Chlorine, in its elemental form, exists as the diatomic molecule Cl2. Understanding the behavior of this simple molecule requires a look at the intermolecular forces present between Cl2 units in the condensed phases. These forces, while significantly weaker than the covalent bonds holding the chlorine atoms together within a molecule, are the primary determinants of chlorine's physical properties, such as its boiling point, melting point, and solubility in various solvents.
Nature of the Cl2 Molecule
To analyze the intermolecular forces, one must first consider the molecule itself. A Cl2 molecule consists of two chlorine atoms sharing a pair of electrons in a covalent bond. This bond is nonpolar because the electronegativity difference between the two identical atoms is zero. Consequently, the molecule has no permanent dipole moment, classifying it as a nonpolar entity. The electron distribution around the molecule is symmetric, but this distribution is not static; the electrons are in constant motion, leading to temporary fluctuations in electron density.
London Dispersion Forces
The dominant intermolecular force for Cl2 is the London dispersion force, also known as induced dipole-induced dipole interactions. These forces arise due to the instantaneous, random movement of electrons within the molecule. At any given moment, the electrons might be asymmetrically distributed, creating a fleeting dipole. This temporary dipole can induce a complementary dipole in a neighboring Cl2 molecule, resulting in a weak, short-lived attraction. While individual London forces are very weak, their cumulative effect across many molecules is significant, especially in larger atoms or molecules with more electrons. Chlorine, being a relatively large atom with 34 electrons, exhibits substantial London dispersion forces compared to smaller noble gases like helium or neon.
Impact on Physical Properties
The strength of these London dispersion forces directly dictates the macroscopic behavior of chlorine. At standard temperature and pressure, chlorine is a gas. This state exists because the kinetic energy of the Cl2 molecules is sufficient to overcome the relatively weak intermolecular attractions. However, cooling the gas reduces the kinetic energy, allowing the dispersion forces to pull the molecules together, condensing chlorine into a liquid at -34.04°C and a solid at -101.5°C. The relatively low boiling point of chlorine is a direct consequence of the weakness of these intermolecular forces compared to, for example, the hydrogen bonding seen in water.
Comparison with Polar Molecules
It is instructive to compare chlorine's intermolecular forces with those of polar molecules, such as hydrogen chloride (HCl). While HCl also exhibits London dispersion forces, it possesses a strong permanent dipole due to the significant electronegativity difference between hydrogen and chlorine. This results in dipole-dipole interactions, which are stronger than the dispersion forces in Cl2. Consequently, HCl has a higher boiling point (-85°C) than chlorine (-34°C), despite chlorine being a larger molecule. This illustrates that for nonpolar molecules like Cl2, the London forces are the sole determinant of their phase behavior under standard conditions.
Role in Solubility
The principle of "like dissolves like" governs the solubility of Cl2. Because Cl2 is a nonpolar molecule, it is poorly soluble in polar solvents like water. The strong hydrogen bonding network in water creates a highly organized structure that nonpolar molecules cannot easily disrupt. However, Cl2 is moderately soluble in nonpolar solvents, such as hexane or carbon tetrachloride. In these environments, the weak London dispersion forces between Cl2 molecules and the solvent molecules are energetically favorable, allowing the chlorine to dissolve. This property is crucial in applications like the extraction of chlorine compounds or in certain industrial processes.
