Problem 4 - Chemistry and Chemical Reactivity

Question

What type of intermolecular forces must be overcome in converting each of the following from a liquid to a gas? (a) \(\mathrm{CO}_{2}\) (c) \(\mathrm{CHCl}_{3}\) (b) \(\mathrm{NH}_{3}\) (d) \(\mathrm{CCl}_{4}\)

Step-by-Step Solution

Answer

(a) London dispersion; (b) Hydrogen bonding; (c) Dipole-dipole; (d) London dispersion.

Step 1: Understanding Intermolecular Forces

First, let's identify the types of intermolecular forces: London dispersion forces, dipole-dipole forces, and hydrogen bonding. London dispersion is present in all molecules, dipole-dipole in polar molecules, and hydrogen bonding in molecules with N-H, O-H, or F-H bonds.

Step 2: Identify Forces in CO2

\(\mathrm{CO}_2\) is a linear and nonpolar molecule. Therefore, the only intermolecular forces present are London dispersion forces.

Step 3: Identify Forces in CHCl3

\(\mathrm{CHCl}_3\) is a polar molecule because of the asymmetrical distribution of its chlorines and hydrogen. It exhibits London dispersion and dipole-dipole forces.

Step 4: Identify Forces in NH3

\(\mathrm{NH}_3\) is a polar molecule with N-H bonds, allowing for hydrogen bonding. It has London dispersion, dipole-dipole, and hydrogen bonding.

Step 5: Identify Forces in CCl4

\(\mathrm{CCl}_4\) is a tetrahedral and nonpolar molecule, meaning it only has London dispersion forces.

Key Concepts

Understanding London Dispersion Forces

London dispersion forces are the weakest type of intermolecular attraction, yet they are ubiquitous, present in every molecule. They arise due to the momentary fluctuation of electron density in molecules, creating temporary

• positive and negative ends.
• These temporary dipoles can induce dipoles in neighboring molecules, resulting in weak attraction.
• The strength of London dispersion forces increases with larger electron clouds, so larger atoms or molecules typically experience stronger dispersion forces.

In the context of the original exercise, carbon dioxide (\(\mathrm{CO_2}\)) and carbon tetrachloride (\(\mathrm{CCl_4}\)) rely solely on these forces. Both are nonpolar and do not have any other type of intermolecular force acting upon them.
Recognizing the universal nature of London dispersion forces can provide insights into molecular interactions and help predict physical properties such as boiling and melting points.

What Are Dipole-Dipole Forces?

Dipole-dipole forces occur between molecules that have permanent dipoles, meaning the molecule is polar. This polarity arises when molecules have an unequal distribution of electrons due to different elements with differing electronegativities within the molecule.

• These opposite charges create attractive forces between the positive end of one polar molecule and the negative end of another.
• Such interactions are generally stronger than London dispersion forces but weaker than hydrogen bonds.

In the textbook exercise, \(\mathrm{CHCl_3}\) exhibits dipole-dipole forces due to its polar nature. Its asymmetrical shape causes a net dipole moment. Recognizing dipole-dipole forces helps in understanding the characteristics like solubility and reactivity, which vary greatly depending on molecular polarity.

Exploring Hydrogen Bonding

Hydrogen bonding is a special, stronger type of dipole-dipole interaction, which occurs when hydrogen is directly bonded to highly electronegative atoms like nitrogen, oxygen, or fluorine. This bond enables a strong attraction between hydrogen and the lone pairs of electrons on neighboring molecules.

• Hydrogen bonds are significantly stronger than both London dispersion and regular dipole-dipole forces.
• They contribute to higher boiling and melting points compared to similar-sized molecules that do not exhibit hydrogen bonding.

In the exercise, \(\mathrm{NH_3}\) (ammonia) displays hydrogen bonding since it has an N-H bond. This contributes to its relatively high boiling point compared to other small molecules.
Understanding hydrogen bonding is essential for grasping many phenomena in chemistry, from water's unique properties to the complexity of biological macromolecules.