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. Well, using different thermometers to determine the freezing points may not be the coolest idea. It's like having one thermometer for your mouth and another one for your armpit to check your body temperature. You want consistent and accurate readings, so it's best to use the same thermometer for both.
2. Ah, the ideal melting and boiling points. It's like the perfect temperature to get a hot cup of tea or melt some chocolate for a delicious treat. In this case, with 0.150 mole of a nonvolatile non-electrolyte solute in 1200g of water, we need to bring out the calculators for some ideal temperature math.
3. Now, let's switch things up a bit. If the solute is potassium sulfate (K2SO4) instead of a covalent solute, we need to factor in the attractions between oppositely charged ions. How romantic, right? So, the ideal melting and boiling points of this solution will be a bit different from before.
4. A liter of water and ethylene glycol getting cozy to create a freezing solution? Sounds like a fun chemistry experiment! To get that -15°C freezing point, we need to determine the right amount of ethylene glycol to add, because let's face it, water alone can't handle the coldness. Time to play with some weights and measurements!
1. Determining the freezing point of a substance requires the use of a reliable thermometer. While it is generally advisable to use the same thermometer for both the pure substance and the solution, it is possible to use different thermometers if certain conditions are met.
To determine if it is advisable to use different thermometers, you need to consider the accuracy and precision of the two thermometers. Accuracy refers to how close the measured value is to the true value, while precision refers to the level of consistency and reproducibility of measurements.
If the two thermometers have high accuracy and precision, and if you calibrate both thermometers using a known reference temperature (such as the triple point of water), you can use one thermometer for the pure substance and another for the solution. However, it is important to keep in mind that using different thermometers may introduce additional sources of error and uncertainty to the measurements. Therefore, it is generally recommended to use the same thermometer for both measurements to minimize potential discrepancies.
2. To determine the ideal melting and boiling points of a solution, you need to apply the concept of colligative properties. Colligative properties depend solely on the number of solute particles in the solution, rather than the specific identity of the solute.
The ideal melting point of a solution can be calculated using the formula: ΔTm = (Km)(molality), where ΔTm is the change in melting point, Km is the molal freezing point depression constant, and molality is the moles of solute per kilogram of solvent.
Similarly, the ideal boiling point of a solution can be calculated using the formula: ΔTb = (Kb)(molality), where ΔTb is the change in boiling point, Kb is the molal boiling point elevation constant, and molality is the moles of solute per kilogram of solvent.
To find the molality of the solution, divide the moles of solute by the mass of the solvent (water in this case) in kilograms. In this case, the moles of solute is given (0.150 mole), and the mass of water is given (1200 g). Convert the mass of water to kilograms by dividing by 1000.
Finally, calculate the ideal melting and boiling points by substituting the values into the respective equations.
3. When dealing with ionic solutes, such as K2SO4, the ideal melting and boiling points can be determined using the same principles as in question 2, but with different constants.
For ionic compounds, the presence of ions in solution affects colligative properties differently than covalent compounds. The ideal melting point depression and boiling point elevation constants, Kf and Kb, are different for ionic solutes.
To calculate the ideal melting and boiling points of the solution with K2SO4, you would use the formulas mentioned in question 2, but substitute the appropriate Kf and Kb values for ionic solutes.
4. To determine the weight of ethylene glycol (C2H6O2) needed to produce a solution that freezes at -15°C, you can use the concept of freezing point depression.
Freezing point depression is the difference between the freezing point of the pure solvent and the freezing point of the solution. It depends on the molality of the solute and the molal freezing point depression constant (Kf) of the solvent.
First, determine the freezing point depression. It is given as -15°C, which is the difference between the freezing point of pure water (0°C) and the desired freezing point of the solution.
Next, calculate the molality of the solution. Molality is the moles of solute per kilogram of solvent. Since we have a 1-liter solution, which is equivalent to 1 kilogram, the molality will be the same as the moles of solute.
Finally, use the formula ΔTf = (Kf)(molality) to solve for the weight of ethylene glycol. Rearrange the equation to solve for molality: molality = ΔTf / Kf. Substitute the given values, and solve for the moles of solute. Since the molar mass of ethylene glycol is known, you can convert the moles to grams by multiplying by the molar mass to obtain the weight of ethylene glycol needed.