Question

Predict which substance in each of the following pairs would have the greater intermolecular forces. a. \(\mathrm{CO}_{2}\) or \(\mathrm{OCS}\) b. \(\mathrm{SeO}_{2}\) or \(\mathrm{SO}_{2}\) c. \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{NH}_{2}\) or \(\mathrm{H}_{2} \mathrm{NCH}_{2} \mathrm{CH}_{2} \mathrm{NH}_{2}\) d. \(\mathrm{CH}_{3} \mathrm{CH}_{3}\) or \(\mathrm{H}_{2} \mathrm{CO}\) e. \(\mathrm{CH}_{3} \mathrm{OH}\) or \(\mathrm{H}_{2} \mathrm{CO}\)

Short answer

a. OCS b. SeO₂ c. H₂NCH₂CH₂NH₂ d. H₂CO e. CH₃OH

Step-by-step solution

a. CO₂ or OCS

First we have to identify the molecular structure and polarity of each molecule: CO₂: linear, non-polar molecule OCS: linear, polar molecule Since OCS is polar, it will have stronger dipole-dipole forces compared to CO₂, which only has weak London dispersion forces. Therefore, OCS has greater intermolecular forces.

b. SeO₂ or SO₂

Both molecules have bent structures and are polar molecules which have dipole-dipole forces. The difference is the size and polarizability of the central atom: Se: larger and more polarizable than S SeO₂ will have stronger London dispersion forces due to its larger size and polarizability. Since both have dipole-dipole forces, but SeO₂ also has stronger London dispersion forces, SeO₂ has greater intermolecular forces.

c. CH₃CH₂CH₂NH₂ or H₂NCH₂CH₂NH₂

Both of these molecules can form hydrogen bonds because they have N-H bonds. Hydrogen bonding is stronger than dipole-dipole and London dispersion forces. So, we need to analyze the strength of the hydrogen bonds: CH₃CH₂CH₂NH₂: has one NH₂ group for hydrogen bonding H₂NCH₂CH₂NH₂: has two NH₂ groups for hydrogen bonding H₂NCH₂CH₂NH₂ can form more hydrogen bonds, thus it has greater intermolecular forces.

d. CH₃CH₃ or H₂CO

We need to classify the types of intermolecular forces for each molecule: CH₃CH₃: non-polar molecule, London dispersion forces H₂CO: polar molecule, dipole-dipole forces H₂CO has stronger dipole-dipole forces compared to CH₃CH₃, which only has weak London dispersion forces. Therefore, H₂CO has greater intermolecular forces.

e. CH₃OH or H₂CO

We need to analyze the types of intermolecular forces for each molecule: CH₃OH: has an OH group, hydrogen bonding forces H₂CO: polar molecule, dipole-dipole forces CH₃OH has stronger hydrogen bonding forces, which are stronger than the dipole-dipole forces in H₂CO. Consequently, CH₃OH has greater intermolecular forces.

Key concepts

Dipole-Dipole Forces

Dipole-dipole forces occur when polar molecules interact with each other. Picture these as tiny magnets attracting each other. A molecule is polar if it has a difference in electronegativity between its atoms, creating a partial positive and negative charge.
These forces are intermediate in strength compared to other types of intermolecular forces.
For example, OCS (carbonyl sulfide) has dipole-dipole forces because it is a polar molecule. The sulfur atom is more electronegative, creating a dipole moment. Conversely, molecules like CO₂ are non-polar and only have weaker London dispersion forces. Knowing whether a molecule is polar helps predict its ability to exhibit dipole-dipole interactions and, thus, its intermolecular force strength.

London Dispersion Forces

London dispersion forces, sometimes simply called dispersion forces, are the weakest intermolecular forces. They occur due to temporary fluctuations in electron distribution, leading to temporary dipoles, even in non-polar molecules.
All molecules, polar or non-polar, exhibit these forces, but they are the only type present in non-polar molecules.
The strength of London dispersion forces is larger in bigger, more massive atoms and molecules due to their greater polarizability. For example, SeO₂ has more potent dispersion forces than SO₂ because selenium is larger than sulfur, enhancing its polarizability.

Hydrogen Bonding

Hydrogen bonding is a special type of intermolecular force that is significantly stronger than both dipole-dipole forces and London dispersion forces. It occurs when hydrogen is bonded to highly electronegative atoms like nitrogen (N), oxygen (O), or fluorine (F).
This creates a scenario where the hydrogen atom acts as a bridge between two electronegative atoms, leading to a strong dipole interaction.
In the example of H₂NCH₂CH₂NH₂, the presence of more than one NH₂ group allows for more hydrogen bonds, meaning it has stronger intermolecular forces compared to CH₃CH₂CH₂NH₂. This is why hydrogen bonding plays such a critical role in determining physical properties.

Molecular Polarity

Molecular polarity is determined by the shape of the molecule and the electronegativity differences between its atoms. A molecule is polar if it has an uneven distribution of electrons, leading to a dipole moment.
Polarity influences the kind of intermolecular forces a molecule can exhibit.
Non-polar molecules, like methane (CH₄), lack such a dipole and rely solely on London dispersion forces. However, polar molecules like formaldehyde (H₂CO) experience stronger intermolecular attractions due to dipole-dipole forces. Understanding polarity helps predict the types of interactions molecules will have and their strength, crucial for predicting behavior in different states of matter.