The oxidation number is a bookkeeping tool used by chemists to track electrons within chemical species, representing a hypothetical charge assigned to an atom. This number describes the degree of electron loss or gain an atom experiences when forming a compound. It is a formal construct, not the atom’s actual electrical charge, but it provides a convenient method for analyzing the distribution of electrons in both ionic and covalent substances. The concept is useful because it allows for a simplified analysis of how atoms change their electronic state during chemical reactions.
Conceptualizing the Oxidation Number
The oxidation number, also known as the oxidation state, is formally defined as the charge an atom would possess if all its bonds were considered entirely ionic. This is an artificial scenario where electrons in every bond are completely transferred to the more electronegative atom. The number can be a positive integer, a negative integer, or zero, indicating a notional loss, gain, or no change in electron possession compared to the neutral element.
A positive oxidation number signifies a hypothetical loss of electrons, while a negative number indicates a hypothetical gain of electrons. This is distinct from an atom’s formal charge, which assumes electrons in a covalent bond are shared equally. The oxidation number assumes the more electronegative atom takes all the shared electrons, making it a powerful tool for understanding electron distribution based on electronegativity differences.
The Priority Rules for Assignment
Assigning oxidation numbers to elements in a compound requires following a set of established rules, which must be applied in a specific priority order. The first rule is that any element in its standard, uncombined state has an oxidation number of zero (e.g., \(\text{O}_2\), \(\text{Na}\)). For monatomic ions, the oxidation number is simply equal to the charge of the ion (e.g., \(+2\) for \(\text{Mg}^{2+}\)).
Fluorine, the most electronegative element, is always assigned an oxidation number of \(-1\) in its compounds. Group 1 and Group 2 metals always have oxidation numbers of \(+1\) and \(+2\), respectively, in their compounds. Hydrogen is usually assigned \(+1\), except when bonded to a metal, forming a metal hydride (e.g., \(\text{NaH}\)), where its oxidation number is \(-1\).
Oxygen is conventionally assigned an oxidation number of \(-2\) in most compounds. Exceptions include peroxides (e.g., \(\text{H}_2\text{O}_2\)), where oxygen is \(-1\), and when bonded to fluorine (e.g., \(\text{OF}_2\)), where its oxidation number is \(+2\). Halogens are generally assigned \(-1\), unless they are bonded to a more electronegative element like oxygen or fluorine.
Calculating Oxidation Numbers in Complex Compounds
The final step in assigning oxidation numbers involves a simple algebraic calculation, necessary when a compound contains an element whose oxidation number is not defined by the priority rules.
For a neutral compound, the sum of the oxidation numbers for all atoms must equal zero. For instance, in \(\text{H}_2\text{SO}_4\), hydrogen is \(+1\) and oxygen is \(-2\). To find the oxidation number of sulfur, represented by \(x\), the equation is \(2(+1) + x + 4(-2) = 0\). Solving this gives \(x = +6\). This algebraic approach allows for the determination of the oxidation state of the unknown atom.
When dealing with a polyatomic ion, the sum of the oxidation numbers of all atoms must equal the overall charge of the ion. Consider the permanganate ion, \(\text{MnO}_4^{-}\), which has an overall charge of \(-1\). With oxygen assigned \(-2\), the calculation to find the oxidation number of manganese, \(y\), is \(y + 4(-2) = -1\). Solving this results in \(y = +7\).
The Utility in Redox Reactions
Oxidation numbers provide the fundamental basis for identifying and understanding reduction-oxidation (redox) reactions. A redox reaction is a chemical process that involves the transfer of electrons between two species, resulting in a change in their oxidation states. Tracking these changes is the primary chemical purpose of the oxidation number concept.
Oxidation is defined as an increase in an atom’s oxidation number, corresponding to a loss of electrons. Conversely, reduction is defined as a decrease in an atom’s oxidation number, which corresponds to a gain of electrons. For example, if a metal goes from \(0\) to \(+2\), it has been oxidized, while a nonmetal going from \(0\) to \(-2\) has been reduced.
By calculating the oxidation numbers of all elements, chemists can determine which species is oxidized and which is reduced. This information forms the foundation for balancing complex redox equations. The change in oxidation numbers dictates the number of electrons transferred, ensuring that the number of electrons lost equals the number gained.