Redox reactions

Nature of electron transitions and historical context

The formation of ionic and polar bonds is always driven by the difference in electronegativity between interacting atoms. Electron density is redistributed: some atoms receive an excess of electrons and a negative charge, while others experience a deficit and a positive charge. A striking example is the synthesis of sodium chloride from simple substances. The electronegativity of sodium is only $0.9$, and that of chlorine is $3.0$. The valence electron from the $3s$ sublevel of sodium transitions to the $3p$ sublevel of chlorine. Sodium is oxidized by losing an electron, and chlorine is reduced. These changes are recorded in electron balance schemes.

The terms themselves came into chemistry from metallurgy. Initially, oxidation referred to the combination of a pure metal with oxygen (which was accompanied by the loss of electrons), and reduction referred to obtaining the starting metal from its oxide (with the acceptance of electrons). Later, these concepts were extended to absolutely any processes involving the donation and acceptance of electrons.

Formal and partial charges

In molecular compounds, unlike ionic crystals, the displacement of bonding electrons is small. They merely spend more time near the more electronegative atom, creating partial charges (denoted as $+\delta$ and $-\delta$). For example, in a hydrogen chloride ($HCl$) molecule, the actual charge of the chlorine atom is $-0.17$, and that of hydrogen is $+0.17$ (which is about 17% of the full charge of an electron).

However, for the convenience of reaction analysis, the true values of partial charges are ignored. Instead, the oxidation state (OS) is used — a conventional formal charge calculated under the assumption that all bonds in a substance are completely ionic. The oxidation state shows the number of electrons formally transferred to a more electronegative element (then the OS is positive) or accepted from a less electronegative one (OS is negative).

Differences from other reactions and relationship with valence

It is extremely important to distinguish redox reactions from acid-base interactions. In acid-base reactions, an electron pair is also transferred from a donor to an acceptor. However, the donor atom has a higher electronegativity, so the transferred electron pair continues to belong to it by more than half. As a result, the oxidation state of the elements in such reactions does not change.

In terms of absolute value, the oxidation state often coincides with valence, but exceptions exist. If an atom is bonded simultaneously to more and less electronegative neighbors, or if there are bonds between identical atoms, the values may diverge. For example, in potassium disulfide ($K_2S_2$), there is a bond between sulfur atoms ($S-S$). Sulfur is divalent here, but its oxidation state is $-1$, since the bond between identical elements does not affect it.

Rules for calculating oxidation state

To calculate oxidation states in complex substances, a strict algorithm and values that are constant for a number of elements are used:

• In simple substances (where bonds are between identical atoms), the OS is always equal to 0.
• Fluorine ($F$) in all compounds has an OS of -1. Other halogens ($Cl$, $Br$, $I$) also typically exhibit an OS of -1 unless bonded to oxygen.
• Oxygen ($O$) almost always exhibits an OS of -2. Exceptions are peroxides with an $–O–O–$ bond (OS is -1$), oxygen difluoride $OF_2$ (OS is +2), and superoxides.
• Hydrogen ($H$) paired with most non-metals yields +1 (except for silicon), and with metals and silicon — -1.
• Metals of group IA have an OS of +1, group IIA elements and zinc — +2, aluminum — +3.

In any polyatomic particle, the algebraic sum of the oxidation states of all atoms is always exactly equal to the net charge of that particle (for example, in the monatomic ion $Cr^{3+}$, the oxidation state of chromium is $+3$).

Redox potentials and practical significance

The strength of various oxidizing and reducing agents can be compared using tables of standard electrode potentials ($E^\circ$, measured in volts), but only under strictly identical environmental conditions. For example, in an acidic medium, the strongest oxidizing agent from the standard list is gaseous fluorine ($F_2$, potential $+2.866$ V), while the hydrogen electrode ($H^+$) is taken as the zero reference point.

In nature and technology, these processes are of fundamental importance. Metallurgy, fuel combustion for energy generation, and the synthesis of inorganic and organic substances are based on electron transfer. In the physiology of higher plants and animals, the most important oxidizing agent is oxygen, which ensures the respiration process that proceeds with the release of energy necessary for life (an exothermic process).

Briefly: answers to topical questions

Class and exam questions

What else to study on this topic

• Drawing up electron balance schemes
• Calculation of oxidation state in complex acid residue anions
• Effect of acidic and alkaline media on the strength of redox couples
• Features of oxidation states in peroxides and superoxides
• Table of standard electrode potentials

Other topics in "Chapter 7. Classification of chemical reactions by electronic processes"

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