How Swapping One Hydrogen for a Halogen Transforms a Molecule

Swapping a hydrogen in a hydrocarbon for a halogen creates compounds used as solvents, refrigerants, pesticides and starting materials for medicines. The polar carbon–halogen bond is the key to how they are made and how they react.

This part covers classification and naming, the nature of the C–X bond, preparing haloalkanes, and preparing haloarenes along with their physical properties.

**Halogen compounds are classified by the number of halogen atoms and by the hybridisation of the carbon carrying the halogen — $s{p}^3$ in alkyl, allylic and benzylic halides, $s{p}^2$ in vinylic and aryl halides — and are named in IUPAC as halo-substituted alkanes or arenes.

By the carbon bearing the halogen:

- Alkyl halides** — halogen on an $s{p}^3$ carbon; primary, secondary or tertiary
- Allylic halides — halogen on an $s{p}^3$ carbon next to a C=C, as in CH${}_2$=CHCH${}_2$Cl
- Benzylic halides — halogen on an $s{p}^3$ carbon attached to a benzene ring, as in C${}_6$H${}_5$CH${}_2$Cl
- Vinylic halides — halogen on an $s{p}^2$ carbon of a C=C, as in CH${}_2$=CHCl
- Aryl halides — halogen directly on a benzene ring, as in C${}_6$H${}_5$Cl

Dihalides are geminal when both halogens sit on one carbon, as in CH${}_3$CHCl${}_2$, and vicinal when they sit on adjacent carbons, as in ClCH${}_2$CH${}_2$Cl.

Naming:

- CH${}_3$CH${}_2$CHClCH${}_3$ — sec-butyl chloride; IUPAC 2-chlorobutane
- (CH${}_3$)${}_3$CBr — tert-butyl bromide; 2-bromo-2-methylpropane

Worked example. CH${}_3$CHBrCH${}_2$CH(CH${}_3$)${}_2$ has a five-carbon chain with substituents at C2 and C4 from either end. With a tie, the alphabetically first substituent gets the lower number, so the name is 2-bromo-4-methylpentane.

An everyday example. PVC water pipes are made by polymerising vinyl chloride, a vinylic halide.

The substance. Allylic and benzylic halides are unusually reactive, while vinylic and aryl halides are unusually unreactive — a difference that comes from the carbon's hybridisation.

Halogens are more electronegative than carbon, so the C–X bond is polar, with a partial positive charge on carbon that attracts nucleophiles; from fluorine to iodine the bond grows longer and weaker, so C–I bonds break most easily.

Polarity:

$$\stackrel{\delta +}{C}-\stackrel{\delta -}{X}$$

**Trends in CH${}_3$X from fluorine to iodine:

- Bond length increases** — $139$, $178$, $193$ and $214$ pm for C–F, C–Cl, C–Br and C–I
- Bond enthalpy decreases — about $452$, $351$, $293$ and $234$ kJ mol${}^{-1}$
- Dipole moment — largest for C–Cl, not C–F, because the much shorter C–F bond offsets fluorine's higher electronegativity

Worked example. Comparing the energy to break one C–I bond with one C–Cl bond:

$$\frac{234}{351}\approx 0.67$$

The C–I bond needs only about two-thirds as much energy, which makes alkyl iodides the most reactive haloalkanes in substitution.

An everyday example. Non-stick cookware coatings are built on C–F bonds, among the strongest single bonds carbon forms, which is why they withstand heat and chemicals.

The substance. Reactivity follows bond strength, not polarity — C–F is highly polar yet the least reactive.

Haloalkanes are made by replacing the –OH of an alcohol with a halogen, by free radical halogenation of alkanes, by adding HX or halogens to alkenes, or by exchanging one halogen for another in the Finkelstein and Swarts reactions.

From alcohols:

- **With HCl and anhydrous ZnCl${}_2$**: $\text{ROH}+\text{HCl}\stackrel{{\text{ZnCl}}_2}{\to }\text{RCl}+{\text{H}}_2\text{O}$
- **With PCl${}_5$, PCl${}_3$ or PBr${}_3$
- With thionyl chloride**: $\text{ROH}+{\text{SOCl}}_2\to \text{RCl}+{\text{SO}}_2+\text{HCl}$

From hydrocarbons:

- Free radical halogenation — Cl${}_2$ in UV light gives mixtures of isomers, so it is rarely used for a single product
- HX addition to alkenes — follows Markovnikov's rule, so propene with HBr gives mainly 2-bromopropane
- Halogen addition — Br${}_2$ adds across a C=C to give a vicinal dibromide

Halogen exchange:

- Finkelstein — $\text{RCl}+\text{NaI}\stackrel{\text{dry acetone}}{\to }\text{RI}+\text{NaCl}$; NaCl precipitates, driving the reaction forward
- Swarts — $\text{RBr}+\text{AgF}\to \text{RF}+\text{AgBr}$

Worked example. Propane has $6$ primary and $2$ secondary hydrogens. If every hydrogen reacted equally, monochlorination would give 1-chloropropane and 2-chloropropane in the ratio $6:2=3:1$ — but secondary hydrogens react faster, so the real mixture holds more 2-chloropropane than this predicts.

An everyday example. Bromine water losing its orange colour is the laboratory test for a C=C double bond, based on halogen addition.

The substance. Thionyl chloride gives the purest product, because both by-products escape as gases.

Haloarenes are made by electrophilic substitution of arenes with chlorine or bromine in the presence of a Lewis acid, or from diazonium salts in the Sandmeyer reaction; halogen compounds boil higher than their parent hydrocarbons, are often denser than water, and are only slightly soluble in water.

Electrophilic substitution:

$${\text{C}}_6{\text{H}}_6+{\text{Cl}}_2\stackrel{{\text{FeCl}}_3}{\to }{\text{C}}_6{\text{H}}_5\text{Cl}+\text{HCl}$$

Sandmeyer reaction. Aniline is diazotised with NaNO${}_2$ and HCl at $273$ to $278$ K, and the diazonium salt is treated with a copper(I) halide:

$${\text{C}}_6{\text{H}}_5{\text{N}}_2^{+}{\text{Cl}}^{-}\stackrel{\text{CuCl/HCl}}{\to }{\text{C}}_6{\text{H}}_5\text{Cl}+{\text{N}}_2$$

Physical properties:

- Boiling points rise with halogen mass — RCl < RBr < RI — and fall with branching
- Dihalobenzenes — isomers boil at similar temperatures, but the para isomer melts highest because its symmetry packs well into a crystal
- Density — bromo, iodo and polychloro compounds are denser than water
- Solubility — only slightly soluble in water, which they cannot hydrogen-bond with strongly, but soluble in organic solvents

Worked example. The isomers of C${}_4$H${}_9$Br boil at $375$ K for 1-bromobutane, $364$ K for 2-bromobutane and $346$ K for 2-bromo-2-methylpropane: each added branch makes the molecule more compact, reducing the surface for van der Waals attraction.

An everyday example. Chlorinated solvents in paint strippers dissolve grease and oils that water cannot touch.

The substance. **The para isomer's higher melting point comes from shape, not stronger bonds** — identical molecules simply pack more neatly.