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

Question

Which pair of the following will not form an ideal solution? (a) \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{Br}+\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{I}\) (b) \(\mathrm{H}_{2} \mathrm{O}+\mathrm{C}_{4} \mathrm{H}_{9} \mathrm{OH}\) (c) \(\mathrm{CCl}_{4}+\mathrm{SiCl}_{4}\) (d) \(\mathrm{C}_{6} \mathrm{H}_{14}+\mathrm{C}_{7} \mathrm{H}_{16}\)

Step-by-Step Solution

Answer

(b) \(\mathrm{H}_{2}\mathrm{O}+\mathrm{C}_{4}\mathrm{H}_{9}\mathrm{OH}\).

Step 1: Understanding Ideal Solutions

An ideal solution follows Raoult's Law at all concentrations and temperatures. For a solution to be ideal, the interactions between the particles of different components must be similar to the interactions between the particles of each component.

Step 2: Identifying Types of Interactions

Examine each pair for the type of molecules: - (a) \(5 ext{ and } 5+\mathrm{I}\) are similar in molecular size and polarity. - (b) \(\mathrm{H}_{2}\mathrm{O}\) is polar with strong hydrogen bonds, while \(\mathrm{C}_{4}\mathrm{H}_{9}\mathrm{OH}\) is less polar, having weaker hydrogen bonds. - (c) \(\mathrm{CCl}_{4}\) and \(\mathrm{SiCl}_{4}\) are similar in molecular size and dispersion forces. - (d) \(\mathrm{C}_{6}\mathrm{H}_{14}\) and \(\mathrm{C}_{7}\mathrm{H}_{16}\) are similar in molecular size and dispersion forces.

Step 3: Comparing Molecular Interactions

Compare interactions: - (a) Both molecules are non-polar, leading to similar van der Waals forces. - (b) One molecule is highly polar and forms hydrogen bonds; the other, less so, results in substantial interaction differences. - (c) Both molecules exhibit similar non-polar, dispersive interactions, making them likely to form ideal solutions. - (d) Both are non-polar hydrocarbons, suggesting similar intermolecular interactions.

Step 4: Concluding Which Pair Does Not Form Ideal Solution

Since the interactions between \(\mathrm{H}_{2}\mathrm{O}\) and \(\mathrm{C}_{4}\mathrm{H}_{9}\mathrm{OH}\) are markedly different (due to the difference in hydrogen bonding strength and polarity), this pair is less likely to form an ideal solution.

Key Concepts

Raoult's Law

Raoult's Law is a key principle in understanding ideal solutions. It states that "the partial vapor pressure of each component in an ideal solution is directly proportional to its mole fraction present in the solution." This means that in a mixture, each component behaves as if it were "alone," influencing the total vapor pressure according to its concentration.

Ideal solutions perfectly follow Raoult's Law without deviation, as their components have similar interactions with one another compared to the interactions they have within themselves.

• If components A and B form an ideal solution, then the mixture's vapor pressure is a straightforward sum of their individual vapor pressures.
• The law works well when the intermolecular forces between dissimilar molecules (A-B) are similar to those of similar molecules (A-A or B-B).

When solutions deviate from Raoult’s Law, it often indicates significant differences in these interactions, suggesting non-ideal behavior.

Intermolecular Forces

Intermolecular forces are interactions between molecules that determine many physical properties of substances, including solubility and boiling points. They include:

• Dispersion Forces: Also known as London forces, these are weak attractions present in all molecules. They arise from temporary shifts in electron density, creating instantaneous dipoles.
• Dipole-Dipole Interactions: These occur between polar molecules due to the positive end of one molecule being attracted to the negative end of another.
• Hydrogen Bonds: A special type of dipole-dipole interaction occurring when hydrogen is bonded to a highly electronegative atom, such as oxygen, nitrogen, or fluorine. These are stronger than regular dipole-dipole forces.

Understanding these forces helps us predict solution behavior. For example, substances with similar intermolecular forces are more likely to form ideal solutions because their interactions within the mixture resemble those in pure states.

Hydrogen Bonding

Hydrogen bonding is a particularly strong type of intermolecular force. It occurs when hydrogen is covalently bonded to a highly electronegative atom like oxygen, nitrogen, or fluorine, creating a strong dipole.

Hydrogen bonds have profound effects on the properties of substances, influencing both boiling and melting points due to the increased energy required to break these bonds. In solutions, significant differences in hydrogen bonding capabilities can lead to non-ideal behavior.

• Examples: Water (H_2O) is well-known for its extensive hydrogen bonding, which accounts for many of its unique physical properties such as high boiling point and surface tension.
• Role in Solutions: When mixing substances, if one component forms strong hydrogen bonds while the other does not, the solution tends to deviate from ideal behavior. This deviation is due to the significant differences in intermolecular interactions.

In the exercise, the combination of water and butanol demonstrates this principle, as water's strong hydrogen bonds differ greatly from the weaker interactions in butanol, leading to non-ideal solution characteristics.