Boiling point refers to the specific temperature at which a liquid changes into a gas. This transformation occurs when a substance’s vapor pressure equals the surrounding atmospheric pressure, allowing bubbles of gas to form throughout the liquid.
Fundamental Factors Influencing Boiling Point
The temperature required for a liquid to boil is primarily determined by the strength of the forces holding its molecules together. These intermolecular forces (IMFs) must be overcome for molecules to escape the liquid phase and enter the gas phase. Stronger intermolecular forces require more energy, resulting in a higher boiling point for the substance.
One of the strongest types of intermolecular forces is hydrogen bonding, which occurs when a hydrogen atom bonded to a highly electronegative atom like nitrogen, oxygen, or fluorine is attracted to another electronegative atom in a different molecule. Molecules capable of forming hydrogen bonds, such as water, often exhibit significantly higher boiling points compared to molecules of similar size that lack this ability.
Another important class of intermolecular forces is dipole-dipole interactions, which arise between polar molecules that possess permanent positive and negative ends. These attractive forces are generally stronger than London Dispersion Forces but weaker than hydrogen bonds. Molecules with significant dipole moments will have higher boiling points than nonpolar molecules of comparable molecular weight.
London Dispersion Forces (LDFs), also known as van der Waals forces, are present in all molecules, whether polar or nonpolar. These temporary, instantaneous dipoles occur due to the constant movement of electrons, inducing temporary dipoles in neighboring molecules. The strength of LDFs increases with molecular size and surface area, as larger molecules have more electrons and a greater surface area for these transient interactions.
Molecular shape also plays a significant role in determining the boiling point, particularly for molecules with similar molecular weights. Straight-chain molecules have a larger surface area for intermolecular contact, allowing for more extensive London Dispersion Forces. In contrast, branched or compact molecules have a smaller surface area available for these interactions, which reduces the overall strength of the LDFs.
Applying the Rules: Comparing Different Compound Types
Determining which compound has the highest boiling point involves systematically evaluating the intermolecular forces present and considering molecular size and shape. If hydrogen bonding is present in one compound but absent in others of similar molecular weight, the hydrogen-bonded compound will have the highest boiling point.
Consider, for example, comparing water (H₂O) and methane (CH₄). Water molecules can form extensive hydrogen bonds due to the presence of hydrogen atoms bonded to oxygen. Methane, on the other hand, is a nonpolar molecule that only exhibits London Dispersion Forces. Despite methane having a slightly higher molecular weight, water’s ability to form strong hydrogen bonds gives it a boiling point of 100°C, while methane boils at -161.5°C.
When comparing compounds that only exhibit London Dispersion Forces, molecular weight becomes the primary factor. For instance, pentane (C₅H₁₂) and octane (C₈H₁₈) are both nonpolar hydrocarbons. Octane has a significantly higher molecular weight than pentane, meaning it has more electrons and a larger surface area for London Dispersion Forces. As a result, octane boils at 125.7°C, which is considerably higher than pentane’s boiling point of 36.1°C.
For compounds with the same molecular formula but different structures, known as isomers, molecular shape becomes important. Normal pentane (n-pentane) is a straight-chain molecule, while neopentane (2,2-dimethylpropane) is a highly branched isomer. Although they have the same molecular weight, n-pentane’s linear shape allows for greater surface contact and stronger London Dispersion Forces, giving it a higher boiling point of 36.1°C compared to neopentane’s boiling point of 9.5°C.
Influence of External Conditions
Beyond a substance’s intrinsic molecular properties, external conditions can also modify its observed boiling point. Atmospheric pressure is a significant external factor influencing when a liquid boils. Therefore, at higher atmospheric pressures, more energy is required to reach this equilibrium, leading to a higher boiling point.
Conversely, at lower atmospheric pressures, such as at high altitudes, the external pressure is reduced. This means less energy is needed for the liquid’s vapor pressure to match the surroundings, resulting in a lower boiling point. For example, water boils at approximately 100°C at sea level but at a lower temperature, around 90°C, in Denver, Colorado, due to the city’s higher altitude and consequently lower atmospheric pressure.
The presence of dissolved impurities in a liquid can also affect its boiling point. Dissolving a non-volatile substance in a liquid typically elevates its boiling point, a phenomenon known as boiling point elevation.