Ag+ ions are reduced to silver metal on the cathode surface, forming a shiny coating.
• Electrorefining: Purification of metals by making the impure metal the anode in an electrolytic cell. For instance, impure copper is oxidized to Cu2+ at the anode, then pure copper is deposited at the cathode, leaving impurities behind.
• Extraction of metals such as aluminum from alumina (Al2O3) dissolved in molten cryolite: aluminum cations are reduced to aluminum metal at the cathode, while oxide anions are oxidized to oxygen at the carbon anode. This is the industrial Hall-Héroult process.

In all electrolysis applications, one must choose appropriate electrolytes, electrode materials, and cell voltages to drive the desired redox reactions while minimizing side reactions (for instance, hydrogen or oxygen evolution).

Batteries are compact galvanic cells (or assemblies of cells) that convert stored chemical energy into electrical energy on demand. Common types include:

• Primary (nonrechargeable) batteries: e.g., alkaline cells (zinc–manganese dioxide), alkaline-carbon cells, lithium primary cells. They undergo irreversible chemical changes and cannot be recharged.
• Secondary (rechargeable) batteries: e.g., lead-acid batteries (Pb–PbO2), nickel–cadmium (Ni–Cd), nickel–metal hydride (Ni–MH), lithium-ion (Li-ion) batteries. Applying an external current reverses the redox reactions, regenerating the reactants.

Example: Lead-acid battery half-reactions (in 1.2 M sulfuric acid):
- Discharge (galvanic) mode:
- Anode (oxidation): Pb(s) + SO4^2−(aq) → PbSO4(s) + 2 e−
- Cathode (reduction): PbO2(s) + 4 H+(aq) + SO4^2−(aq) + 2 e− → PbSO4(s) + 2 H2O(l)

Corrosion is the undesired oxidation of metals, often resulting in metal oxide or hydroxide formation. Common examples include rusting of iron and tarnishing of silver. Understanding corrosion as a redox process allows strategies for protection:

1. Protective Coatings: Paint, varnish, plastic coatings, or other impermeable layers prevent contact with oxygen and moisture.
2. Galvanization: Coating iron or steel with a more easily oxidized metal (often zinc). Zinc acts as a sacrificial anode: it corrodes preferentially (zinc → Zn2+ + 2 e−), protecting the underlying iron.
3. Cathodic Protection: Attaching a more negative (more easily oxidized) metal block (sacrificial anode) to a steel structure in contact with soil or water. The sacrificial anode dissolves instead of the protected metal.
4. Alloying: Stainless steel (iron alloyed with chromium and nickel) forms a thin, passive chromium oxide layer that prevents further oxidation of iron.
5. Design Considerations: Avoiding crevices where moisture can accumulate, preventing contact between dissimilar metals (which can cause galvanic corrosion), and controlling pH and electrolyte composition.
   Understanding the half-reactions in corrosion (for example, iron oxidation to Fe2+, oxygen reduction to hydroxide) helps in designing inhibitors and protective measures.

Redox titrations are volumetric analyses in which a solution of an oxidizing agent is titrated against a solution of a reducing agent (or vice versa) until the equivalence point, where stoichiometrically equivalent amounts have reacted. Common titrations include:

• Potassium permanganate titration: KMnO4 (a strong oxidizer) titrated against a solution containing Fe2+ (ferrous sulfate) or oxalate ions (C2O42–). Permanganate’s purple color disappears as it is reduced to colorless Mn2+, so the end point is the first persistent pale pink color. Reaction in acidic medium with Fe2+: 5 Fe2+ + MnO4− + 8 H+ → 5 Fe3+ + Mn2+ + 4 H2O
• Dichromate titration: K2Cr2O7 (potassium dichromate) is another common oxidizing titrant. For example, it can titrate Fe2+ in acidic solution by converting orange dichromate (Cr2O72–) to green Cr3+. An external indicator such as diphenylamine sulfonate may be used, since Cr3+ color change is not always sharp.
• Iodometry: Iodine/iodide redox titrations often involve generation of I2 by oxidation of I– with some oxidizer (e.g., Cu2+, ClO–), then titration of the liberated I2 with standardized thiosulfate solution. The end point is detected with starch indicator (blue-black complex with I2 disappears). Key steps for performing a redox titration include:
• Write the balanced redox reaction under the titration conditions (acidic, basic, neutral).
• Determine stoichiometry (moles of electrons exchanged per mole of analyte or titrant).
• Use a suitable indicator or a potentiometric method (measuring potential changes) to identify the end point.
• Calculate the analyte concentration from volume and concentration of titrant used at equivalence.

Living cells depend on controlled redox reactions to convert nutrients into usable energy (ATP). Two principal redox-driven processes are:

• Cellular respiration: Glucose is oxidized stepwise to carbon dioxide, and oxygen is reduced to water. Electrons travel through a chain of redox centers (NADH, flavins, iron-sulfur proteins, cytochromes, copper centers) in the mitochondrial inner membrane. This electron flow pumps protons across the membrane, creating a proton gradient. The return flow of protons through ATP synthase drives the synthesis of ATP from ADP and inorganic phosphate. Overall simplified equation: C6H12O6 + 6 O2 → 6 CO2 + 6 H2O + energy (ATP + heat)
• Photosynthesis: In chloroplasts of plants and photosynthetic bacteria, water is oxidized (releasing O2) while carbon dioxide is reduced to carbohydrate. Light energy excites electrons in chlorophyll, which move through a series of redox carriers, pumping protons and generating a proton gradient. That gradient drives ATP synthesis. Electrons ultimately reduce NADP+ to NADPH, which is used in the Calvin cycle to fix CO2 into sugars. Overall simplified equation: 6 CO2 + 6 H2O + light energy → C6H12O6 + 6 O2
  In both processes, the controlled movement of electrons through protein complexes and cofactors exemplifies redox chemistry in living systems.

Many large-scale industrial operations rely on redox chemistry:

• Hall-Héroult process: Extraction of aluminum metal from alumina (Al2O3) dissolved in molten cryolite. In a large electrolytic cell, aluminum cations gain electrons (are reduced) at the cathode to form molten aluminum metal, while oxide anions are oxidized at carbon anodes to form carbon dioxide.
• Electrolytic refining of copper: Impure copper is cast into an anode. In an electrolytic cell, Cu(s) from the anode is oxidized to Cu2+ and migrates to the cathode to deposit as pure copper. Impurities that are less noble (e.g., iron, zinc) remain in solution; those that are more noble (e.g., silver, gold) collect beneath the anode as anode sludge, which can be further processed.
• Electrolytic production of chlorine: In the chlor-alkali process, brine (sodium chloride solution) is electrolyzed. Chloride ions are oxidized at the anode to chlorine gas; water is reduced at the cathode to hydrogen gas and hydroxide ions. Resulting products are chlorine, hydrogen, and sodium hydroxide (caustic soda).
• Production of nitric acid: Ammonia is oxidized (using air and a platinum catalyst) to nitrogen monoxide, which is further oxidized to nitrogen dioxide, then absorbed in water to yield nitric acid. These are multiple redox steps essential to fertilizer production.
  Industrial redox processes are engineered to maximize yield, minimize energy consumption, manage by-products, and control operational costs.