Table of Contents
Overview of Redox Reactions
Redox reactions are chemical reactions in which electrons are transferred between particles. In such processes, the oxidation state (oxidation number) of at least one element changes. Because of electron transfer, redox reactions are closely linked to energy conversion (for example in batteries, corrosion, respiration, and combustion).
Redox reactions always consist of two inseparable partial processes:
- an oxidation (electron loss)
- a reduction (electron gain)
No electron can disappear or appear out of nowhere, so oxidation and reduction always occur together in a coupled way.
In this chapter, the general nature, appearance, and significance of redox reactions are introduced; details such as the formal determination of oxidation numbers, the systematic balancing of redox equations, and redox potentials are treated in later chapters.
Typical Characteristics of Redox Reactions
Redox reactions show one or more of the following characteristics:
- Change in oxidation state of elements
At least one element changes its oxidation state between educts and products.
Example (metal and nonmetal):
$$
\text{Zn} + \text{Cu}^{2+} \rightarrow \text{Zn}^{2+} + \text{Cu}
$$ - Zn: $0 \rightarrow +2$ (oxidation)
- Cu: $+2 \rightarrow 0$ (reduction)
- Electron transfer (often visible in ionic equations)
Simplified ionic representation highlights electron transfer:
$$
\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^- \quad (\text{oxidation})
$$
$$
\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu} \quad (\text{reduction})
$$ - Energy conversion
Many redox reactions release large amounts of energy (often as heat or electrical energy). Combustion reactions and electrochemical cells are key examples. - Change in composition or color
Because different oxidation states of an element often have different colors and reactivities, redox reactions can be recognized by color changes or formation of new substances.
Oxidation and Reduction – Modern View
Historically, oxidation and reduction were defined via oxygen uptake and loss. Today, they are defined via electron transfer or changes in oxidation number:
- Oxidation
- Loss of electrons
- Increase in oxidation number of an element
- Reduction
- Gain of electrons
- Decrease in oxidation number of an element
A simple example is the reaction of metallic magnesium with oxygen:
$$
2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}
$$
Formally:
- Mg: $0 \rightarrow +2$ (each Mg atom gives up two electrons → oxidation)
- O: $0 \rightarrow -2$ (each O atom takes up two electrons → reduction)
Oxidizing and Reducing Agents
Because oxidation and reduction are always coupled, substances come in pairs:
- Reducing agent
- Donates electrons to another species
- Is itself oxidized
- Its oxidation state increases
- Oxidizing agent
- Accepts electrons from another species
- Is itself reduced
- Its oxidation state decreases
In the magnesium–oxygen reaction:
- Mg is the reducing agent (it donates electrons and is oxidized).
- O$_2$ is the oxidizing agent (it accepts electrons and is reduced).
In the zinc–copper ion reaction:
- Zn is the reducing agent.
- Cu$^{2+}$ is the oxidizing agent.
A helpful rule of thumb:
- “Oxidizing agent” causes oxidation of the other substance, but is itself reduced.
- “Reducing agent” causes reduction of the other substance, but is itself oxidized.
Redox Reactions in Different Contexts
Redox reactions can take place in different media and have different appearances:
Redox Reactions in Aqueous Solutions
In aqueous solution, ions are present and redox reactions often occur via ionic species. For example, an iron nail in copper sulfate solution:
$$
\text{Fe} + \text{CuSO}_4 \rightarrow \text{FeSO}_4 + \text{Cu}
$$
Ionic form (simplified):
$$
\text{Fe} + \text{Cu}^{2+} \rightarrow \text{Fe}^{2+} + \text{Cu}
$$
- Fe is oxidized ($0 \rightarrow +2$), Fe is the reducing agent.
- Cu$^{2+}$ is reduced ($+2 \rightarrow 0$), Cu$^{2+}$ is the oxidizing agent.
Typical features:
- Metal surface dissolves or is coated with another metal.
- Solution color changes (e.g. blue Cu$^{2+}$ fades).
Redox Reactions in Gaseous Phase and Combustion
Many combustion processes are redox reactions between a fuel and oxygen:
- Burning of carbon:
$$
\text{C} + \text{O}_2 \rightarrow \text{CO}_2
$$ - Burning of methane:
$$
\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}
$$
Here, the carbon and hydrogen atoms are oxidized, and oxygen is reduced. Energy is released as heat and light.
Redox Reactions in Electrochemical Systems
In electrochemical cells, redox reactions are spatially separated into two half-cells, and the electrons flow through an external circuit. This makes it possible to convert chemical energy into electrical energy (or vice versa).
The underlying processes are redox reactions, for example:
- At the anode: oxidation (electron release)
- At the cathode: reduction (electron uptake)
The details of construction and analysis of such cells are treated in electrochemistry, but the central chemistry is always a redox process.
Biochemical Redox Reactions (Overview)
In biological systems, many central metabolic pathways are based on redox reactions. Instead of simple metal ions, special organic molecules (coenzymes such as NAD$^+$/NADH, FAD/FADH$_2$) serve as electron carriers.
Typical biochemical redox process:
- Nutrients are oxidized (e.g. glucose to CO$_2$).
- Oxygen is reduced to water.
- The released energy is stored in biochemical “energy carriers”.
The mechanistic and energetic aspects are dealt with in detail in biochemical chapters, but they rest on the same redox principles described here.
Recognizing Redox Reactions
To decide whether a reaction is a redox reaction, you can use the following approach:
- Assign oxidation numbers to all atoms in educts and products.
- Compare oxidation numbers for each element.
- If at least one element shows a change in oxidation number, the reaction is a redox reaction.
Examples:
- Reaction of hydrogen with chlorine:
$$
\text{H}_2 + \text{Cl}_2 \rightarrow 2\text{HCl}
$$ - H: $0 \rightarrow +1$ (oxidation)
- Cl: $0 \rightarrow -1$ (reduction)
⇒ Redox reaction. - Neutralization of hydrochloric acid with sodium hydroxide:
$$
\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}
$$
Oxidation numbers remain unchanged for all elements.
⇒ Not a redox reaction; it is mainly an acid–base reaction.
Redox Pairs and Half-Reactions
For every redox process, you can write two conceptual “half-reactions”:
- Oxidation half-reaction: shows the electron loss.
- Reduction half-reaction: shows the electron gain.
Each half-reaction involves a redox pair (also called conjugate redox pair), for example:
- Zn/Zn$^{2+}$
- Fe$^{2+}$/Fe$^{3+}$
- Cl$_2$/Cl$^-$
Using the zinc–copper reaction again:
- Oxidation:
$$
\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-
$$ - Reduction:
$$
\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}
$$
Redox pairs are useful for:
- Separating complex redox processes into simpler parts.
- Comparing tendencies to be oxidized or reduced.
- Describing redox equilibria and cell voltages (treated later).
Types of Redox Reactions
Depending on the substances involved and the pattern of electron transfer, several types of redox reactions can be distinguished.
Direct Electron Transfer Between Two Different Substances
This is the “classical” redox case, such as:
- Metal + metal ion:
$$
\text{Zn} + \text{Cu}^{2+} \rightarrow \text{Zn}^{2+} + \text{Cu}
$$ - Nonmetal + metal:
$$
2\text{Na} + \text{Cl}_2 \rightarrow 2\text{NaCl}
$$
Combustion and Oxidation by Oxygen
Many oxidation processes involve gaseous O$_2$ as oxidizing agent:
- Metals:
$$
2\text{Fe} + \frac{3}{2}\text{O}_2 \rightarrow \text{Fe}_2\text{O}_3
$$ - Organic compounds:
$$
\text{C}_3\text{H}_8 + 5\text{O}_2 \rightarrow 3\text{CO}_2 + 4\text{H}_2\text{O}
$$
Oxygen-containing oxidizing agents (e.g. permanganate, dichromate, peroxides) also belong to this group, although their detailed mechanisms can be more complex.
Disproportionation and Comproportionation (Overview)
In some redox reactions, only one element changes its oxidation state in two directions at once:
- Disproportionation:
One species is simultaneously oxidized and reduced, forming two products with different oxidation states of the same element.
Example:
$$
2\text{H}_2\text{O}_2 \rightarrow 2\text{H}_2\text{O} + \text{O}_2
$$
Oxygen in H$_2$O$_2$ has oxidation number $-1$.
In H$_2$O: $-2$ (reduction), in O$_2$: $0$ (oxidation).
- Comproportionation (or synproportionation):
Two species with different oxidation states react to form a product with an intermediate oxidation state.
These special forms of redox reactions illustrate the central role of oxidation numbers; exact treatment and balancing appear in later chapters.
Importance of Redox Reactions in Everyday Life and Technology
Redox reactions play a fundamental role in many areas of daily life and technology:
- Corrosion of metals
Rusting of iron is a redox process in which iron is oxidized (often by oxygen and water).
Example (simplified):
$$
4\text{Fe} + 3\text{O}_2 + 6\text{H}_2\text{O} \rightarrow 4\text{Fe(OH)}_3
$$ - Batteries and accumulators
In galvanic cells, redox reactions provide electrical energy.
Different materials act as reducing and oxidizing agents at the electrodes. - Combustion processes and energy supply
Burning fossil fuels, running engines, and many heating systems are based on redox reactions between fuels and oxygen. - Bleaching and disinfection
Many bleaches and disinfectants (e.g. chlorine, hypochlorite, hydrogen peroxide) act via redox reactions that oxidize dyes or damage microorganisms. - Biological energy conversion
Cellular respiration and photosynthesis are based on cascade-like redox processes that convert energy and store it in chemically usable form.
Understanding redox reactions thus provides insight not only into abstract chemical theory, but also into numerous practical applications in technology, environment, and life sciences.