A solution contains 4.08 g of chloroform (CHCl3) and 9.29 g of acetone (CH3COCH3). The vapor pressures at 35 °C of pure chloroform and pure acetone are 295 torr and 332 torr, respectively. Assuming ideal behavior, calculate the vapor pressures of each of the components and the total vapor pressure above the solution. The experimentally measured total vapor pressure of the solution at 35 °C is 312 torr. Is the solution ideal? If not, what can you say about the relative strength of chloroform–acetone interactions compared to the acetone–acetone and chloroform–chloroform interactions?

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1. Calculate the mole fraction of chloroform (CHCl_3) in the solution. Use the formula: \( \text{mole fraction of CHCl}_3 = \frac{\text{moles of CHCl}_3}{\text{moles of CHCl}_3 + \text{moles of CH}_3\text{COCH}_3} \).
2. Calculate the mole fraction of acetone (CH_3COCH_3) in the solution. Use the formula: \( \text{mole fraction of CH}_3\text{COCH}_3 = \frac{\text{moles of CH}_3\text{COCH}_3}{\text{moles of CHCl}_3 + \text{moles of CH}_3\text{COCH}_3} \).
3. Use Raoult's Law to calculate the partial vapor pressure of chloroform: \( P_{\text{CHCl}_3} = \text{mole fraction of CHCl}_3 \times P^0_{\text{CHCl}_3} \), where \( P^0_{\text{CHCl}_3} \) is the vapor pressure of pure chloroform.
4. Use Raoult's Law to calculate the partial vapor pressure of acetone: \( P_{\text{CH}_3\text{COCH}_3} = \text{mole fraction of CH}_3\text{COCH}_3 \times P^0_{\text{CH}_3\text{COCH}_3} \), where \( P^0_{\text{CH}_3\text{COCH}_3} \) is the vapor pressure of pure acetone.
5. Calculate the total vapor pressure of the solution by adding the partial pressures: \( P_{\text{total}} = P_{\text{CHCl}_3} + P_{\text{CH}_3\text{COCH}_3} \). Compare this with the experimentally measured total vapor pressure to determine if the solution is ideal. If the calculated pressure is different from the experimental value, discuss the relative strength of interactions.