At 63.5 °C, the vapor pressure of H 2 O is 175 torr, and that of ethanol (C 2 H 5 OH) is 400 torr. A solution is made by mixing equal masses of H 2 O and C 2 H 5 OH. (b) Assuming ideal solution behavior, what is the vapor pressure of the solution at 63.5 °C?

Hello everyone. Today. We are being given the following question and asked to solve for it. So it says the vapor pressure of benzene is at 20 degrees Celsius as Tour. While that of chloroform is 1 60 Tour, calculate the vapor pressure of the solution prepared by dissolving equal amounts in molds of the two liquids at 20 degrees Celsius. And we are to assume that there is ideal behavior going on. So the first you want to do is you want to find the vapor pressure of a component. And so by doing that we can see the vapor pressure as you go to the P. A. Times the mole fraction times our partial pressures. And so what we have to do is we have to calculate the vapor pressure of benzene in the solution. And since they're made up with equal molar amounts we can go ahead and say that that is going to be equal to 1/ plus one giving us 0.5 for a mole fraction. And so solving for both Benzene and chloroform, We get 0.5 for both mole fractions for the benzene we use 75 tour as indicated in the question and for chloroform we use 160 tour. And when we saw for both of these for Benzene we get 37.5 tour and a detour with chloroform To find the total pressure, we simply have to add these pressures together. And so we take that 37.5 tour and we added to that tour to get a total of 117. tour for our final answer, I hope this helped until next time.

Ch.13 - Properties of Solutions

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Brown 14th Edition

Ch.13 - Properties of Solutions

Problem 67b

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Chapter 13, Problem 67b

At 63.5 °C, the vapor pressure of H₂O is 175 torr, and that of ethanol (C₂H₅OH) is 400 torr. A solution is made by mixing equal masses of H₂O and C₂H₅OH. (b) Assuming ideal solution behavior, what is the vapor pressure of the solution at 63.5 °C?

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Calculate the moles of H_2O and C_2H_5OH using their respective molar masses.

Determine the mole fraction of each component in the solution.

Use Raoult's Law to calculate the partial vapor pressure of each component: P_i = X_i * P_i^0, where P_i is the partial pressure, X_i is the mole fraction, and P_i^0 is the pure component vapor pressure.

Add the partial pressures of H_2O and C_2H_5OH to find the total vapor pressure of the solution.

Ensure the solution behaves ideally by checking that the sum of the mole fractions equals 1.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Vapor Pressure

Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid or solid phase at a given temperature. It reflects the tendency of particles to escape from the liquid phase into the vapor phase. The higher the vapor pressure, the more volatile the substance. In this question, the vapor pressures of pure water and ethanol at 63.5 °C are provided, which are essential for calculating the vapor pressure of the solution.

Raoult's Law

Raoult's Law states that the vapor pressure of a solvent in a solution is equal to the vapor pressure of the pure solvent multiplied by its mole fraction in the solution. This law applies to ideal solutions, where interactions between different molecules are similar to those between like molecules. In this scenario, Raoult's Law will be used to determine the vapor pressure of the mixed solution of water and ethanol.

Mole Fraction

Mole fraction is a way of expressing the concentration of a component in a mixture, defined as the number of moles of that component divided by the total number of moles of all components in the mixture. It is a crucial factor in Raoult's Law, as it helps to quantify the proportion of each component in the solution. In this case, calculating the mole fractions of water and ethanol will be necessary to find the overall vapor pressure of the solution.

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Related Practice

Open Question

(a) Calculate the vapor pressure of water above a solution prepared by adding 22.5 g of lactose (C12H22O11) to 200.0 g of water at 338 K. (Vapor–pressure data for water are given in Appendix B.)

Open Question

(b) Calculate the mass of ethylene glycol (C2H6O2) that must be added to 1.00 kg of ethanol (C2H5OH) to reduce its vapor pressure by 10.0 torr at 35 °C. The vapor pressure of pure ethanol at 35 °C is 1.00 x 10^2 torr.

Textbook Question

At 63.5 °C, the vapor pressure of H₂O is 175 torr, and that of ethanol (C₂H₅OH) is 400 torr. A solution is made by mixing equal masses of H₂O and C₂H₅OH. (a) What is the mole fraction of ethanol in the solution?

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Textbook Question

At 20 °C, the vapor pressure of benzene (C₆H₆) is 75 torr, and that of toluene (C₇H₈) is 22 torr. Assume that benzene and toluene form an ideal solution. (b) What is the mole fraction of benzene in the vapor above the solution described in part (a)?

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List the following aqueous solutions in order of increasing boiling point: 0.120 m glucose, 0.050 m LiBr, 0.050 m Zn(NO₃)₂.

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