Applications of Redox Reactions

Everyday and Technical Uses of Redox Reactions

Redox reactions are not just an abstract concept; they are among the most widely used reaction types in technology, industry, and living systems. This chapter focuses on where and how redox reactions are used, without re‑deriving the basic redox concepts, oxidation numbers, or balancing methods.

We will look at typical application areas:

• Energy conversion and storage (batteries, fuel cells, combustion)
• Corrosion and corrosion protection
• Industrial production of important substances
• Biological redox processes
• Analytical and household uses of redox chemistry

Energy Conversion and Storage

Combustion Processes

Combustion is a redox process in which a substance (the fuel) is oxidized by an oxidizing agent, usually oxygen from the air.

Examples:

• Combustion of carbon:
  $$\text{C} + \text{O}_2 \rightarrow \text{CO}_2$$
• Combustion of methane (main component of natural gas):
  $$\text{CH}_4 + 2\,\text{O}_2 \rightarrow \text{CO}_2 + 2\,\text{H}_2\text{O}$$

Characteristics:

• Strongly exothermic → heat can be used directly (heating) or to produce mechanical/electrical energy (engines, power plants).
• Often incomplete in practice (e.g. formation of CO and soot), which has environmental and health consequences.

Key idea for applications: in combustion the electron transfer is tightly coupled to energy release. Technical combustion systems are designed to control this energy release efficiently and as cleanly as possible.

Galvanic Cells and Batteries

In galvanic cells (batteries), redox reactions are used to directly convert chemical energy into electrical energy.

General principle:

• Two half-reactions (oxidation and reduction) are spatially separated into two half-cells.
• Electrons flow externally through a wire from the anode (oxidation) to the cathode (reduction).
• Ions move through the electrolyte or a salt bridge to maintain charge balance.
• The cell voltage depends on the involved redox couples and their standard potentials.

Example: Zinc–Carbon (Leclanché) Cell (Simplified)

• Anode (oxidation): zinc is oxidized.
• Cathode (reduction): manganese(IV) oxide is reduced in the presence of ammonium ions.

Overall simplified reaction:
$$\text{Zn} + 2\,\text{MnO}_2 + 2\,\text{NH}_4^+ \rightarrow \text{Zn}^{2+} + \text{Mn}_2\text{O}_3 + 2\,\text{NH}_3 + \text{H}_2\text{O}$$

Properties:

• Primary cell: essentially not rechargeable (irreversible or poorly reversible reactions).
• Used in cheap “dry” batteries for low‑drain devices.

Example: Lead–Acid Battery

Used in car starter batteries.

Main half-reactions during discharge:

• Anode (negative electrode):
  $$\text{Pb} + \text{SO}_4^{2-} \rightarrow \text{PbSO}_4 + 2\,\text{e}^-$$
• Cathode (positive electrode):
  $$\text{PbO}_2 + 4\,\text{H}^+ + \text{SO}_4^{2-} + 2\,\text{e}^- \rightarrow \text{PbSO}_4 + 2\,\text{H}_2\text{O}$$

Overall:
$$\text{Pb} + \text{PbO}_2 + 2\,\text{H}_2\text{SO}_4 \rightarrow 2\,\text{PbSO}_4 + 2\,\text{H}_2\text{O}$$

Key points for applications:

• Rechargeable: by applying an external voltage, the redox reactions are driven in reverse (electrolytic charging process).
• High current capability → suitable for starting engines.

Example: Lithium‑Ion Battery (Principle)

In lithium‑ion cells, lithium ions and electrons move between two host materials (e.g. graphite and a metal oxide).

Simplified view:

• During discharge, Li atoms in the negative electrode are oxidized:
  $$\text{LiC}_6 \rightarrow \text{C}_6 + \text{Li}^+ + \text{e}^-$$
• Li⁺ ions migrate through the electrolyte to the positive electrode, where they are reduced and inserted into a metal oxide host.

Applications:

• Portable electronics, electric vehicles, energy storage.
• High energy density due to large potential differences between the redox couples used.

Across all batteries, design focuses on choosing suitable redox couples, electrolytes, and electrode materials to achieve desired voltage, capacity, lifetime, and safety.

Fuel Cells

Fuel cells, like galvanic cells, convert chemical energy to electrical energy. The difference: reactants are continuously supplied from outside.

Typical example: hydrogen–oxygen fuel cell.

• Anode (H₂ oxidation):
  $$\text{H}_2 \rightarrow 2\,\text{H}^+ + 2\,\text{e}^-$$
• Cathode (O₂ reduction, in acidic medium):
  $$\tfrac{1}{2}\,\text{O}_2 + 2\,\text{H}^+ + 2\,\text{e}^- \rightarrow \text{H}_2\text{O}$$

Overall:
$$\text{H}_2 + \tfrac{1}{2}\,\text{O}_2 \rightarrow \text{H}_2\text{O}$$

Advantages in applications:

• Higher efficiency than internal combustion engines.
• Direct production of electrical power; the only product (for pure H₂) is water.
• Used in space technology, prototypes and some commercial fuel‑cell vehicles, and stationary power systems.

Challenges include hydrogen production and storage and the cost/durability of catalysts.

Corrosion and Corrosion Protection

Corrosion as an Unwanted Redox Process

Corrosion is the undesired oxidation of metals by their environment, often with a corresponding reduction of oxygen or other oxidizing agents.

Example: rusting of iron in moist air:

• Anodic partial reaction (iron oxidation):
  $$\text{Fe} \rightarrow \text{Fe}^{2+} + 2\,\text{e}^-$$
• Cathodic partial reaction (oxygen reduction in neutral/alkaline solution, simplified):
  $$\text{O}_2 + 2\,\text{H}_2\text{O} + 4\,\text{e}^- \rightarrow 4\,\text{OH}^-$$

Resulting iron hydroxides/oxides form “rust”, which usually does not adhere well and does not protect the underlying metal.

Conditions that promote corrosion:

• Presence of water (electrolyte) to conduct ions.
• Presence of dissolved oxygen or other oxidizers.
• Inhomogeneities in the metal (e.g. impurities, stress, different phases) that create local galvanic cells.

Methods of Corrosion Protection

Because corrosion is a redox process, protection methods aim to:

• Prevent electron transfer, or
• Force another, more “sacrificial” material to be oxidized instead.

Common methods:

1. Coatings
• Paint, organic coatings, plastics, or inorganic layers (e.g. oxide layers).
• They physically separate the metal from oxygen and moisture.
• If the coating is damaged, localized corrosion may occur at defects.
2. Alloying and Passivation
• Chromium in stainless steels enables formation of a thin, adherent, protective Cr₂O₃ layer.
• This passive layer slows further oxidation.
3. Sacrificial Anodes (Cathodic Protection)
• A more easily oxidized metal (e.g. Zn or Mg) is electrically connected to the protected metal (e.g. steel hull of a ship, underground pipelines).
• The sacrificial metal acts as the anode and is preferentially oxidized:
  $$\text{Zn} \rightarrow \text{Zn}^{2+} + 2\,\text{e}^-$$
• The protected structure becomes the cathode and is thus reduced (or at least not oxidized).
4. Impressed Current Cathodic Protection
• An external DC power source forces the metal structure to be the cathode by supplying electrons.
• Used for large pipelines, storage tanks, reinforced concrete structures.

Industrial Redox Processes

Many large‑scale industrial processes are redox based. Only some key examples are highlighted here; several are covered in detailed form in the chapter “Selected Chemical Engineering Processes”.

Production of Metals from Ores

Most metals occur in nature as compounds, often as oxides or sulfides. To obtain the pure metal, the metal cation must be reduced.

Example: Iron Production in the Blast Furnace (Simplified)

Ores mainly contain iron(III) oxide.

Stepwise reduction:

1. Carbon monoxide acts as reducing agent:
   $$\text{Fe}_2\text{O}_3 + 3\,\text{CO} \rightarrow 2\,\text{Fe} + 3\,\text{CO}_2$$
2. Carbon (coke) can also directly reduce iron oxides at high temperatures.

Coal (coke) itself is oxidized (to CO and CO₂), providing both energy and the reducing agent.

Example: Aluminum Production by Electrolysis of Molten Alumina

Aluminum oxide is very stable; chemical reduction with carbon is impractical. Instead, electrolytic reduction in molten salt (Hall–Héroult process) is used.

At the cathode (reduction):
$$\text{Al}^{3+} + 3\,\text{e}^- \rightarrow \text{Al}$$

The process consumes large amounts of electrical energy because a large potential must be applied to drive the redox reaction.

Industrial Oxidations and Reductions of Inorganic Substances

Examples:

• Chlor‑alkali electrolysis:
• Oxidation at the anode (e.g. in brine):
  $\,\text{Cl}^- \rightarrow \text{Cl}_2 + 2\,\text{e}^-$$
• Reduction at the cathode:
  $\,\text{H}_2\text{O} + 2\,\text{e}^- \rightarrow \text{H}_2 + 2\,\text{OH}^-$$
• Products: chlorine gas, hydrogen gas, sodium hydroxide solution.
• Production of nitric acid (Ostwald process):
• Ammonia is catalytically oxidized:
  $\,\text{NH}_3 + 5\,\text{O}_2 \rightarrow 4\,\text{NO} + 6\,\text{H}_2\text{O}$$
• NO is further oxidized to NO₂, which is absorbed in water to form nitric acid.
• These are multi‑step redox processes controlled by catalysts and reaction conditions.

Such processes illustrate how technical chemistry uses redox reactions to build large‑scale industrial value chains.

Biological Redox Processes

In living organisms, redox reactions are central to energy metabolism and biosynthesis. Biological systems use specialized molecules and enzymes to control and couple redox processes.

Cellular Respiration (Aerobic)

In aerobic respiration, organic substrates (e.g. glucose) are oxidized, and oxygen is reduced to water. The released energy is captured in the form of ATP.

Global equation for glucose respiration:
$$\text{C}_6\text{H}_{12}\text{O}_6 + 6\,\text{O}_2 \rightarrow 6\,\text{CO}_2 + 6\,\text{H}_2\text{O}$$

Underlying redox principles:

• Carbon atoms in glucose are stepwise oxidized (increase in oxidation number).
• Oxygen atoms in O₂ are reduced (from 0 to −2).
• Electron carriers such as NAD⁺/NADH, FAD/FADH₂ mediate electron transfer.
• In the electron transport chain, electrons move “downhill” in redox potential, and the energy released drives the synthesis of ATP.

Although the chemistry is complex, it follows the same redox logic as in technical fuel cells.

Photosynthesis

Photosynthesis is essentially the reverse redox process of respiration. Light energy is used to reduce CO₂ and oxidize water.

Overall simplified equation:
$$6\,\text{CO}_2 + 6\,\text{H}_2\text{O} \xrightarrow{\text{light}} \text{C}_6\text{H}_{12}\text{O}_6 + 6\,\text{O}_2$$

Key redox aspect:

• Water is oxidized to O₂, releasing electrons.
• CO₂ is reduced to carbohydrates.
• Light‑absorbing pigments (e.g. chlorophyll) and redox chains direct electron flow.

Thus, redox reactions connect the solar energy input to the chemical energy storage in biomass.

Analytical and Everyday Uses of Redox Reactions

Redox Titrations

In analytical chemistry, redox reactions are used to determine the concentration of oxidizing or reducing agents.

Typical features:

• A titrant with known concentration undergoes a well‑defined redox reaction with the analyte.
• The equivalence point can often be detected via:
• Color change of the analyte (e.g. permanganate self‑indicator),
• Added redox indicator,
• Electrochemical detection (change in potential).

Examples:

• Determination of iron(II) with potassium permanganate:
  $\,\text{Fe}^{2+} + \text{MnO}_4^- + 8\,\text{H}^+ \rightarrow 5\,\text{Fe}^{3+} + \text{Mn}^{2+} + 4\,\text{H}_2\text{O}$$
• Iodometric methods (using iodine/iodide redox couple) to determine oxidizing agents.

Redox titrations are widely used because many analytes can be either oxidized or reduced in a selective and stoichiometrically clear way.

Disinfectants and Bleaching Agents

Many disinfectants and bleaches kill microorganisms or decolorize dyes by oxidizing key components.

Examples:

• Hydrogen peroxide ($\text{H}_2\text{O}_2$):
• Can act as oxidizing or reducing agent, depending on conditions.
• Used as disinfectant, bleaching agent (e.g. hair, textiles), and in some environmental treatment processes.
• Sodium hypochlorite (in bleach solutions):
• Oxidizes organic dyes and biomolecules, leading to discoloration and disinfection.
• Redox reactions involve species such as $\text{ClO}^-$ and $\text{Cl}_2$.
• Ozone ($\text{O}_3$):
• Strong oxidizing agent for water treatment and disinfection.

Mechanistic details vary, but in all cases redox reactions chemically transform or destroy target molecules.

Redox Reactions in Everyday Life

Some common phenomena and products based on redox chemistry:

• “Self‑heating” packs:
• Use exothermic oxidation of metals (e.g. Fe → Fe₂O₃) to produce heat.
• Breathalyzers (older types):
• Oxidation of ethanol in exhaled air is linked to a measurable color or current change.
• Food preservation:
• Ascorbic acid (vitamin C) and sulfites act as antioxidants by undergoing oxidation themselves and thereby preventing or slowing oxidation of food components.
• Photography (traditional silver halide):
• Exposure and development involve redox reactions of silver ions and reducing agents.

Summary of Application Principles

Across this diversity of uses, some common redox principles appear:

• Energy conversion: Redox couples with large potential differences can provide useful electrical or thermal energy (batteries, fuel cells, combustion, respiration).
• Material transformation: Redox processes allow extraction and refining of metals, synthesis of chemicals, and controlled decomposition or modification of substances.
• Protection and control: Understanding redox allows the design of corrosion protection systems, antioxidants, and passivation strategies.
• Analysis and monitoring: Redox reactions form the basis of many quantitative analytical techniques.

Recognizing redox processes in these applications helps connect the abstract concept of electron transfer to many practical and technological contexts.