Problem 6 - Chemistry and Chemical Reactivity

Question

Considering intermolecular forces in the pure substance, which of these substances exists as a gas at \(25^{\circ} \mathrm{C}\) and \(1 \mathrm{atm} ?\) (a) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{3}\) (butane) (b) \(\mathrm{CH}_{3} \mathrm{OH}\) (methanol) (c) \(\mathrm{Ar}\)

Step-by-Step Solution

Answer

Argon (\(\mathrm{Ar}\)) is a gas at \(25^{\circ} \mathrm{C}\) and \(1 \mathrm{atm}\).

Step 1: Identify Possible Intermolecular Forces

Examine each substance to determine the types of intermolecular forces present. - Butane (\(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{CH}_{3}\)) is nonpolar and only exhibits London dispersion forces.- Methanol (\(\mathrm{CH}_{3} \mathrm{OH}\)) is polar with hydrogen bonding due to the O-H group, as well as dipole-dipole interactions and dispersion forces.- Argon (\(\mathrm{Ar}\)) is a noble gas and exhibits only London dispersion forces.

Step 2: Assess Strengths of Intermolecular Forces

Recognize that substances with weaker intermolecular forces are more likely to exist as gases. - London dispersion forces are generally weaker than dipole-dipole interactions and hydrogen bonding. - Methanol, with hydrogen bonding, has stronger intermolecular forces compared to butane and argon, which only have dispersion forces.

Step 3: Compare Boiling Points

Factors influencing whether a substance is a gas at a given temperature include its boiling point compared to that temperature.- Substances with boiling points below \(25^{\circ} \mathrm{C}\) will exist as gases at this temperature and pressure.- Butane's boiling point is about \(-1^{\circ} \mathrm{C}\), methanol's boiling point is around \(65^{\circ} \mathrm{C}\), and argon's boiling point is \(-185.8^{\circ} \mathrm{C}\).

Step 4: Conclusion

Substances that have boiling points below \(25^{\circ} \mathrm{C}\) are most likely to be gases at room temperature and atmospheric pressure.- Butane and argon have boiling points below \(25^{\circ} \mathrm{C}\), whereas methanol does not. Among them, argon has the lowest boiling point and is significantly gaseous.- Argon (\(\mathrm{Ar}\)) is a gas at \(25^{\circ} \mathrm{C}\) and \(1 \mathrm{atm}\).

Key Concepts

London Dispersion Forces

London dispersion forces, also known as van der Waals forces, are the weakest type of intermolecular forces. They occur when electron clouds in atoms or molecules, usually nonpolar, experience momentary shifts in their positions, creating temporary dipoles that attract each other. Even though these dipoles are fleeting, they enable molecules to interact.
These forces are universal, present in all molecules, but they are especially significant in nonpolar substances like butane (\(\text{CH}_3\text{CH}_2\text{CH}_2\text{CH}_3\)) and argon (\(\text{Ar}\)), a noble gas. Both substances rely solely on London dispersion forces for intermolecular interactions because they lack permanent dipole moments. The strength of London dispersion forces can increase with the size of the molecules or the number of electrons they possess.

• Always present in all molecules
• Predominant in nonpolar substances
• Strength increases with larger and heavier molecules

Understanding London dispersion forces is crucial for determining the physical state of a substance at a given temperature. Because these forces are generally weak, substances like argon, which rely only on them, tend to exist as gases at room temperature and pressure.

Hydrogen Bonding

Hydrogen bonding is a special type of dipole-dipole interaction that occurs when hydrogen is covalently bonded to a highly electronegative atom, such as nitrogen, oxygen, or fluorine. This bonding results in a significant polarity in the molecule and creates a strong attraction between molecules. Methanol (\(\text{CH}_3\text{OH}\)) is a classic example of a substance capable of hydrogen bonding, due to the presence of the O-H group in its structure.
Hydrogen bonds are much stronger than London dispersion forces. This strength gives substances with hydrogen bonding much higher boiling points compared to those that lack it. In methanol's case, this strong intermolecular force is responsible for its relatively high boiling point of approximately 65°C, which is much above room temperature.

• Strongest form of dipole-dipole interactions
• Requires hydrogen bonded to N, O, or F
• Significantly increases boiling points and influences substance state

Thus, the presence of hydrogen bonds means that methanol remains in the liquid state under ambient conditions while substances that lack such bonds may exist as gases.

Boiling Point Comparison

Comparing boiling points is essential to predict whether a substance will be a gas at a specific temperature. The boiling point is the temperature at which a substance changes from a liquid to a gas. Substances with lower boiling points than the ambient temperature are typically gases under those conditions.
Considering butane (\(\text{CH}_3\text{CH}_2\text{CH}_2\text{CH}_3\)), methanol (\(\text{CH}_3\text{OH}\)), and argon (\(\text{Ar}\)), we observe the following:

• Butane has a boiling point around -1°C, making it a gas at 25°C.
• Methanol has a higher boiling point of around 65°C, so it remains a liquid at 25°C.
• Argon, with a boiling point of -185.8°C, easily remains gaseous at standard room conditions.

The role of intermolecular forces is evident here: stronger forces like hydrogen bonding in methanol lead to a higher boiling point. Weaker forces, such as London dispersion in argon, account for the much lower boiling point, explaining why argon and butane exist as gases at 25°C.