Problem 32 - A Complete Resource Book in Chemistry for JEE Main

Question

Which pair of the following will not form an ideal solution? (a) \(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{Br}+\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{I}\) (b) \(\mathrm{H}_{2} \mathrm{O}+\mathrm{C}_{4} \mathrm{H}_{9} \mathrm{OH}\) (c) \(\mathrm{CCl}_{4}+\mathrm{SiCl}_{4}\) (d) \(\mathrm{C}_{6} \mathrm{H}_{14}+\mathrm{C}_{\gamma} \mathrm{H}_{16}\)

Step-by-Step Solution

Answer

(b) \(\mathrm{H}_{2} \mathrm{O}+\mathrm{C}_{4} \mathrm{H}_{9} \mathrm{OH}\) will not form an ideal solution.

Step 1: Understanding Ideal Solutions

An ideal solution obeys Raoult's Law, which means the interactions between different molecules in the solution must be similar to the interactions between like molecules. This occurs when the intermolecular forces (like hydrogen bonding, Van der Waals forces) are of similar strength for both components of the mixture.

Step 2: Analyzing Each Pair for Similar Interactions

Let's check each pair to see if they likely form ideal solutions by comparing their interactions:(a) \(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{Br}+\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{I}\): Both are haloalkanes with similar molecular structures, suggesting similar Van der Waals dispersion forces.(b) \(\mathrm{H}_{2} \mathrm{O}+\mathrm{C}_{4} \mathrm{H}_{9} \mathrm{OH}\): Water and butanol can form hydrogen bonds, but their strength and nature differ, as water has stronger hydrogen bonds.(c) \(\mathrm{CCl}_{4}+\mathrm{SiCl}_{4}\): Both are nonpolar, with similar dispersion forces, due to their similar molecular shapes.(d) \(\mathrm{C}_{6} \mathrm{H}_{14}+\mathrm{C}_{\gamma} \mathrm{H}_{16}\): These are hydrocarbons with similar Van der Waals forces.

Step 3: Identifying the Non-Ideal Pair

From the analysis, pairs (a), (c), and (d) involve components that have similar types and strengths of intermolecular forces, making them likely to form ideal solutions. However, pair (b) has a significant difference in the strength of hydrogen bonding between water and butanol, leading to non-ideal behavior.

Step 4: Conclusion

The pair that will not form an ideal solution is (b) \((\mathrm{H}_{2} \mathrm{O}+\mathrm{C}_{4} \mathrm{H}_{9} \mathrm{OH})\), due to the significantly different interactions from strong hydrogen bonding in water compared to the weaker interactions possible in butanol.

Key Concepts

Raoult's Law

Raoult's Law is fundamental to understanding ideal solutions. It states that the partial vapor pressure of each component in a solution is directly proportional to its mole fraction.
For an ideal solution, the interactions between unlike molecules must mirror the interactions between like molecules.
A good analogy is that of dancers at a party. If everyone dances comfortably with everyone else, the mixture behaves ideally; otherwise, things become complicated!

• This law helps predict how the vapor pressure of a solution compared to pure substances behaves.
• It applies under the assumption that the solution's components have similar intermolecular forces.
• Raoult's Law is typically accurate for solutions of similar substances, such as benzene and toluene.

Intermolecular Forces

Intermolecular forces are the forces that exist between molecules, affecting their physical properties and interactions in solutions.
These forces include Van der Waals forces, dipole-dipole interactions, and hydrogen bonds.
Understanding these forces helps in predicting whether a solution will be ideal.

• Van der Waals forces: These are weak attractions between molecules due to transient shifts in electron density. They include dispersion forces, which are particularly important in nonpolar molecules.
• Dipole-dipole interactions: Occur between molecules with permanent dipoles, like hydrogen fluoride. They are stronger than dispersion forces and are significant in polar substances.
• Hydrogen bonding: A special case of dipole interactions, exceptionally strong, often seen in compounds like water or ammonia.

Molecules with similar forces mix more easily into an ideal solution as their energy states and interactions align closely.

Hydrogen Bonding

Hydrogen bonding is a strong type of dipole-dipole interaction and plays a critical role in the behavior of many compounds.
It occurs when hydrogen is covalently bonded to highly electronegative elements such as oxygen, nitrogen, or fluorine.
This results in unique properties, especially for water, and influences solution behavior.

• Strength: Hydrogen bonds are stronger than other dipole interactions owing to their highly directional nature. However, capillary action and surface tension depend on them as well.
• Impact on Ideal Solutions: When two substances exhibit drastically different hydrogen bonding capabilities, forming an ideal solution is less likely due to significant differences in interaction strengths.
• Example with Water: Water's extensive hydrogen bonding network leads to high boiling and melting points compared to its molecular size.

In the given problem, the mixture of butanol with water does not produce an ideal solution due to water's significantly stronger hydrogen bonding ability compared to butanol. This illustrates how hydrogen bonding can affect solution behavior.