Multiple Choice
The osmotic pressure of a solution containing 7.0 g of insulin per liter is 23 torr at 25ºC. What is the molar mass of insulin? (1 atm = 760 torr)
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9. Solutions
Osmotic Pressure
Multiple Choice
The osmotic pressure of blood is 5950.8 mmHg at 41ºC. What mass of glucose, C6H12O6, is needed to prepare 5.51 L of solution. The osmotic pressure of the glucose solution is equal to the osmotic pressure of blood.
A
54.7 g
B
0.304 g
C
419 g
D
302 g
0 Comments
Verified step by step guidance1
Start by understanding the formula for osmotic pressure: \( \Pi = iMRT \), where \( \Pi \) is the osmotic pressure, \( i \) is the van't Hoff factor (which is 1 for glucose as it does not ionize), \( M \) is the molarity, \( R \) is the ideal gas constant (0.0821 L·atm/mol·K), and \( T \) is the temperature in Kelvin.
Convert the temperature from Celsius to Kelvin: \( T = 41 + 273.15 = 314.15 \) K.
Convert the osmotic pressure from mmHg to atm: \( \Pi = \frac{5950.8 \text{ mmHg}}{760 \text{ mmHg/atm}} \approx 7.83 \text{ atm} \).
Rearrange the osmotic pressure formula to solve for molarity \( M \): \( M = \frac{\Pi}{iRT} \). Substitute the values: \( M = \frac{7.83}{1 \times 0.0821 \times 314.15} \).
Calculate the mass of glucose needed using the molarity: \( M = \frac{\text{moles of solute}}{\text{volume of solution in liters}} \). Use the molar mass of glucose (180.18 g/mol) to find the mass: \( \text{mass} = \text{moles} \times 180.18 \text{ g/mol} \).
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