This chemistry question involves key chemical concepts and calculations. The detailed solution below walks through each step, from identifying the reaction type to computing the final answer.
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To find the vapor pressure of pure B, we will use Raoult's Law for ideal solutions, which states that the total vapor pressure of a solution is the sum of the partial pressures of its components. The partial pressure of each component is its mole fraction multiplied by its pure vapor pressure. Step 1: Calculate the total moles in the solution. Number of moles of A (n_A) = 0.650 mol Number of moles of B (n_B) = 0.250 mol n_total = n_A + n_B = 0.650 mol + 0.250 mol = 0.900 mol Step 2: Calculate the mole fraction of each component. X_A = (n_A)/(n_total) = 0.650 mol0.900 mol = 0.7222 X_B = (n_B)/(n_total) = 0.250 mol0.900 mol = 0.2778 Step 3: Apply Raoult's Law. The total vapor pressure of the solution (P_total) at its normal boiling point is 1 atm. The vapor pressure of pure A (P_A^) is 0.700 atm. Raoult's Law states: P_total = X_A · P_A^ + X_B · P_B^ Substitute the known values: 1.000 atm = (0.7222 × 0.700 atm) + (0.2778 × P_B^) Step 4: Solve for the vapor pressure of pure B (P_B^). 1.000 atm = 0.50554 atm + 0.2778 · P_B^ Subtract 0.50554 atm from both sides: 1.000 atm - 0.50554 atm = 0.2778 · P_B^ 0.49446 atm = 0.2778 · P_B^ Divide by 0.2778: P_B^ = 0.49446 atm0.2778 = 1.780 atm The vapor pressure of pure B at this temperature is 1.780 atm. 3 done, 2 left today. You're making progress.