Question

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

The substances with greater intermolecular forces in each pair are: a. OCS b. SeO2 c. H2NCH2CH2NH2 d. H2CO e. CH3OH

Step-by-step solution

Pair a: CO2 or OCS

The structure of CO2 (carbon dioxide) is linear with a symmetrical distribution of electron density around the carbon atom, making it a nonpolar molecule. The only intermolecular force present in CO2 is London dispersion forces. OCS (carbonyl sulfide) is also linear, but it is polar because of the asymmetrical distribution of electron density around the central atom. In addition to London dispersion forces, OCS has dipole-dipole interactions. Thus, OCS has greater intermolecular forces than CO2.

Pair b: SeO2 or SO2

SeO2 (selenium dioxide) and SO2 (sulfur dioxide) are both bent molecules. Both have polar bonds and are polar molecules. However, selenium has a greater atomic size and more electrons than sulfur. This leads to larger London dispersion forces in SeO2 than in SO2. Since both molecules have dipole-dipole interactions and dispersion forces, but SeO2 has greater dispersion forces, SeO2 has greater overall intermolecular forces.

Pair c: CH3CH2CH2NH2 or H2NCH2CH2NH2

CH3CH2CH2NH2 (propylamine) has covalent bonds between carbon and hydrogen atoms, and a polar bond between nitrogen and hydrogen. There is no hydrogen bond within the molecule. H2NCH2CH2NH2 (ethylenediamine) also has covalent bonds between carbon and hydrogen atoms, and polar bonds between nitrogen and hydrogen. However, ethylenediamine can form hydrogen bonds with other molecules through its two nitrogen atoms. Therefore, H2NCH2CH2NH2 has greater intermolecular forces due to the presence of hydrogen bonding.

Pair d: CH3CH3 or H2CO

CH3CH3 (ethane) is a nonpolar hydrocarbon molecule, which means it only has London dispersion forces as its intermolecular forces. H2CO (formaldehyde) is a polar molecule due to the presence of a polar carbon-oxygen bond and a trigonal planar molecular geometry. Since it is polar, H2CO has both dispersion forces and dipole-dipole interactions. Consequently, H2CO has greater intermolecular forces than CH3CH3.

Pair e: CH3OH or H2CO

CH3OH (methanol) is a polar molecule with a polar bond between oxygen and hydrogen atoms. It can form hydrogen bonds with other molecules due to the presence of a hydrogen atom bonded to an oxygen atom. H2CO (formaldehyde), as discussed previously, is a polar molecule with dipole-dipole forces and London dispersion forces. However, it cannot form hydrogen bonds. Hence, CH3OH has greater intermolecular forces than H2CO due to the presence of hydrogen bonding. In summary, the substances with greater intermolecular forces in each pair are: a. OCS b. SeO2 c. H2NCH2CH2NH2 d. H2CO e. CH3OH

Key concepts

London Dispersion Forces

London dispersion forces are a type of intermolecular force that occurs in all molecules, whether they are polar or nonpolar. They arise from temporary fluctuations in the electron cloud of a molecule, which create an immediate dipole. This dipole can induce a similar dipole in a neighboring molecule, leading to an attraction between them. These forces are generally weak, but their strength increases with the size of the electron cloud. The more electrons a molecule has, the stronger its London dispersion forces.

Think of these forces like a brief handshake between two people passing by; they're quick and usually not very strong, but they do exist. The larger the molecule, like SeO2 versus SO2, the stronger the dispersion force due to having more electrons to shuffle around. In the context of intermolecular forces, CO2 would rely solely on these weak forces, while something more complex like OCS would have a bit more going on due to its polar nature.

Dipole-Dipole Interactions

Dipole-dipole interactions occur in polar molecules. These forces result from the attraction between the positive end of one molecule and the negative end of another. It's like a magnetic effect where opposites attract. This interaction is stronger than London dispersion forces but weaker than hydrogen bonding.

For example, in the analysis of OCS for intermolecular forces, dipole-dipole interactions play an essential role due to its molecular polarity. Compared to CO2, which lacks these interactions, OCS experiences a more robust attraction between its molecules. Similarly, formaldehyde (H2CO) benefits from such interactions, making its intermolecular forces greater compared to nonpolar molecules like ethane (CH3CH3).

• These interactions are generally stronger in molecules that have a large difference in electronegativity between atoms.
• The effectiveness of these forces depends on the orientation of the polar molecules to one another.

Hydrogen Bonding

Hydrogen bonding is the strongest of the three types of intermolecular forces discussed here. It occurs in molecules where hydrogen is directly bonded to very electronegative atoms like nitrogen, oxygen, or fluorine. These atoms pull the shared electrons closer, leaving hydrogen relatively exposed and positive. This positive hydrogen can then attract the lone pair electrons of other electronegative atoms in nearby molecules.

A perfect example is methanol (CH3OH), where hydrogen bonding is potent due to its -OH group. This type of bonding is a strong, directional force that significantly increases a substance's boiling and melting points. In the case of ethylenediamine (H2NCH2CH2NH2), the two nitrogen atoms allow the formation of hydrogen bonds, enhancing its intermolecular strength compared to molecules that can't hydrogen bond, such as propylamine.

• Hydrogen bonds are crucial for the unique properties of water and biological molecules like proteins and DNA.
• They are much stronger than dipole-dipole interactions and significantly influence physical properties.