Open Question
What is the required concentration (in percent by mass) for an aqueous ethylene glycol (C2H6O2) solution to have a boiling point of 104.0 °C?
14. Solutions
Boiling Point Elevation
Multiple Choice
100.0 g of urea (CO(NH2)2) is dissolved in 550.0 g of liquid chloroform. Given that urea is a covalent compound with a molar mass of 60.07 g/mol and the boiling point elevation constant (Kb) for chloroform is 3.63 °C/m, what is the boiling point of the solution if the normal boiling point of chloroform is 61.2 °C?
A
61.2 °C
B
64.5 °C
C
60.0 °C
D
62.8 °C
Verified step by step guidance1
Calculate the number of moles of urea using its molar mass. Use the formula: \( \text{moles of urea} = \frac{\text{mass of urea}}{\text{molar mass of urea}} \). Substitute the given values: \( \text{mass of urea} = 100.0 \text{ g} \) and \( \text{molar mass of urea} = 60.07 \text{ g/mol} \).
Determine the molality of the solution. Molality (m) is defined as the number of moles of solute per kilogram of solvent. Use the formula: \( m = \frac{\text{moles of urea}}{\text{mass of chloroform in kg}} \). Convert the mass of chloroform from grams to kilograms: \( 550.0 \text{ g} = 0.550 \text{ kg} \).
Calculate the boiling point elevation using the formula: \( \Delta T_b = K_b \times m \), where \( K_b \) is the boiling point elevation constant for chloroform and \( m \) is the molality. Substitute the given \( K_b = 3.63 \text{ °C/m} \) and the calculated molality.
Determine the new boiling point of the solution by adding the boiling point elevation to the normal boiling point of chloroform. Use the formula: \( \text{new boiling point} = \text{normal boiling point} + \Delta T_b \). Substitute the normal boiling point of chloroform: \( 61.2 \text{ °C} \).
Review the calculated boiling point and compare it with the provided options to ensure the correct answer is identified.
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