Question

Suppose you use a miscalibrated thermometer that reads 1.0 degree C higher than the actual temperature throughout the entire experiment. Consider how the miscalibration affects each temperature measurement, and then determine what happens to your T_f values. What would be the overall effect on your molecular weight calculations? Suppose a small amount of your solid sample is lost and does not get dissolved in the solvent. How does this affect your number of solute moles and your molality? What happens to T_f? What would be the overall effect on your molecular weight calculations? Suppose a small amount of the cyclohexane vaporizes, so that there is actually less solvent mass than what you had recorded. How does this affect your molality and your T_f? What happens to your calculations for the number of moles? What would be the overall effect on your molecular weight calculations? Suppose a small amount of your dissolved solute dissociate into ions. This makes "i" (the Van't Hoff factor) slightly larger than 1, but less than 2 (for complete dissociation). Using T = ik_f m, determine what happens to your T_f. How are your calculations for molality, moles, and molecular weight affected if the "i" is neglected and assumed to be equal to 1? (You may refer to the discussion of sodium chloride on page 183 in your lab manual, as well as to "BP/FP of Ionic Solutions in the chapter 12 class notes.) Suppose that your solvent T_f is measured correctly, but you interpret your graphs so that the solution T_f'S values are higher than actual. How does this affect your T_f values? How does this affect your calculations for molality, moles, and molecular weight?

Answer 1

1. With each temperature measurement it will read higher temperature by 1 ^oC.

        Suppose actual temperature = T ^oC

    Read Temperature by miscalibrated thermometer = (T + 1 )^oC

    We know that :

              ΔT = T_f° –T_f = K_f m

where T_f° and T_f are the freezing-point temperatures of the pure solvent and the solution, respectively and K_f is the freezing point depression constant which is a function of the solvent not the solute. m is the molality of the solution, which is defined as the number of moles of solute per 1000 g of solvent. Thus with each temperature measurement, T_f° and T_f will read 1^oC higher but overall no change in ΔT.

ΔT =( T_f° + 1) – (T_f + 1)

       ΔT = T_f° –T_f

Therefore, with no change in ΔT, molality would be the same and no effect on molecular weight as molality can be expressed in terms of molecular weight.

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