The boiling point of water is \(100^{\circ}\text{C}\) at standard pressure, while chlorine boils at a frigid \(-34^{\circ}\text{C}\). This vast gap in thermal behavior is dictated entirely by the microscopic forces that hold the molecules together. The fundamental reason for the disparity lies in the distinct types and relative strengths of the attractive forces operating between water and chlorine molecules.
What Determines a Boiling Point
The boiling point of any substance is the temperature at which its liquid form changes into a gas. This phase transition requires energy, typically heat, to overcome the forces of attraction between neighboring molecules in the liquid state.
As the temperature rises, molecules gain kinetic energy and move faster. Boiling occurs when this motion is energetic enough to separate the molecules from the attractive pull of their neighbors. The stronger these forces are, the more heat energy is required, resulting in a higher boiling temperature.
The energy used during boiling does not break the chemical bonds within the molecules, such as the covalent bonds holding atoms together. Instead, the energy targets the weaker attractions between one molecule and another. These external forces are the primary factor that determines the boiling point of a substance.
The Forces Between Molecules
The attractive forces that exist between separate molecules are collectively known as Intermolecular Forces (IMFs). The strength of these IMFs determines how easily a liquid will vaporize, directly influencing its boiling point.
There are three main categories of intermolecular forces. The weakest are the London Dispersion Forces (LDFs), present in all molecules but the only force in non-polar molecules. Next in strength are the Dipole-Dipole interactions, which occur when two polar molecules align so the positive end of one attracts the negative end of the other.
The strongest of the three principal IMFs is Hydrogen Bonding, a specialized type of dipole-dipole interaction. These forces exist only when a hydrogen atom is covalently bonded to a highly electronegative atom: nitrogen, oxygen, or fluorine. Substances primarily governed by hydrogen bonding will have much higher boiling points than those relying on weaker dipole-dipole or dispersion forces.
Water’s Strong Hydrogen Bonding
Water (H2O) has an exceptionally high boiling point because its molecular structure allows for the formation of powerful hydrogen bonds. Oxygen is far more electronegative than hydrogen, pulling shared electrons closer to itself. This unequal sharing creates a permanent partial negative charge on the oxygen atom and partial positive charges on the two hydrogen atoms, making the molecule highly polar.
This strong polarity allows the positively charged hydrogen atom of one water molecule to form a powerful attraction with the negatively charged oxygen atom of a neighboring molecule, creating a hydrogen bond. In liquid water, each molecule forms, on average, about 3.5 hydrogen bonds. This results in an extensive three-dimensional network of attractions that must be overcome for the water to boil.
The energy required to break a single hydrogen bond is approximately 21 kJ/mole, which is substantial for an intermolecular force. Because so many hydrogen bonds are formed per molecule, the total energy needed to separate the molecules and transition them into the gas phase is very high. This massive collective force anchors water molecules together, forcing the boiling point up to 100°C.
Chlorine’s Weak Dispersion Forces
The elemental chlorine molecule (Cl2) presents a stark contrast to water, having a boiling point of -34°C. Chlorine is composed of two identical atoms sharing electrons equally, resulting in a perfectly non-polar molecule. Since there are no permanent positive or negative poles, the stronger dipole-dipole and hydrogen bonding forces cannot form between chlorine molecules.
The only intermolecular forces operating between chlorine molecules are the weak London Dispersion Forces (LDFs). LDFs arise from the constant motion of electrons within a molecule. Electrons may temporarily cluster on one side, creating a fleeting, induced dipole moment.
This momentary dipole influences a neighbor, inducing a corresponding opposite dipole, leading to a very weak, short-lived attraction. Because these attractions are temporary and easily disrupted, they are the weakest of all intermolecular forces. The energy needed to overcome these forces is minimal, estimated to be in the range of 0 to 10 kJ.
The minimal energy required to separate the non-polar chlorine molecules means that very little thermal energy is sufficient to cause the liquid-to-gas phase change. This explains why chlorine is a gas at room temperature and only condenses into a liquid at temperatures far below zero.
The Energy Gap: Why Water Wins
The dramatic difference between the boiling points of water (100°C) and chlorine (-34°C) is a direct consequence of the energy differential between their respective intermolecular forces. Water molecules are locked into a robust, extensive network of strong hydrogen bonds, which requires a significant amount of heat energy to break.
Chlorine molecules, conversely, are held together only by the feeble, transient London Dispersion Forces. These weak attractions are overcome by minimal thermal energy, allowing the molecules to escape into the gas phase at a very low temperature. The energy gap between breaking the vast network of hydrogen bonds in water and the sparse dispersion forces in chlorine is the fundamental reason for water’s superior boiling point.