Bromine is a chemical element with the symbol Br and atomic number 35. It belongs to the halogen group (Group 17) in the periodic table and is known for its reddish-brown liquid form at room temperature. Among the halogens, bromine is the only element that exists as a liquid under standard conditions. One of its most important physical properties is its boiling point, which is crucial for various chemical and industrial applications.
What Is the Boiling Point of Bromine?
The boiling point of bromine is 58.8°C (137.8°F or 332 K). This means that at 58.8 degrees Celsius, bromine transitions from a liquid to a gas under normal atmospheric pressure (1 atm).
Bromine has a relatively low boiling point compared to most other nonmetals and halogens, making it volatile and easily evaporates at room temperature. Its high vapor pressure gives bromine its characteristic strong odor and toxicity when inhaled.
Understanding the Boiling Point of Bromine
1. Why Does Bromine Have a Low Boiling Point?
The boiling point of any substance depends on the intermolecular forces between its molecules. Bromine exists as a diatomic molecule (Br₂) and is held together by:
- Van der Waals forces (London dispersion forces) – These are weak intermolecular attractions that increase with molecular size.
- No hydrogen bonding or strong dipole interactions – Unlike water or ammonia, bromine lacks hydrogen bonds, which contribute to its lower boiling point.
Although bromine is a relatively large molecule, its intermolecular forces are still weaker compared to ionic or strongly polar compounds, resulting in a low boiling point.
2. Comparison of Bromine’s Boiling Point With Other Halogens
Bromine is part of the halogen group, which includes fluorine (F), chlorine (Cl), iodine (I), and astatine (At). The boiling points of halogens follow a trend based on molecular weight and intermolecular forces:
| Halogen | Chemical Symbol | Boiling Point (°C) | Physical State at Room Temperature |
|---|---|---|---|
| Fluorine | F₂ | -188.1°C | Gas |
| Chlorine | Cl₂ | -34.0°C | Gas |
| Bromine | Br₂ | 58.8°C | Liquid |
| Iodine | I₂ | 184.3°C | Solid |
| Astatine | At₂ | ~300°C (estimated) | Solid |
As seen in the table, the boiling point increases down the group because larger molecules have stronger dispersion forces. Bromine is the only halogen that is liquid at room temperature, while chlorine and fluorine exist as gases, and iodine and astatine are solids.
3. Factors Affecting the Boiling Point of Bromine
Several factors influence bromine’s boiling point, including:
- Atmospheric Pressure: At lower pressures, bromine will boil at a lower temperature. Conversely, at higher pressures, its boiling point increases.
- Purity of the Substance: Impurities can alter the boiling point slightly, though bromine is typically studied in its pure diatomic form (Br₂).
- Intermolecular Forces: The strength of London dispersion forces determines how much energy is needed to break molecular attractions.
Applications of Bromine’s Boiling Point in Industry
1. Chemical Synthesis and Industrial Use
Bromine’s boiling point and volatility make it useful in:
- Pesticide and flame retardant production
- Pharmaceuticals and chemical intermediates
- Dyes, solvents, and water treatment
2. Laboratory and Experimental Use
- Evaporation studies: Scientists study bromine’s volatility to understand halogen behavior.
- Reagent in organic reactions: Its liquid state at room temperature makes it convenient for chemical reactions.
3. Storage and Handling Considerations
Since bromine is volatile and toxic, it requires proper storage:
- Sealed containers to prevent evaporation.
- Cool and well-ventilated areas to avoid vapor accumulation.
- Protective equipment when handling due to its corrosive nature.
Bromine has a boiling point of 58.8°C, which makes it volatile and easily converted into gas at slightly elevated temperatures. Compared to other halogens, it stands out as the only liquid at room temperature. Understanding its boiling point is essential for its safe handling, industrial applications, and chemical reactions.