A redox reaction, short for reduction-oxidation reaction, is defined by the transfer of electrons between two chemical species. This electron exchange results in a change in the oxidation states of the atoms involved. These reactions drive energy production in biological metabolism, the functioning of modern batteries, and the corrosion of metals. Identifying a redox reaction is a foundational skill for analyzing many chemical changes.
Defining Oxidation and Reduction
The complementary actions of oxidation and reduction must always occur together in a redox process. Oxidation describes the loss of electrons by a chemical species, while reduction refers to the gain of electrons. The mnemonic “OIL RIG” (“Oxidation Is Loss, Reduction Is Gain”) helps remember this relationship.
When an atom loses electrons, its oxidation state becomes more positive. Conversely, when an atom gains electrons, its oxidation state becomes more negative. The species that loses electrons (is oxidized) is called the reducing agent because it causes reduction in the other species. The species that gains electrons (is reduced) is known as the oxidizing agent because it facilitates the oxidation of the other reactant.
Rules for Assigning Oxidation Numbers
To move beyond the simple conceptual definition, chemists use a formal accounting tool called oxidation numbers, or oxidation states. The oxidation number is a hypothetical charge assigned to an atom in a molecule or ion, assuming all bonds are purely ionic. These numbers provide the necessary mathematical framework to track electron shifts.
The assignment of oxidation numbers follows a specific hierarchy of rules:
- An atom in its elemental form (e.g., \(\text{O}_2\) or \(\text{Zn}\)) always has an oxidation number of zero.
- For a monatomic ion (e.g., \(\text{Cl}^-\) or \(\text{Na}^+\)), the oxidation number equals the charge of the ion.
- Fluorine is always assigned an oxidation number of \(-1\) in all its compounds.
- Oxygen is typically assigned \(-2\), except in peroxides where it is \(-1\).
- Hydrogen is assigned \(+1\) when bonded to nonmetals (e.g., \(\text{H}_2\text{O}\)), but \(-1\) when bonded to metals (e.g., \(\text{NaH}\)).
- Group 1 metals (e.g., lithium) are always \(+1\), and Group 2 metals (e.g., magnesium) are always \(+2\).
The final, overarching rule is that the sum of the oxidation numbers for all atoms in a neutral compound must equal zero. If the species is a polyatomic ion, the sum must equal the overall charge of that ion. These rules allow for the calculation of an unknown oxidation number for one element if the numbers for the other elements in the compound are known.
Confirming the Electron Transfer
Once the rules for assigning oxidation numbers are understood, confirming a redox reaction becomes a straightforward comparison exercise. The first step is to methodically assign the oxidation number to every atom in the reactants on the left side of the chemical equation. Then, the oxidation numbers for every atom in the products on the right side must also be determined.
The core of the identification process involves comparing the oxidation numbers of the same element before and after the reaction. If an atom’s oxidation number increased (became more positive), that atom lost electrons and was oxidized. Conversely, if the number decreased (became more negative), that atom gained electrons and was reduced.
If the comparison reveals that at least one atom’s oxidation number increased and another atom’s oxidation number decreased, the reaction is unequivocally classified as a redox reaction. Consider the reaction between solid zinc and copper ions: \(\text{Zn} + \text{Cu}^{2+} \rightarrow \text{Zn}^{2+} + \text{Cu}\).
Initially, zinc metal (\(\text{Zn}\)) has an oxidation number of \(0\), and the copper ion (\(\text{Cu}^{2+}\)) has a number of \(+2\). In the products, the zinc ion (\(\text{Zn}^{2+}\)) has an oxidation number of \(+2\), and the copper metal (\(\text{Cu}\)) has a number of \(0\). Zinc’s number changed from \(0\) to \(+2\), indicating it was oxidized. Copper’s number changed from \(+2\) to \(0\), meaning it was reduced. Since both oxidation and reduction occurred, this reaction is confirmed as a redox process.
Recognizing Non-Redox Reactions
It is helpful to understand reactions that do not involve electron transfer, as these are non-redox reactions. In these cases, the oxidation numbers of all atoms remain unchanged throughout the chemical change. The most common examples are acid-base reactions and double displacement reactions.
In acid-base neutralization (e.g., \(\text{HCl}\) and \(\text{NaOH}\)), ions combine to form a salt and water. The oxidation numbers of all elements remain constant because the ions merely change partners without exchanging electrons.
Double displacement reactions, also known as metathesis reactions, involve two compounds exchanging components to form two new compounds. Since the ions retain their charges and oxidation numbers while swapping partners, no electron transfer occurs, confirming the process is not redox.