The boiling point, the temperature at which a liquid turns into a gas, measures the energy needed to separate its molecules. This property reveals the attractive forces holding a substance together. Comparing the halogens chlorine (\(\text{Cl}_2\)) and bromine (\(\text{Br}_2\)), a clear difference emerges: chlorine boils at \(-34.0^\circ\text{C}\), while bromine remains liquid until \(58.8^\circ\text{C}\). Understanding this disparity requires examining the microscopic forces acting between these diatomic molecules.
The Forces Governing Boiling Point
The boiling process involves supplying thermal energy to overcome intermolecular forces (IMFs). Strong IMFs require high energy, resulting in a high boiling point. Conversely, weak IMFs translate to a low boiling point.
Since both \(\text{Cl}_2\) and \(\text{Br}_2\) are composed of two identical atoms, they are nonpolar and electrically symmetrical. They do not possess permanent positive or negative ends, ruling out stronger dipole-dipole attractions or hydrogen bonds. Consequently, the only relevant intermolecular force acting between these molecules is the London Dispersion Force (LDF).
LDFs arise from the constant, random movement of electrons within a molecule. At any instant, electrons can become unevenly distributed, creating a momentary, or instantaneous, dipole. This temporary charge imbalance induces a corresponding dipole in a neighboring molecule, resulting in a fleeting attractive force. LDFs are universally present, but they are the sole determinant of the boiling point for nonpolar substances like chlorine and bromine.
Molecular Size, Mass, and Polarizability
The strength of the London Dispersion Force increases significantly with the size and mass of the molecule. This increase is directly related to polarizability, which describes the ease with which a molecule’s electron cloud can be temporarily distorted by a nearby instantaneous dipole.
As a molecule grows larger, it contains a greater number of electrons, and these electrons are located farther from the positively charged nucleus. This increased distance means the outer electrons are held more loosely. A large, diffuse electron cloud is therefore much easier to distort or “polarize” than a small, tightly held cloud.
When the electron cloud is highly polarizable, the temporary dipoles that form are larger and last longer. This results in stronger and more persistent attractive forces between neighboring molecules. The stronger the induced-dipole attraction, the greater the energy required to break them apart, leading to a higher boiling point. This trend is consistently observed moving down the halogen group, where atomic size and mass increase.
The Direct Comparison: Chlorine Versus Bromine
The fundamental difference in boiling points between chlorine and bromine is a direct consequence of this size-dependent polarizability. Chlorine (\(\text{Cl}_2\)) has a total of 34 electrons. Bromine (\(\text{Br}_2\)), located one row below chlorine, is significantly larger and heavier, possessing a total of 70 electrons.
The \(\text{Br}_2\) molecule has more than twice the number of electrons as the \(\text{Cl}_2\) molecule, and these electrons are distributed over a much larger volume. This large, loose electron cloud makes bromine highly polarizable, generating substantial London Dispersion Forces between \(\text{Br}_2\) molecules. Due to these stronger forces, greater thermal energy must be applied to overcome the intermolecular attraction, which explains its elevated boiling point of nearly \(59^\circ\text{C}\).
In contrast, the smaller \(\text{Cl}_2\) molecule has a smaller, more compact electron cloud that is less easily distorted, resulting in much weaker London Dispersion Forces. Because the attraction between \(\text{Cl}_2\) molecules is weaker, less energy is needed to separate them and cause the phase change from liquid to gas. This requirement for less energy is why chlorine boils at a much lower temperature, below freezing, compared to bromine.