Hydrocarbons are organic compounds composed entirely of hydrogen and carbon atoms. The boiling point is the temperature at which a substance transitions from a liquid to a gas. Whether hydrocarbons have low boiling points depends almost entirely on the size of the molecule. Small hydrocarbons, such as methane, are gases with extremely low boiling points, while very large hydrocarbons, like those found in asphalt, are solids at room temperature with very high boiling points. The broad range of boiling points is governed by the forces of attraction between individual molecules.
The General Role of Intermolecular Forces in Boiling
Boiling requires energy to overcome the attractive forces holding molecules together in the liquid state. This involves overcoming the weaker attractions that exist between neighboring molecules, known as intermolecular forces, not the strong covalent bonds within the molecule. The energy required to break these forces dictates the boiling point.
There are three primary types of intermolecular forces. The weakest are London Dispersion Forces (LDFs), which are the sole attractive force in nonpolar molecules. Stronger forces include hydrogen bonding and dipole-dipole interactions. A substance relying only on LDFs will generally boil at a much lower temperature than one with strong hydrogen bonds, provided the molecules are small.
Why Hydrocarbons Rely Solely on London Dispersion Forces
The physical properties of hydrocarbons are dictated by the nonpolar nature of the carbon-hydrogen (C-H) bond. The electronegativity difference between carbon and hydrogen is very small, meaning electron sharing is nearly equal. This classifies the C-H bond as nonpolar, making the overall hydrocarbon molecule nonpolar.
Since hydrocarbons lack permanent positive or negative ends, they cannot engage in stronger attractions like hydrogen bonding or dipole-dipole interactions. Consequently, the only attractive force operating between them is the London Dispersion Force (LDF). LDFs arise from the continuous movement of electrons, which can momentarily distribute unevenly, creating a temporary dipole. This transient attraction influences neighboring molecules, governing the boiling points of all hydrocarbons.
How Molecular Size and Structure Determine Boiling Point Ranges
The strength of London Dispersion Forces is not constant; it scales dramatically with both the size and the shape of the molecule. The larger a hydrocarbon molecule is, the more electrons it contains, and the greater its surface area becomes. More electrons mean the electron cloud is more easily distorted, a property called polarizability, which allows for stronger temporary dipoles to form. A larger surface area also increases the number of points of contact between adjacent molecules, which allows the weak LDFs to add up to a substantial overall force of attraction.
This direct relationship between size and boiling point explains the vast range of physical states seen in hydrocarbons. Methane, the smallest hydrocarbon with just one carbon atom, has a boiling point of about -164 degrees Celsius and is a gas at room temperature because its LDFs are minimal. Hexane, a straight-chain hydrocarbon with six carbon atoms, is a liquid that boils at 69 degrees Celsius, demonstrating the increase in LDF strength with chain length. Conversely, paraffin wax, which consists of hydrocarbon chains with twenty or more carbon atoms, has a boiling point above 370 degrees Celsius and is a solid at room temperature.
Molecular structure also plays a significant role through branching. Hydrocarbon isomers, which are molecules with the same chemical formula but different arrangements of atoms, exhibit different boiling points. A straight-chain molecule, such as n-pentane, has a larger surface area and can pack closely with its neighbors, maximizing the strength of the LDFs.
In contrast, a branched isomer, like 2,2-dimethylpropane, is more compact and spherical. This spherical shape reduces the overall surface area available for contact between molecules, which in turn weakens the total London Dispersion Forces. For example, n-pentane boils at 36 degrees Celsius, while its branched isomer, 2,2-dimethylpropane, boils at a much lower temperature of 9.5 degrees Celsius. Therefore, while small hydrocarbons have low boiling points due to their reliance on weak LDFs, the strength of these same forces increases with molecular size, causing large hydrocarbons to have very high boiling points.