Simulation · Chemistry · Class 11
Oxidation numbers that add up
From the lesson Oxidation number and its rules in Redox Reactions. Change the values and watch what happens.
The idea behind it
NCERT §7.3
- In reactions between covalent molecules, such as H₂ + Cl₂ → 2HCl, no electron is fully transferred; the shared pair only shifts toward the more electronegative atom. To treat such reactions as redox anyway, chemists assign each atom an oxidation number.
- Oxidation number is the charge an atom would carry in a compound if every bonding pair were handed entirely to the more electronegative partner.
- Rule 1: an element in the free state has oxidation number 0, whatever its form (H₂, O₂, Cl₂, O₃, S₈, P₄, Na, Mg, Al).
- Rule 2: a monatomic ion has an oxidation number equal to its charge. In their compounds alkali metals are always +1, alkaline earth metals +2 and aluminium +3.
- Rule 3: oxygen is usually −2, but −1 in peroxides (H₂O₂, Na₂O₂), −½ in superoxides (KO₂, RbO₂), +2 in OF₂ and +1 in O₂F₂. It takes positive values only when bonded to fluorine.
- Rule 4: hydrogen is usually +1, but −1 in binary hydrides of active metals such as LiH, NaH and CaH₂.
- Rule 5: fluorine is −1 in all its compounds. Chlorine, bromine and iodine are −1 in halides but take positive values when combined with oxygen, as in oxoacids and oxoanions.
- Rule 6: the oxidation numbers of all atoms in a neutral compound add to zero; in a polyatomic ion they add to the ion's charge. In CO₃²⁻, x + 3(−2) = −2 gives carbon +4.
- For main-group elements the highest oxidation number usually equals the group number for groups 1 and 2, and the group number minus 10 for groups 13 to 17. Across period 3 it rises from +1 in NaCl to +7 in HClO₄.
More simulations in Redox Reactions
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