Bromine (Br)

Introduction: Why Bromine Matters

Bromine (Br) is a fascinating and reactive non-metallic element, unique among non-metals for being a liquid at standard temperature and pressure. Its reddish-brown, fuming liquid state, and pungent odor are distinctive. As a member of the halogen group (Group 17), bromine exhibits characteristic high reactivity, strong oxidizing properties, and a tendency to form ionic bromides or covalent compounds. Its versatile chemistry underpins its widespread applications in various industries, making it a crucial element to understand for competitive examinations.

CBSE/JEE Quick Revision Notes

Symbol: Br
Atomic Number: 35
Atomic Mass: 79.904 u
Group: 17 (Halogens)
Period: 4
Block: p-block
State at Room Temperature: Reddish-brown liquid with an irritating odor.
Electronic Configuration: [Ar] 3d¹⁰ 4s² 4p⁵
Valency: -1 (most common), can exhibit +1, +3, +5, +7 in compounds with more electronegative elements (e.g., in oxoacids or interhalogens).
Electronegativity (Pauling scale): 2.96
Electron Affinity: 324.6 kJ/mol
First Ionization Enthalpy: 1139.9 kJ/mol
Standard Electrode Potential (E° Br₂/Br⁻): +1.07 V (indicates its oxidizing power)
Oxidizing Nature: Strong oxidizing agent, but less than F₂ and Cl₂.

Electron Configuration & Bonding Behavior

Bromine’s electron configuration, [Ar] 3d¹⁰ 4s² 4p⁵, reveals seven valence electrons. This configuration drives its chemical behavior:

Electron Gain: Bromine readily gains one electron to achieve a stable octet, forming the bromide ion (Br⁻). This is the most common and stable oxidation state (-1) for bromine.
Covalent Bonding: Due to its high electronegativity, bromine forms covalent bonds with less electronegative elements (e.g., in HBr, CBr₄). In elemental form, it exists as a diatomic molecule, Br₂, held by a single covalent bond.
Variable Oxidation States: When bonding with more electronegative elements like oxygen or fluorine, bromine can exhibit positive oxidation states (+1, +3, +5, +7) by expanding its octet using empty d-orbitals. Examples include BrF₃, HBrO₃ (bromic acid), and HBrO₄ (perbromic acid).
Oxidizing Agent: Bromine acts as an oxidizing agent by accepting electrons, converting itself into bromide ions. Its oxidizing strength is lower than fluorine and chlorine but higher than iodine.

Crucial Chemical Reactions

1. Reaction with Metals

Bromine reacts with many metals to form ionic bromides.

With Sodium: 2Na(s) + Br₂(l) → 2NaBr(s)
With Magnesium: Mg(s) + Br₂(l) → MgBr₂(s)

2. Reaction with Non-metals

With Hydrogen: Hydrogen reacts with bromine to form hydrogen bromide, requiring heating or light. H₂(g) + Br₂(l) → 2HBr(g)
With Phosphorus (e.g., PCl₃/PCl₅ analogs): 2P(s) + 3Br₂(l) → 2PBr₃(l) (Phosphorus tribromide)

3. Displacement Reactions

Bromine, being more reactive than iodine, can displace iodide ions from their salts. However, it is displaced by more reactive halogens like fluorine and chlorine.

Displacement of Iodide: Br₂(aq) + 2I⁻(aq) → 2Br⁻(aq) + I₂(aq)
Displacement by Fluorine: F₂(g) + 2Br⁻(aq) → 2F⁻(aq) + Br₂(aq)

4. Reaction with Water

Bromine dissolves sparingly in water to form bromine water, undergoing disproportionation to form hydrobromic acid and hypobromous acid. This reaction is reversible.

Br₂(l) + H₂O(l) ⇌ HBr(aq) + HOBr(aq)

5. Reaction with Alkalis (NaOH)

Bromine reacts differently with cold, dilute alkalis versus hot, concentrated alkalis, exhibiting disproportionation.

Cold, dilute NaOH: Br₂(l) + 2NaOH(aq) → NaBr(aq) + NaOBr(aq) + H₂O(l) (Sodium bromide and Sodium hypobromite)
Hot, concentrated NaOH: 3Br₂(l) + 6NaOH(aq) → 5NaBr(aq) + NaBrO₃(aq) + 3H₂O(l) (Sodium bromide and Sodium bromate)

6. Reaction with Organic Compounds

Addition to Alkenes (Test for Unsaturation): Bromine water (reddish-brown) is decolorized when added to an alkene, indicating the presence of a carbon-carbon double bond. CH₂=CH₂(g) + Br₂(l) → BrCH₂-CH₂Br(l) (1,2-Dibromoethane)
Electrophilic Substitution in Aromatics: In the presence of a Lewis acid catalyst (e.g., FeBr₃), bromine substitutes hydrogen on an aromatic ring. C₆H₆(l) + Br₂(l) --(FeBr₃)--> C₆H₅Br(l) + HBr(g) (Bromobenzene)

Industrial and Biological Importance

Industrial Importance

Flame Retardants: Brominated organic compounds are extensively used as flame retardants in plastics, textiles, and electronics, reducing flammability.
Agriculture: Bromine compounds serve as pesticides and soil fumigants. Methyl bromide (CH₃Br) was historically used but is being phased out due to ozone depletion concerns.
Pharmaceuticals: Historically, sodium bromide (NaBr) was used as a sedative. Bromine derivatives are still used in some pharmaceutical syntheses.
Photography: Silver bromide (AgBr) is a light-sensitive compound crucial in traditional photographic films and papers due to its ability to undergo photolysis.
Dyes: Bromine is used in the synthesis of certain dyes, including derivatives of indigo.
Water Treatment: Bromine compounds, like bromochlorodimethylhydantoin, are used as disinfectants in swimming pools and hot tubs, especially where chlorine is not preferred.

Biological Importance

Marine Organisms: Bromine is an essential trace element for certain marine organisms, particularly algae and sponges, where it’s incorporated into various organobromine compounds.
Enzymatic Activity: Some enzymes, such as vanadium bromoperoxidase found in marine life, utilize bromine in their catalytic mechanisms.
Human Biology: Trace amounts of bromide ions are present in the human body. Recent research indicates a potential role for bromide in the formation of collagen IV, a structural component of basement membranes, suggesting a more fundamental role than previously thought.

Bromine (Br) - Properties, Reactions & Uses