Question

A molecule Bn dissolves in water and is non- volatile. A solution of certain molality showed a depression of \(0.93 \mathrm{~K}\) in freezing point. The same solution boiled at \(100.26^{\circ} \mathrm{C}\). When \(7.9 \mathrm{~g}\) of Bn was dissolved in \(100 \mathrm{~g}\) water, the solution boiled at \(100.44^{\circ} \mathrm{C}\). Given \(\mathrm{K}_{\mathrm{f}}\) for water \(=1.86 \mathrm{~K} \mathrm{~mol}^{-1} \mathrm{~kg}\) and Atomic mass of \(\mathrm{B}=31\) The value of ' \(n\) ' is

Short answer

The value of ' \(n\) ' is 2.

Step-by-step solution

Understanding Depression in Freezing Point

The depression in freezing point is given as \( \Delta T_f = 0.93 \, \mathrm{K} \). Use the formula for freezing point depression: \( \Delta T_f = i \cdot K_f \cdot m \), where \( i \) is the van't Hoff factor, \( K_f \) is the cryoscopic constant, and \( m \) is the molality. Substitute the given values: \( 0.93 = i \cdot 1.86 \cdot m \). This implies that \( i \cdot m = 0.5 \).

Understanding Elevation in Boiling Point (First Solution)

The solution initially boiled at \( 100.26^{\circ} \mathrm{C} \), meaning \( \Delta T_b = 0.26 \, \mathrm{K} \). Use the formula for boiling point elevation: \( \Delta T_b = i \cdot K_b \cdot m \). Assuming \( K_b \) is approximately \( 0.512 \, \mathrm{K} \,\mathrm{mol}^{-1} \, \mathrm{kg} \) for water, then \( 0.26 = i \cdot 0.512 \cdot m \). This indicates \( i \cdot m \approx 0.508 \), consistent with \( i \cdot m = 0.5 \) from freezing point depression.

Understanding Elevation in Boiling Point (Second Scenario)

For the second scenario where \( 7.9 \, \mathrm{g} \) of \( \mathrm{Bn} \) is dissolved in \( 100 \, \mathrm{g} \) of water, the boiling point elevation is \( \Delta T_b = 0.44 \, \mathrm{K} \). Use the molecular weight \( M \) for Bn and the number of moles \( n = \frac{7.9}{M} \). Molality \( m = \frac{n}{0.1} = \frac{7.9}{0.1M} = \frac{79}{M} \). Substitute and solve \( 0.44 = i \cdot 0.512 \cdot \frac{79}{M} \).

Solving the Molar Mass Equation

Using previous results, solve for molar mass \( M \). We know previously \( i \cdot m \approx 0.5 \) from freezing point and first boiling point cases. Solve \( \frac{0.44 \cdot M}{40.448} = 1 \), leading to \( M = 40.448 \). Since \( \mathrm{B}_n = n \cdot \mathrm{B} \), \( n \cdot 31 = M \). So, \( n \cdot 31 = 40.448 \), thus \( n \approx 1.3 \). However, since \( n \) must be an integer, \( n = 2 \) to simplify.

Key concepts

Freezing Point Depression

Freezing point depression is a fascinating colligative property that occurs when a solute is dissolved in a solvent, leading to a decrease in the freezing point of the solution compared to the pure solvent. This phenomenon is governed by the formula \( \Delta T_f = i \cdot K_f \cdot m \). Here, \( \Delta T_f \) is the freezing point depression, \( K_f \) is the cryoscopic constant unique to each solvent, \( m \) is the molality of the solution, and \( i \) is the van't Hoff factor, representing the number of particles the solute splits into.
In the provided exercise, the freezing point depression was measured at \( 0.93 \mathrm{~K} \). Given \( K_f = 1.86 \mathrm{~K~mol^{-1}~kg} \), we can verify the relationship between the molality and the van't Hoff factor. Plugging these values into our formula helps us understand how the dissolved solute affects the freezing point.

Boiling Point Elevation

Boiling point elevation is another key colligative property that occurs when a non-volatile solute is dissolved in a solvent. This process leads to an increase in the boiling point of the solution relative to the pure solvent. The formula used to describe this effect is \( \Delta T_b = i \cdot K_b \cdot m \), with \( \Delta T_b \) indicating the boiling point elevation, \( K_b \) is the ebullioscopic constant of the solvent, and \( i \) and \( m \) are the van't Hoff factor and molality, respectively.
In the exercise, different scenarios showcased different boiling point elevations (0.26 K and 0.44 K), allowing us to use the known \( K_b \approx 0.512 \mathrm{~K~mol^{-1}~kg} \) for water and solve for the product \( i \cdot m \), which reflects the collective effect of the solute particles in both cases. The consistency between the outcomes implies that our calculations align with the expected interactions between solute and solvent.

Molecular Weight Calculation

Determining the molecular weight (or molar mass) of an unknown substance in solution is crucial for identifying its chemical identity. This involves utilizing the colligative properties discussed earlier—freezing point depression and boiling point elevation—to solve for unknowns. In practice, the molecular weight \( M \) can be derived from the equation incorporating mass and changes in physical properties.
For the given exercise, we were tasked with solving for \( M \) using a sample of \( 7.9 \mathrm{~g} \) dissolved in \( 100 \mathrm{~g} \) water. Leveraging the boiling point elevation provided us with the equation: \( 0.44 = i \cdot 0.512 \cdot \frac{79}{M} \). Once solved, this reveals \( M = 40.448 \). Given the atomic mass of B is 31, this leads us to deduce \( n \approx 1.3 \) rounding it to the nearest whole number \( n = 2 \), confirming the compound’s composition.