1. Would it be advisable to determine the freezing point of pure p-dichlorobenzene with one thermometer and the freezing point of the solution with a different one? Explain.
2. If 0.150 mole of a nonvolatile non-electrolyte solute is present in 1200g of water, what are the ideal melting and boiling points of the solution?
3. If the solute above is K2SO4 instead of a covalent solute, what are the ideal (assuming no attraction between oppositely charged ions) melting and boiling points?
4. What weight of ethylene glycol, C2H6O2, must be added to a liter of water to yield a solution that freezes at -15°C?
1. Absolutely not a good idea to use different thermometers. You want the DIFFERENCE in temperature. If either thermometer reads only slightly different than the other, then that error will be the final calculation.
2. m = mols/kg solvent = 0.150/1.2 = ?
Then delta T = Kf*m
Solve for delta T.
For the ideal freezing point subtract delta T from 0 C.
For the ideal b.p., delta = Kb*m add delta T to 100 C.
3. Change delta T = k*m to
delta T = i*Kf*m where i = 3.
4. delta T = Kf*m
You know delta T and Kf for water, solve for m
Then m = mols/kg solvent.
You know m and kg solvent = 1 kg, solve for mols ethylene clycol.
Then mol x molar mass = grams eth glycol
1. It would not be advisable to determine the freezing point of pure p-dichlorobenzene with one thermometer and the freezing point of the solution with a different one.
To accurately determine the freezing point of a substance, it is crucial to use the same thermometer for both measurements. Different thermometers may have slight variations in their calibration, which could lead to inaccurate results. Therefore, using two different thermometers for the pure substance and the solution introduces an additional source of error and may compromise the reliability of the data.
To determine the freezing point accurately, it is best to use the same thermometer for both measurements under controlled experimental conditions. This allows for consistent and reliable results.
2. To determine the ideal melting and boiling points of a solution, you need to consider the nature of the solute and the concentration of the solution.
In this case, you are given that the solute is a nonvolatile non-electrolyte, which means it does not evaporate and does not dissociate into ions in the water. The presence of a nonvolatile solute in water will cause changes in the freezing point and boiling point of the solution.
To calculate the ideal melting and boiling points, you can use the formula ΔT = Kf * m for the freezing point depression and ΔT = Kb * m for the boiling point elevation.
Kf and Kb are the cryoscopic and ebullioscopic constants, respectively, which are specific to the solvent (water in this case).
m is the molality of the solution, which is the number of moles of solute divided by the mass of the solvent in kg.
You can find the values of Kf and Kb for water in reference tables.
By substituting the values into the equations and solving for ΔT, you can then calculate the ideal melting and boiling points of the solution by adding or subtracting these temperature changes from the normal melting and boiling points of water, which are 0°C and 100°C, respectively.
3. If the solute is K2SO4 instead of a covalent solute, the nature of the solute changes, as K2SO4 is an ionic compound. This means that it dissociates into potassium ions (K+) and sulfate ions (SO4²-) when dissolved in water.
Considering ideal conditions (no attraction between oppositely charged ions), the freezing point depression and boiling point elevation can be calculated using the same formulas as in the previous question.
However, since K2SO4 dissociates into multiple ions in solution, it is important to account for the total number of particles formed. In this case, K2SO4 dissociates into three ions (2K+ and 1SO4²-).
Therefore, the total concentration of particles in the solution is multiplied by the molality of the solution (as in the previous question) to determine the effective concentration used in the ΔT equations.
4. To calculate the weight of ethylene glycol, C2H6O2, needed to yield a solution with a specific freezing point (-15°C in this case), you can make use of the freezing point depression equation.
ΔT = Kf * m * i
In this equation:
- ΔT is the change in freezing point (the difference between the normal freezing point of the solvent and the freezing point of the solution).
- Kf is the cryoscopic constant for the solvent (water in this case).
- m is the molality of the solution, which is the number of moles of solute divided by the mass of the solvent in kg.
- i is the van 't Hoff factor, which accounts for the number of particles formed when the solute dissolves.
By rearranging the equation, you can solve for the molality (m) of the solution. Then, using the given volume (1 liter) and the density of ethylene glycol, you can calculate the mass of ethylene glycol required to achieve the desired freezing point of -15°C.