Question

What type of intermolecular forces is active in the liquid state of each of the following substances? a. Ne b. CO c. \(\mathrm{CH}_{3} \mathrm{OH}\) d. \(\mathrm{Cl}_{2}\)

Short answer

a. Ne: Dispersion (London) forces only b. CO: Dispersion forces and dipole-dipole interactions c. \(\mathrm{CH}_{3} \mathrm{OH}\): Dispersion forces, dipole-dipole interactions, and hydrogen bonding d. \(\mathrm{Cl}_{2}\): Dispersion (London) forces only

Step-by-step solution

Identify the structure and polarity of each molecule

Ne: Neon is a noble gas, so in its liquid state, it exists as Ne atoms without any covalent bonds. CO: Carbon monoxide is a covalent molecule consisting of one carbon atom and one oxygen atom, with a triply bonded structure. It is a polar molecule because oxygen is more electronegative than carbon, creating a dipole. \(\mathrm{CH}_{3} \mathrm{OH}\): Methanol (CH3OH) is a polar covalent molecule, with the oxygen atom bonded to one of the hydrogen atoms in a hydroxyl group (-OH). \(\mathrm{Cl}_{2}\): Chlorine molecules are diatomic and covalent, consisting of two chlorine atoms bonded together. Since both atoms are chlorine, the molecule is nonpolar.

Identify the type of intermolecular forces present in each molecule

Ne: As Ne has no covalent bonds or polarity, the intermolecular forces present are only dispersion (London) forces, which are weak forces caused by temporary dipoles. CO: In carbon monoxide, the dipole-dipole interactions are present due to the molecule's polarity. Dispersion forces are also present, as they exist between all molecules. \(\mathrm{CH}_{3} \mathrm{OH}\): Methanol's hydroxyl group allows it to form hydrogen bonds with other molecules, as well as dipole-dipole interactions due to its polarity. Dispersion forces are also present, as they are in all molecules. \(\mathrm{Cl}_{2}\): Since the chlorine molecule is nonpolar, only dispersion forces are present in the liquid state.

Final summary of intermolecular forces for each substance

a. Ne: Dispersion (London) forces only b. CO: Dispersion forces and dipole-dipole interactions c. \(\mathrm{CH}_{3} \mathrm{OH}\): Dispersion forces, dipole-dipole interactions, and hydrogen bonding d. \(\mathrm{Cl}_{2}\): Dispersion (London) forces only

Key concepts

London Dispersion Forces

London Dispersion Forces are the weakest type of intermolecular attraction, but they play a crucial role, especially in nonpolar molecules. These forces arise due to temporary fluctuations in electron density within molecules, creating small instantaneous dipoles. This happens when, at any given moment, the electrons within an atom or molecule are unevenly distributed. These temporary dipoles induce neighboring molecules to also become temporarily polarized, leading to an attractive force between these entities.

All atoms and molecules exhibit these forces to some extent; however, they are particularly significant in nonpolar substances, such as noble gases like Neon (Ne) and nonpolar molecules like Chlorine ( Cl_2 ). As an example, even though Ne and Cl_2 have no permanent dipole to interact with one another, their electron clouds can shift and overlap, creating these transient attractive forces that allow them to condense into liquids under appropriate conditions.

Dipole-Dipole Interactions

Dipole-Dipole interactions occur between polar molecules, which have permanent dipoles due to differences in electronegativity between bonded atoms. This type of force is generally stronger than London Dispersion Forces because it involves interactions between permanent dipoles rather than transient ones.

In molecules like Carbon Monoxide (CO), the difference in electronegativity between carbon and oxygen leads to a separation of charge, creating a dipole. When such polar molecules approach each other, the positive end of one molecule is attracted to the negative end of another, forming a dipole-dipole attraction.

• These interactions play a significant role in determining the physical properties of substances, such as boiling and melting points.
• Stronger dipole interactions usually result in higher boiling points compared to those involving only London Dispersion Forces.

Hydrogen Bonding

Hydrogen Bonding is a special case of dipole-dipole interaction, but much stronger, often impacting the substance's structure and properties dramatically. It occurs when hydrogen is covalently bonded to very electronegative atoms like oxygen, nitrogen, or fluorine. In such molecules, the hydrogen carries a partial positive charge, making it strongly attracted to lone pairs of electrons on neighboring electronegative atoms.

Methanol ( CH_3 OH) is a prime example where hydrogen bonding is prevalent. The oxygen in the hydroxyl group is highly electronegative, drawing electron density away from hydrogen. This interaction leads to hydrogen bonds forming with the oxygen atoms of adjacent methanol molecules.

• Hydrogen bonding significantly elevates methanol's boiling point.
• This interaction is crucial for understanding the behavior of water and biological molecules like DNA and proteins.
• It adds to the cohesive nature of methanol, owing to the strength of these interactions compared to regular dipole-dipole attractions.