Some General Rules For Determining Oxidation States Without Use of Lewis Structures
Here are general rules for simple compounds without structural formulae:
- Any pure element (even if it forms diatomic molecules like chlorine, Cl2) has an oxidation state (OS) of zero. Examples of this are Cu or O2.
- For monatomic ions, the OS is the same as the charge of the ion. For example, S2− has an OS of −2, whereas Li+ has an OS of +1.
- The sum of OSs for all atoms in a molecule or polyatomic ion is equal to the charge of the molecule or ion, so that the OS of one element can be calculated from the OS of the other elements. For example, in SO32− (sulfite ion), the total charge of the ion is −2, and each oxygen is assumed to have its usual oxidation state of −2. The sum of OSs is then OS(S) + 3(−2) = −2, so that OS(S) = +4.
To sum up: The algebraic sum of oxidation states of all atoms in a neutral molecule must be zero, while in polyatomic ions the algebraic sum of the oxidation states of the constituent atoms must be equal to the charge on the ion. This fact, combined with the fact that some elements almost always have certain oxidation states (due to their very high electropositivity or electronegativity), allows one to compute the oxidation states for the remaining atoms (such as transition metals) in simple compounds.
The following rules that are used for initially assigning oxidation states for certain elements, in simple compounds:
- Fluorine has an oxidation state of −1 when bonded to any other element, since it has the highest electronegativity of all reactive elements.
- Halogens other than fluorine have an oxidation state of −1 except when they are bonded to oxygen, nitrogen, or another (more electronegative) halogen. For example, the oxidation state of chlorine in chlorine monofluoride (ClF) is +1. However, in chlorine bromide (or bromine chloride) (BrCl) the oxidation state of Cl is −1.
- Hydrogen has an oxidation state of +1 except when bonded to more electropositive elements such as sodium, aluminium, and boron, as in NaH, NaBH4, LiAlH4, where each H has an oxidation state of −1.
- In compounds, oxygen typically has an oxidation state of −2, though there are exceptions that are listed below. An example is peroxides (e.g. hydrogen peroxide H2O2) when it has an OS of −1.
- Alkali metals have an oxidation state of +1 in virtually all of their compounds (exception, see alkalide).
- Alkaline earth metals have an oxidation state of +2 in virtually all of their compounds.
Example for a complex salt: In Cr(OH)3, oxygen has an oxidation state of −2 (no fluorine or O–O bonds present), and hydrogen has a state of +1 (bonded to oxygen). So, each of the three hydroxide groups has an oxidation state of −2 + 1 = −1. As the compound is neutral, Cr has an oxidation state of +3.
For molecules with inequivalent atoms of the same element, the algebraic sum method gives only an average oxidation state. We will consider below how to find the oxidation state of each atom with the help of a Lewis structure.
Read more about this topic: Oxidation State
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