Intermolecular forces (IMFs) are the attractive forces that exist between molecules, acting like a faint glue that holds them together. These forces are entirely distinct from the much stronger covalent bonds that hold atoms together within a single molecule. The collective strength of IMFs directly determines the physical state of a substance (solid, liquid, or gas). Identifying the specific forces in a substance like ethane (\(\text{C}_2\text{H}_6\)) is essential for understanding its characteristic physical properties.
Understanding Intermolecular Forces
IMFs are categorized into three main types that vary significantly in strength. The strongest is hydrogen bonding, a special interaction that only occurs when a hydrogen atom is directly bonded to a highly electronegative atom like nitrogen, oxygen, or fluorine.
Dipole-dipole interaction occurs between molecules that possess a permanent electrical polarity. This polarity arises from an unequal sharing of electrons, creating slight positive and negative ends. These opposite partial charges on neighboring molecules then attract one another.
The third force, and the weakest, is the London Dispersion Force (LDF). LDF is present in all substances regardless of their polarity. LDFs result from temporary, fleeting dipoles that form as electrons constantly move around the molecule. The presence and molecular structure of a substance determines which of these forces will be the dominant factor.
Analyzing Ethane’s Molecular Structure
Ethane (\(\text{C}_2\text{H}_6\)) is a simple hydrocarbon composed of two carbon atoms and six hydrogen atoms. It has a single bond between the carbons, and the geometry around each carbon is tetrahedral. This structure gives the molecule an overall symmetrical shape.
A slight difference in electronegativity creates small polarity in each \(\text{C-H}\) bond. However, the symmetrical arrangement is key. The individual bond polarities point away from the center in opposing directions and effectively cancel each other out. This results in a net molecular dipole moment of zero, meaning ethane is nonpolar.
Because ethane is nonpolar, it cannot engage in dipole-dipole interactions. Since ethane contains no hydrogen bonded to nitrogen, oxygen, or fluorine, hydrogen bonding is also not possible.
Identifying the Specific Force in Ethane
The only intermolecular force acting between ethane molecules is the London Dispersion Force (LDF), also known as a van der Waals force. LDFs exist in all molecules, but they are the sole force of attraction for nonpolar molecules like ethane. This force arises from the continuous, random motion of electrons within the molecule.
At any given instant, the electrons may be unevenly distributed, creating a temporary, instantaneous dipole. This momentary imbalance influences the electron cloud of an adjacent ethane molecule, inducing a corresponding temporary dipole. The resulting attraction between these two fleeting opposite charges is the LDF.
LDFs increase in strength with molecular size and mass. Larger molecules have more electrons, making the electron cloud more easily distortable. This ease of distortion is called polarizability. For a small molecule like ethane, the LDFs are relatively weak, but they are the only mechanism of attraction.
How London Dispersion Forces Affect Ethane
The presence of only weak London Dispersion Forces impacts ethane’s physical state. To change a substance from a liquid to a gas, energy must be supplied to overcome the forces holding the molecules together. Because the LDFs in ethane are weak, less energy is required to separate the molecules.
This low energy requirement is reflected in ethane’s low boiling point, approximately \(-88.5^\circ\text{C}\). At standard room temperature and pressure, ethane exists as a colorless, odorless gas. The equally low melting point of about \(-182.8^\circ\text{C}\) demonstrates the ease with which ethane transitions from solid to liquid. Ethane’s behavior contrasts sharply with substances like water, which require much more energy for phase changes due to stronger hydrogen bonds.