Problem 80 - Chemistry: The Central Science

Question

Ethylene glycol \(\left(\mathrm{HOCH}_{2} \mathrm{CH}_{2} \mathrm{OH}\right)\) is the major component of antifreeze. It is a slightly viscous liquid, not very volatile at room temperature, with a boiling point of \(198^{\circ} \mathrm{C}\). Pentane \(\left(\mathrm{C}_{5} \mathrm{H}_{12}\right)\), which has about the same molecular weight, is a nonviscous liquid that is highly volatile at room temperature and whose boiling point is \(36.1^{\circ} \mathrm{C}\). Explain the differences in the physical properties of the two substances.

Step-by-Step Solution

Answer

The differences in the physical properties of ethylene glycol and pentane can be attributed to their chemical structures, which dictate the types of intermolecular forces they experience. Ethylene glycol has two hydroxyl groups, enabling it to form strong hydrogen bonds with other ethylene glycol molecules. Pentane, on the other hand, is a hydrocarbon without these functional groups and can only form weaker London dispersion forces with its neighbors. Consequently, ethylene glycol is more viscous, less volatile, and has a higher boiling point (198°C) than pentane, which has weaker intermolecular forces and a lower boiling point (36.1°C).

Step 1: Examine the chemical structures of each molecule

Ethylene glycol has a chemical formula of HOCH₂CH₂OH, which means it has two hydroxyl (-OH) groups, one on each end of the molecule. These hydroxyl groups facilitate the formation of hydrogen bonds with other ethylene glycol molecules. Pentane, on the other hand, has a chemical formula of C₅H₁₂. It is a simple hydrocarbon with no functional groups that can form hydrogen bonds like those in ethylene glycol.

Step 2: Identify the types of intermolecular forces

The presence of hydroxyl groups in ethylene glycol allows it to form strong hydrogen bonds with other ethylene glycol molecules, while pentane can only form weaker London dispersion forces (also called van der Waals forces) with other pentane molecules. London dispersion forces are generally weaker than hydrogen bonds because they result from temporary, induced dipoles and are therefore less effective at keeping molecules together at higher temperatures.

Step 3: Relate intermolecular forces to viscosity

The stronger hydrogen bonds in ethylene glycol cause it to have a higher viscosity than pentane, which only contains weaker London dispersion forces. Viscosity is a measure of a fluid's resistance to flow, so the stronger intermolecular forces in ethylene glycol cause its molecules to "stick together" more, making it more resistant to flow and therefore more viscous.

Step 4: Relate intermolecular forces to volatility

Volatility refers to how readily a substance evaporates or turns into a gas. Since ethylene glycol has stronger intermolecular forces due to hydrogen bonding, it is less volatile than pentane, which only has weaker London dispersion forces holding its molecules together. The weaker intermolecular forces in pentane mean its molecules can more easily overcome these forces and escape into the gas phase, making it more volatile at room temperature.

Step 5: Relate intermolecular forces to boiling points

The boiling point of a substance is the temperature at which it transitions from a liquid to a gas. The stronger hydrogen bonds in ethylene glycol require more energy to break, meaning its boiling point is higher (\(198^{\circ} \mathrm{C}\)) than that of pentane, which has weaker London dispersion forces and therefore a lower boiling point (\(36.1^{\circ} \mathrm{C}\)). Since more energy is needed to break the hydrogen bonds in ethylene glycol, its boiling point is significantly higher than that of pentane. In conclusion, the unique physical properties of ethylene glycol and pentane can be attributed to the differences in their chemical structures, specifically the presence of hydrogen bonding in ethylene glycol and the absence of it in pentane. These differences in intermolecular forces directly affect their viscosity, volatility, and boiling points.

Key Concepts

Hydrogen Bonding

Understanding why ethylene glycol has different properties from pentane starts with recognizing the role of hydrogen bonding. This type of intermolecular force occurs when a hydrogen atom covalently bonded to an electronegative atom, like oxygen, interacts with another electronegative atom.

In ethylene glycol ( HOCH₂CH₂OH ), there are -OH (hydroxyl) groups at both ends of the molecule. These groups can form hydrogen bonds with other ethylene glycol molecules.

• Hydrogen bonds are relatively strong compared to other intermolecular forces.
• They create a network between molecules, increasing the energy required to separate them.

This is why you find ethylene glycol as a thick, non-volatile liquid at room temperature, distinctively different from pentane, which cannot form such bonds due to the absence of polar functional groups.

London Dispersion Forces

London dispersion forces are the weakest form of intermolecular attraction, yet they are present in all molecules, regardless if they are polar or non-polar. These forces arise from transient electron cloud distortions that lead to momentary dipoles.

Pentane ( C₅H₁₂ ) is a non-polar molecule that primarily relies on London dispersion forces for intermolecular attractions.

• These forces are weaker because they are temporary and fluctuate with electron movement.
• Their effectiveness decreases as molecules move apart, which happens more readily at higher temperatures.

Hence, without stronger forces like hydrogen bonding, pentane is more volatile and less viscous than ethylene glycol, having a low boiling point due to its inability to sustain significant intermolecular attractions.

Boiling Point

The boiling point of a substance is significantly determined by the strength of its intermolecular forces. It is the temperature where a liquid becomes a gas.

In the case of ethylene glycol and pentane, ethylene glycol’s strong hydrogen bonds mean it requires more energy (higher temperature) to break these bonds. Consequently, its boiling point is much higher at \(198^{\circ} \text{C}\) compared to pentane's \(36.1^{\circ} \text{C}\).

• Stronger intermolecular forces, such as hydrogen bonds, increase boiling points.
• Weaker forces, like London dispersion, lead to lower boiling points as less energy is needed for molecules to enter the gas phase.

Thus, the presence and type of intermolecular force directly influence the boiling point, with hydrogen bonding dominating in ethylene glycol to substantially raise its boiling temperature.