18.11 Occurrence, Preparation, and Properties of Halogens

Occurrence and Preparation of Halogens

The halogens (Group 17) are among the most reactive nonmetals on the periodic table. They never occur as free elements in nature because they're so reactive, so extracting them requires specific chemical or electrochemical methods. Understanding where they're found and how they're isolated connects directly to their reactivity trends.

Natural Sources

Each halogen has characteristic natural sources:

• Fluorine is found primarily in the minerals fluorite ($CaF_2$) and cryolite ($Na_3AlF_6$). These are mined from geological deposits.
• Chlorine occurs most abundantly as sodium chloride ($NaCl$) in seawater and underground salt deposits. Seawater contains roughly 1.9% chloride ions by mass.
• Bromine is present as bromide salts (like $NaBr$) in seawater, salt lakes, and underground brines. It's much less abundant than chlorine.
• Iodine exists in trace amounts in seawater and is concentrated by certain seaweeds. It's also found alongside Chile saltpeter ($NaNO_3$) deposits as sodium iodate ($NaIO_3$).

Preparation Methods

Commercial preparation depends on the halogen:

• Chlorine is most commonly produced by electrolysis of brine (aqueous $NaCl$). This process also yields sodium hydroxide ($NaOH$) and hydrogen gas, making it economically efficient. Electrolysis of molten $NaCl$ produces liquid sodium metal and chlorine gas separately.
• Fluorine is prepared by electrolysis of a molten mixture of $KF$ and $HF$. Because fluorine is the most electronegative element, no chemical oxidizing agent is strong enough to oxidize fluoride ions, so electrolysis is the only option.
• Bromine is prepared by treating bromide-containing brines with chlorine gas. Chlorine is more reactive, so it displaces bromine: $Cl_2 + 2Br^- \rightarrow 2Cl^- + Br_2$
• Iodine is similarly extracted by displacement, using chlorine or bromine to oxidize iodide salts: $Cl_2 + 2I^- \rightarrow 2Cl^- + I_2$

The displacement reactions for bromine and iodine work because a more reactive halogen will always oxidize the halide ion of a less reactive one. This pattern follows directly from the reactivity trend: $F_2 > Cl_2 > Br_2 > I_2$.

Properties and Applications of Halogens

Physical Properties

All halogens exist as diatomic molecules ($F_2$, $Cl_2$, $Br_2$, $I_2$) in their elemental form. Their physical properties change predictably down the group because of increasing atomic mass and stronger London dispersion forces (a type of van der Waals force):

Melting and boiling points increase steadily down the group for the same reason: larger electron clouds create stronger London dispersion forces between molecules.

Chemical Properties

Several key chemical trends run through the halogen group:

Reactivity and oxidizing power decrease down the group. Fluorine is the strongest oxidizing agent of all elements. This trend exists because smaller halogens have a greater ability to attract and accept an electron into their valence shell.

Electronegativity decreases down the group: F (4.0), Cl (3.0), Br (2.8), I (2.5). Fluorine is the most electronegative element on the entire periodic table.

Reactions with water differ by halogen:

• $F_2$ reacts violently with water, actually oxidizing the oxygen in water to produce $O_2$ (or $OF_2$) and $HF$.
• $Cl_2$ reacts with water to form hydrochloric acid and hypochlorous acid: $Cl_2 + H_2O \rightarrow HCl + HOCl$
• $Br_2$ and $I_2$ react only slightly with water.

Reactions with metals produce ionic metal halides (like $NaCl$, $MgBr_2$). Reactivity with metals decreases down the group, matching the overall reactivity trend.

Hydrogen halides ($HF$, $HCl$, $HBr$, $HI$) form by direct combination of the elements. An important distinction: acid strength increases down the group ($HI > HBr > HCl > HF$) because the $H–X$ bond gets weaker and longer as the halogen gets bigger, making it easier to donate a proton. $HF$ is actually a weak acid despite fluorine being the most reactive halogen.

Interhalogen Compounds

When two different halogens combine, they form interhalogen compounds such as $ClF_3$, $BrF_5$, or $ICl$. The less electronegative halogen is always the central atom. These compounds tend to be highly reactive and are strong oxidizing or fluorinating agents.

Applications

Halogens and their compounds show up across many industries:

• Water treatment: Chlorine and its compounds (like hypochlorite) are the primary disinfectants used in municipal water purification.
• Polymers: $PVC$ (polyvinyl chloride) contains chlorine; $PTFE$ (Teflon) contains fluorine.
• Pharmaceuticals: Many drugs incorporate fluorine or chlorine atoms to improve biological activity (e.g., fluoxetine/Prozac).
• Agriculture: Chlorine and bromine compounds serve as pesticides and herbicides, though methyl bromide use has been restricted due to ozone concerns.
• Flame retardants: Bromine compounds reduce the flammability of textiles and electronics.
• Photography: Silver bromide ($AgBr$) and silver iodide ($AgI$) were the light-sensitive compounds in traditional photographic film.
• Refrigerants: Chlorofluorocarbons (CFCs) were once standard refrigerants but depleted the ozone layer. They've been largely replaced by hydrofluorocarbons (HFCs), which don't contain chlorine.
• Organic synthesis: Halogenation reactions are a fundamental tool for introducing halogen atoms into organic molecules, enabling further chemical transformations.