Physical Properties of Haloalkanes and Haloarenes

Physical State at Room Temperature

The state in which a haloalkane exists at room temperature is decided by the strength of its intermolecular forces, which in turn rise steadily with molecular size. According to NCERT §6.6, methyl chloride, methyl bromide, ethyl chloride and some chlorofluoromethanes are gases at room temperature. As the carbon chain lengthens, the higher members become liquids and eventually solids.

NIOS §25.3.2 sharpens this picture: the lower alkyl halides $\ce{CH3F}$, $\ce{CH3Cl}$, $\ce{CH3Br}$ and $\ce{C2H5Cl}$ are gases, while alkyl halides up to about $\ce{C18}$ are liquids with high boiling points. Among the aromatic compounds, all monohalobenzenes such as chlorobenzene are colourless liquids at room temperature. The progression gas → liquid → solid simply tracks the growing van der Waals attraction between larger, heavier molecules.

Why Halocompounds Are Polar

Every physical property in this topic rests on one structural fact: the carbon–halogen bond is polar. Halogens are more electronegative than carbon, so the bonding electrons are displaced toward the halogen. The carbon therefore carries a partial positive charge and the halogen a partial negative charge, written as $\ce{C^{\delta+}-X^{\delta-}}$. This permanent dipole means neighbouring molecules attract one another through dipole–dipole forces, on top of the ever-present van der Waals (London dispersion) forces.

Both NCERT and NIOS describe halocompounds as polar yet only moderately so — polar enough to raise boiling points above the parent hydrocarbon, but not polar enough to dissolve in water. Keeping this dual character in mind resolves nearly every apparent contradiction in the trends below. The deeper origin of the dipole is treated in the sibling note on the nature of the C–X bond.

Boiling Point Trends

The boiling point measures how much thermal energy is needed to overcome intermolecular attraction. NCERT §6.6 establishes the master comparison first: the boiling points of chlorides, bromides and iodides are considerably higher than those of the hydrocarbons of comparable molecular mass. The reason is twofold — greater polarity (the permanent dipole adds dipole–dipole attraction) and higher molecular mass (more electrons, larger van der Waals forces). NIOS §25.3.2 gives the identical two-point explanation.

Within the halogen family, for the same alkyl group the boiling point falls in the order:

$$\ce{RI > RBr > RCl > RF}$$

As we move from fluorine to iodine, the halogen atom grows larger and heavier and carries more electrons. NCERT attributes the rising boiling point to the increasing magnitude of van der Waals forces with the size and mass of the halogen atom — equivalently, larger atoms are more polarisable, so the temporary dipoles that drive dispersion forces are stronger. The figure below traces this rise for the methyl and ethyl halides using the data of NIOS Table 25.2.

A second comparison runs along the carbon chain. For the same halogen, the boiling point rises as the alkyl group lengthens, because each added $\ce{-CH2-}$ unit increases molecular mass and surface area, strengthening van der Waals contact. This is why a longer chloroalkane boils higher than a shorter one of the same halogen.

Branching and Isomeric Haloalkanes

When two haloalkanes share the same molecular formula, mass can no longer separate them — shape decides the boiling point. NCERT §6.6 states that the boiling points of isomeric haloalkanes decrease with an increase in branching. A branched molecule is more compact and spherical, with less surface available for van der Waals contact, so its intermolecular forces are weaker and it boils lower.

The textbook example is the set of isomers of $\ce{C4H9Br}$. Among them, the most branched isomer, 2-bromo-2-methylpropane $\ce{(CH3)3CBr}$, has the lowest boiling point — its near-spherical shape minimises surface contact. The straight-chain 1-bromobutane, with the largest surface area, boils highest.

Melting Points and the Para Anomaly

Melting point measures how easily a crystal lattice is broken, so it depends not only on intermolecular force strength but also on how well the molecules pack into the solid. This packing factor produces the single most-tested anomaly of the chapter. NCERT §6.6 records that the boiling points of the three isomeric dihalobenzenes are very nearly the same, but the para isomers are high-melting compared with the ortho and meta isomers.

The reason is symmetry. The para isomer is the most symmetrical of the three; its compact, regular shape fits into the crystal lattice far better than the lower-symmetry ortho and meta forms. A tightly packed lattice has stronger overall attraction and a higher lattice energy, so more heat is required to melt it. NIOS §25.3.2 gives the explicit melting points for the dichlorobenzenes, which the figure below visualises.

Density Trends

Density behaviour separates the lighter chloroalkanes from the heavier bromo and iodo compounds. NCERT §6.6 states the rule plainly: bromo, iodo and polychloro derivatives of hydrocarbons are heavier than water. Density increases with an increase in (i) the number of carbon atoms, (ii) the number of halogen atoms, and (iii) the atomic mass of the halogen atom. The dominant factor is the heavy halogen packed into a relatively small molecular volume.

Two patterns from NCERT Table 6.3 are worth memorising. First, holding the alkyl group fixed at $\ce{n-C3H7}$, density climbs $\ce{Cl(0.89) < Br(1.335) < I(1.747)}$ — confirming that a heavier halogen raises density. Second, increasing the number of chlorine atoms on a single carbon raises density from $\ce{CH2Cl2 (1.336)}$ to $\ce{CHCl3 (1.489)}$ to $\ce{CCl4 (1.595)}$. Note that n-propyl chloride at 0.89 g/mL is the one entry lighter than water; a single chlorine on a small chain is not enough to exceed water's density.

Solubility in Water and Organic Solvents

Although haloalkanes are polar, they are very slightly soluble in water — a result that often surprises students who equate polarity with water solubility. The explanation is energetic. NCERT §6.6 reasons as follows: to dissolve a haloalkane in water, energy must be supplied to overcome the attractions between the haloalkane molecules and to break the hydrogen bonds between water molecules. When new attractions form between haloalkane and water molecules, less energy is released because these new forces are not as strong as water's original hydrogen bonds. The energy account does not balance, so dissolution is unfavourable and solubility stays low.

NIOS §25.3.2 states the same conclusion more directly: halocompounds are immiscible in water because of their inability to form hydrogen bonds with water molecules. A haloalkane has no $\ce{-OH}$ or $\ce{-NH}$ to donate a hydrogen bond, and its $\ce{C-X}$ dipole is too weak to compete with water's hydrogen-bond network.

The opposite holds for organic solvents. Haloalkanes dissolve readily in solvents such as benzene, ether and chloroform because the new intermolecular attractions formed between haloalkane and solvent molecules are of much the same strength as the attractions being broken in the separate liquids. The energy balance is roughly neutral, so mixing proceeds — the chemical paraphrase of "like dissolves like."