The oxidation number of bromine (Br) in HBr is -1.
Understanding Oxidation Numbers
Oxidation numbers, also known as oxidation states, represent the hypothetical charge an atom would have if all its bonds were ionic. They are crucial for understanding chemical reactions, especially in identifying oxidation-reduction (redox) processes. These numbers help us track the distribution of electrons among atoms in a compound.
Determining Oxidation Numbers in HBr
To determine the oxidation number of bromine in hydrogen bromide (HBr), we apply a set of fundamental rules for assigning oxidation numbers:
- Neutral Compound Rule: The sum of the oxidation numbers of all atoms in a neutral compound must equal zero. HBr is a neutral molecule, so the sum of the oxidation numbers of hydrogen and bromine must be 0.
- Hydrogen Rule: Hydrogen typically has an oxidation number of +1 when it is bonded to a non-metal. Bromine is a non-metal, so hydrogen in HBr has an oxidation number of +1.
- Electronegativity Principle: In a covalent bond, the more electronegative atom is assigned the negative oxidation number, assuming it "gains" electrons from the less electronegative atom. In HBr, bromine (Br) is covalently bonded to hydrogen (H). Hydrogen is less electronegative than bromine. This means that the electron pair in the H-Br bond is pulled closer to the more electronegative bromine atom. Consequently, hydrogen is assigned +1, and bromine is assigned -1 to reflect this electron distribution.
Based on these rules:
- Let the oxidation number of Hydrogen (H) be +1.
- Let the oxidation number of Bromine (Br) be x.
- For the neutral compound HBr: (+1) + x = 0
- Solving for x: x = -1
Therefore, the oxidation number of bromine in HBr is -1.
General Rules for Assigning Oxidation Numbers
Understanding common oxidation number rules can help determine them in various compounds. Here's a brief overview:
- Free Elements: An atom in its elemental form (e.g., O₂, Br₂, Na) has an oxidation number of 0.
- Monatomic Ions: The oxidation number of a monatomic ion is equal to its charge (e.g., Cl⁻ is -1, Mg²⁺ is +2).
- Fluorine: Fluorine (F) always has an oxidation number of -1 because it is the most electronegative element.
- Oxygen: Oxygen (O) usually has an oxidation number of -2, except in peroxides (like H₂O₂) where it is -1, or when bonded to fluorine (where it can be positive).
- Alkali Metals: Group 1 elements (Li, Na, K, etc.) always have an oxidation number of +1 in compounds.
- Alkaline Earth Metals: Group 2 elements (Be, Mg, Ca, etc.) always have an oxidation number of +2 in compounds.
- Halogens: Other halogens (Cl, Br, I) usually have an oxidation number of -1 when they are the most electronegative element in a compound or in simple binary compounds, but they can have positive oxidation numbers when bonded to more electronegative elements like oxygen or fluorine.
| Element/Type | Typical Oxidation Number | Notes |
|---|---|---|
| Free elements (e.g., O₂) | 0 | |
| Monatomic ions | Equal to ion's charge | |
| Hydrogen (H) | +1 | -1 in metal hydrides (e.g., NaH) |
| Oxygen (O) | -2 | -1 in peroxides (e.g., H₂O₂) |
| Fluorine (F) | -1 | Always |
| Group 1 metals | +1 | Always in compounds |
| Group 2 metals | +2 | Always in compounds |
| Sum in neutral compound | 0 | |
| Sum in polyatomic ion | Equal to ion's charge |
For more detailed information on oxidation number rules, you can refer to resources on chemical oxidation states.
