Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), and Astatine (At) are the five elements that make up the halogen group of the periodic table. In Greek, 'halos' means sea salts, ‘genes’ means born. Except for astatine, which is radioactive and only exists in trace amounts, these elements react with metals to produce salts and are plentiful in nature. Because the elements' valence shells have seven electrons, the typical electronic configuration of Br is $\mathrm{[Ar]\:4s^2 \:4p^4}$.
Bromine is the 47th most prevalent element in the earth's crust, with a concentration of 2.5 ppm by weight. It's mostly found as bromides in seawater, mineral springs, and deposits. It was discovered in 1826 by A J Balardin, who used chlorine to purify the mother liquid left behind after the crystallisation of NaCl from sea water. Due to its unpleasant odour, he dubbed the liquid as bromine (Greek bromos, stink).
$$\mathrm{Cl_2 + 2Br^−\:\rightarrow\:2Cl^− + Br_2}$$
A stream of air removes the released bromine, which is then passed through a solution of $\mathrm{Na_2CO_3}$ to produce a mixture of $\mathrm{NaBr}$ and $\mathrm{NaBrO_3}$.
$$\mathrm{3Br_2 + 3Na_2CO_3\:\rightarrow\:5NaBr + NaBrO_3 + 3CO_2}$$
To obtain pure bromine, the solution is acidified and distilled.
$$\mathrm{5NaBr + NaBrO_3 + 3H_2SO_4\:\rightarrow\:3Na_2SO_4 + 5HBr + HBrO_3}$$
$$\mathrm{5HBr + HBrO_3 \:\rightarrow\:3Br_2 + 3H_2O}$$
The physical properties of Bromine and the other halogen congeners can be compared using the table below:
| Property | F | Cl | Br | I |
|---|---|---|---|---|
| Colour (in liquid) | Clear Yellow | Amber Yellow | Reddish Brown | Dark Solid |
| Colour (in vapour) | Pale Yellow | Greenish Yellow | Orange Red | Violet |
| Enthalpy of Fusion (kJ/mol) | 0.51 | 6.41 | 10.57 | 15.52 |
| Enthalpy of Vapourization (kJ/mol) | 6.54 | 20.41 | 29.56 | 41.95 |
| Standard Electrode Potential (V) | + 2.87 | + 1.36 | + 1.06 | + 0.53 |
| Bond Energy (kJ/mol) | 158.8 | 242.6 | 192.8 | 151.1 |
| Density $\mathrm{(g/cm^{-3})}$ | 1.513 | 1.655 | 3.187 | 3.960 |
| Boiling Point (K) | 85 | 239 | 333 | 458 |
| Melting Point (K) | 54 | 172 | 266 | 387 |
| Electronegativity | 4.0 | 3.0 | 2.8 | 2.5 |
| Electron Affinity (kJ/mol) | 333 | 349 | 325 | 296 |
| Ionisation Energy (kJ/mol) | 1681 | 1256 | 1143 | 1009 |
| Ionic Size (pm) | 133 | 184 | 196 | 220 |
| Atomic Size (pm) | 72 | 99 | 114 | 133 |
Bromine is less reactive than chlorine but more active than iodine, and it reacts with a wide range of metals, non-metals, and compounds. For oxidising, bleaching, and disinfection, bromine aqueous solutions are utilised. The following are some instances of its oxidising properties:
$$\mathrm{Na_2S_2O_3 + Br_2 + H_2O\:\rightarrow\:Na_2SO_4 + S + 2HBr}$$
$$\mathrm{Na_3AsO_3 + Br_2 + H_2O\:\rightarrow\:Na_3AsO_4 + 2HBr}$$
Bromine forms hypo brominates when it interacts with a cold NaOH solution. Bromates are formed when hot solutions are used.
$$\mathrm{Br_2 + 2NaOH (cold)\:\rightarrow\:NaBr + NaOBr + H_2O}$$
$$\mathrm{3Br_2 + 6NaOH (hot)\:\rightarrow\: 5NaBr + NaBrO_3 + 3H_2O}$$
Hydrogen Bromide: It is made by passing a hydrogen and bromine combination through an electrically heated platinum spiral. Conc. $\mathrm{H_2SO_4}$ treatment of metal bromides does not yield HBr because the resultant HBr is oxidised to $\mathrm{Br_2}$ by conc. $\mathrm{H_2SO_4}$
Oxides of Bromine: Many of the oxides formed by Bromine decompose even below room temperature. Some important oxides of Bromine are:
Bromine Monoxide $\mathrm{(Br_2O)}$: Bromine vapours react with dry HgO at 323–343 K to produce it.
Bromine Dioxide $\mathrm{(BrO_2)}$: Bromine is ozonolyzed at a low temperature to make Bromine Dioxide
Bromine Trioxide $\mathrm{(BrO_3)}$: It is made via ozonolysis of bromine at temperatures ranging from –5 to 10 degrees Celsius.
Oxyacids of Bromine: Hypobromous acid (HOBr) and bromic acid $\mathrm{(HBrO_3)}$ are two oxyacids formed by bromine. Bromites, on the other hand, are bromous acid salts $\mathrm{(HBrO_2)}$. Bromine water is treated with newly precipitated mercuric oxide or silver oxide to produce hypobromous acid $\mathrm{(HOBr)}$.
It is employed as an oxidizer as well as in the production of disinfectants and other key organic compounds like tetraethyl lead.
In photography, silver bromide is utilised.
As a laboratory reagent, HBr is employed.
Bromine oxides and oxyacids are utilised as oxidising agents.
Bromine bears an atomic number of 35 in the periodic table. It has a characteristic odour and is fairly abundant in the earth’s crust. The unique property of Bromine would be that it is only non-metal present in liquid state at room temperature and ambient pressure. Bromine has ample uses in daily life and in industries.
Q1. What are some of the properties of Bromine?
Ans: It is a reddish brown liquid with a boiling point of 58.8°C and a freezing point of –7.3°C. It's very flammable and produces powerful fumes that irritate the throat and lungs. Bromine is less reactive than chlorine but more active than iodine, and it reacts with a wide variety of metals, nonmetals, and compounds.
Q2. How would you produce HBr in the laboratory?
Ans: $\mathrm{HBr}$ can be produced by the action of conc. $\mathrm{H_3PO_4}$ on metal bromides. The chemical reaction involved is
$$\mathrm{3NaBr + H_3PO_4 \:\xrightarrow{\Delta}\:Na_3PO_4 + 3HBr}$$
Q3. How does $\mathrm{BrO_2}$ react with alkalis?
Ans: It decomposes at 0°C and is a yellow solid only at –40°C. It forms a combination of bromides and bromates when it dissolves in alkalies.
$$\mathrm{6BrO_2 + 6NaOH\:\rightarrow\:5NaBrO_3 + NaBr + 3H_2O}$$
Q4. Write a brief explanation of Bromine Monoxide $\mathrm{(Br_2O)}$.
(i) Bromine vapours react with dry HgO at 323–343 K to produce Bromine Monoxide.
$$\mathrm{2HgO + 2Br_2\:\rightarrow\:HgO.HgBr_2 + Br_2O}$$
(ii) It's a dark brown liquid with a freezing point of –17.5 degrees Celsius. In the presence of alkalis, it disproportionates.
$$\mathrm{6NaOH + 6Br_2O\:\rightarrow\:5NaBrO_3 + NaBr + 3H_2O}$$
(iii) It is an anhydride of $\mathrm{HBrO}$ because it dissolves in water to form hypobromous acid. It operates as an oxidizing agent, oxidizing iodine to iodine pentoxide.
$$\mathrm{5Br_2O + I_2\:\rightarrow\:I_2O_5 + 5Br_2}$$
Q5. Why does dilute solution of $\mathrm{HBr}$ tends to turn yellow?
Ans: Because of the oxidation of $\mathrm{HBr}$ to $\mathrm{Br_2}$, dilute aqueous solutions of $\mathrm{HBr}$ are particularly reactive and become yellow.
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