Question

Consider the following molecules: \(\mathrm{CH}_{3} \mathrm{OH}, \mathrm{CH}_{4}, \mathrm{H}_{2} \mathrm{O}, \mathrm{C}_{2} \mathrm{H}_{6}\). a. Draw the Lewis structure for each molecule, and indicate whether each is polar or nonpolar. b. At room temperature, two of these compounds exist as a liquid, and two of these compounds exist as a gas. State which two compounds are liquids at room temperature and which two are gases. Be sure to justify your answer completely. Include discussions of intermolecular forces in your response. c. Rank the compounds from lowest to highest boiling point. Justify your answer.

Short answer

The Lewis structures and polarities of the molecules are as follows: CH3OH is polar, CH4 is nonpolar, H2O is polar, and C2H6 is nonpolar. At room temperature, CH3OH and H2O exist as liquids due to strong hydrogen bonding, while CH4 and C2H6 exist as gases due to weak London dispersion forces. The boiling point order is CH4 < C2H6 < CH3OH < H2O, with the strength of intermolecular forces increasing from CH4 to H2O.

Step-by-step solution

Draw Lewis structures and determine polarity

For each molecule, we will draw the Lewis structure by arranging the atoms and distributing the valence electrons to satisfy the octet rule for each atom (except hydrogen, which should have two electrons). 1. CH3OH: Oxygen (6 valence electrons) is the central atom bonded to one carbon and one hydrogen. Carbon has four valence electrons and bonds with three hydrogen atoms. The Lewis structure of CH3OH has a single unshared pair of electrons on oxygen. This creates a dipole moment as oxygen is more electronegative than carbon and hydrogen, making CH3OH a polar molecule. 2. CH4: Carbon (4 valence electrons) bonds with four hydrogen atoms, each having one valence electron. Since all the four bonding electrons shared between carbon and hydrogen are evenly distributed, no dipole moment is created, making CH4 a nonpolar molecule. 3. H2O: Oxygen (6 valence electrons) bonds with two hydrogen atoms, each having one valence electron. The Lewis structure of H2O shows that oxygen has two unshared pairs of electrons. Oxygen being more electronegative than hydrogen creates a dipole moment, making H2O a polar molecule. 4. C2H6: Ethane contains two carbon atoms bonded together with six hydrogen atoms bonded around them. The Lewis structure of C2H6 shows that electrons are evenly distributed among the atoms. No dipole moment is created, making C2H6 a nonpolar molecule.

Identify the state at room temperature and discuss intermolecular forces

At room temperature: 1. CH3OH (Methanol) has hydrogen bonding due to its polar nature. This force is stronger than other intermolecular forces, causing methanol to be a liquid at room temperature. 2. CH4 (Methane) only has weak London dispersion forces due to its non-polar nature. These forces are not strong enough to keep the molecules together as a liquid, making methane a gas at room temperature. 3. H2O (Water) also has hydrogen bonding due to its polar nature. This strong force causes water to be a liquid at room temperature. 4. C2H6 (Ethane) has weak London dispersion forces due to its non-polar nature. As a result, ethane is a gas at room temperature. So, at room temperature, CH3OH and H2O are liquids, and CH4 and C2H6 are gases.

Rank the compounds based on boiling points

The boiling points of the compounds will depend on the strength of their intermolecular forces. The greater the intermolecular forces, the higher the boiling point. Here is the order based on intermolecular forces: 1. CH4 - weakest London dispersion forces (nonpolar) 2. C2H6 - slightly stronger London dispersion forces due to more electrons (nonpolar) 3. CH3OH - hydrogen bonding (polar) 4. H2O - stronger hydrogen bonding due to higher electronegativity difference (polar) Therefore, the order of the compounds from lowest to highest boiling point is CH4 < C2H6 < CH3OH < H2O.

Key concepts

Intermolecular Forces

Understanding intermolecular forces (IMFs) is crucial when analyzing how molecules interact with one another. These forces are what hold molecules together in the liquid or solid states and significantly influence properties like boiling points, melting points, and solubilities.

Types of Intermolecular Forces

There are several types of IMFs, with varying strengths that affect molecular behavior:

• London Dispersion Forces: Present in all molecules, these are the weakest forces caused by temporary dipoles in electron clouds.
• Dipole-Dipole Interactions: These occur between polar molecules due to permanent dipoles attracting each other.
• Hydrogen Bonds: A special type of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms (N, O, or F), resulting in very strong attractions.

Certain molecules, such as water (H2O), exhibit strong hydrogen bonds because they have a polar structure where hydrogen is bound to a highly electronegative oxygen atom. This significantly increases the boiling point.

Molecules such as methane (CH4) and ethane (C2H6), which are nonpolar, only have London dispersion forces, leading to much lower boiling points. When answering questions about intermolecular forces, focus on the type of atoms involved, their electronegativity, and the molecule's shape, which can affect the distribution of charge and the resultant IMFs.

Boiling Point Trends

The boiling point of a substance is an important physical property that can be impacted by molecular structure and intermolecular forces.

Factors Affecting Boiling Points

The strength and type of IMFs directly correlate with boiling points:

• Stronger IMFs lead to higher boiling points because more energy is needed to separate the molecules.
• Molecules with hydrogen bonding typically have higher boiling points than those with just dipole-dipole interactions or London dispersion forces.
• The molecular mass and shape can also influence boiling points due to increased London dispersion forces in larger and more complex molecules.

When considering the provided molecules, methanol (CH3OH) and water (H2O) have higher boiling points than methane (CH4) and ethane (C2H6) due to their ability to form hydrogen bonds. Polarity plays a vital role as well; polar molecules with strong dipole moments tend to have higher boiling points than nonpolar molecules of similar size and shape.

Remember, when predicting boiling point trends, pay close attention to the types of intermolecular forces present and the molecular weights, as these factors dramatically influence a substance's ability to transition to a gaseous state at a given temperature.

Molecular Geometry

Molecular geometry, or the three-dimensional shape of a molecule, affects physical and chemical properties, such as reactivity, polarity, and phases of matter at given temperatures.

Impact of Molecular Geometry

The orientation of atoms within a molecule dictates its geometry, which can be linear, bent, tetrahedral, trigonal planar, and so on. The specific shape:

• Dictates the distribution of electrons around the molecule, influencing its polarity.
• Alters the type and strength of intermolecular forces between molecules.
• Can affect how molecules pack together in the solid state, which in turn affects melting points and boiling points.

For instance, water has a bent geometry that leads to a highly polar molecule with strong hydrogen bonds, thus a high boiling point. Methane and ethane have tetrahedral shapes, but because they are nonpolar, their boiling points are much lower.

Considering molecular geometry in conjunction with electronegativity and bond polarity is essential when determining a molecule's properties. This integrated approach gives a holistic understanding of why substances behave the way they do in different environments.