AP+Chem+Lab+8

Solutions are created by adding a solute to a solvent. The temperature at which a solution freezes is always lower than the freezing temperature of the pure solvent; the freezing temperature is lowered in proportion to the number of moles of solute added. This property, known as freezing-point depression, is a colligative property; that is, it depends on the ratio of solute and solvent particles, not on the nature of the substance itself. The equation that shows this relationship is ΔT = Kfm where ΔT is the freezing point depression, Kf is the freezing point depression constant for a particular solvent, and m is the molality of the solution. The magnitude of this depression depends on the concentration of solute particles in the solution, not the identity of the solute. For compounds that do not dissociate, the freezing point depression of the solution depends on the number of moles of that compound per kilogram of solvent. For compounds that do dissociate, such as soluble ionic compounds, freezing point depression depends on the number of dissociated ions per kilogram of solvent. In this experiment, you will determine the molar mass of an unknown solute. Ask your teacher if your unknown compound dissociates, and, if so, how many ions are contained in a single formula unit. The freezing point depression is directly related to the number of moles of particles per kilogram of water. By measuring the mass of water used and the freezing temperature of the solution, one can calculate the number of moles of the compound that were dissolved in the solution using the formula, ΔTf = imKf(water) The freezing constant, Kf for water is 1.86 K  kg/mol, and the van’t Hoff factor, //i// is the number of ions in one formula unit of the compound. The total mass of solute added must be measured before it is dissolved in the water that is to be frozen. The molar mass is found by dividing the mass of the solute, in grams, by the number of moles of solute. To determine the freezing points of the pure solvent and solution, you must plot cooling curves for each. Examples of such plots are shown below.
 * AP Lab #8 Determination of Molar Mass by Freezing Point Depression **
 * __Introduction:__**
 * __Background:__**

Some liquids supercool before they solidify. During supercooling the solvent or solution remains in the liquid phase after the temperature has dropped below the freezing temperature. At the base of the supercooling curve, solid crystals form. This causes the temperature of the system to increase dramatically, as heat is released by the solvent particles during the process of solidification. The method for determining the freezing point for the pure solvent differs slightly from that of the solution. For a pure solvent, the freezing temperature is found by following the plotted horizontal line to the right of the supercooling curve directly across the graph to the y-axis. For a solution, you need to draw two lines. The first line follows the slope of the cooling liquid before it starts to level off. The second line follows the slope of the majority of the plotted points between the supercooling curve and the point at which the entire system solidifies. The y-coordinate for the point at which these two lines intersect is the freezing temperature for the solution. You may have noticed that the solid-liquid line in the cooling curve for the solution is not horizontal. This is because the solid crystals that form are composed of pure solvent. The decreasing volume of liquid solvent causes the concentration of solute particles to increase in the remaining liquid solution. As you already know, freezing temperature decreases as the concentration of solute particles increases.
 * __Purpose:__**
 * To determine the molar mass of an unknown soluble compound using the freezing point depression method.
 * __Safety:__**
 * Wear your lab coat and safety goggles at all times during this lab.
 * You will be working with an unknown substance and it should be treated with respect.
 * __Materials:__**
 * Balance
 * Thermometer
 * Large test tube
 * Two hole stopper
 * Thin piece of wire to be used as a stir stick
 * Test tube stopper without holes
 * Ice bath
 * Ring stand
 * Buret clamp
 * Stir stick
 * 100 mL graduated cylinder
 * funnel
 * paper towel
 * Ice
 * Rock salt
 * Distilled water
 * Unknown solute

__Part 1: Determining the freezing point of pure water__ The thermometer that you have will likely be precise, but it may not be accurate. Use the same thermometer for both parts of the experiment. The value for the freezing point depression, ΔTf, will be correct so long as you use the same thermometer for the entire experiment. __Part 2: Measuring the mass of the solvent and unknown solute__ __Part 3: Determining the freezing point of the solution__
 * __Procedure:__**
 * 1) Clean the test tube and add approximately 20 mL of distilled water.
 * 2) Insert the stopper with thermometer and wire stir stick. Make sure that the thermometer is immersed in water but is not touching the sides or bottom of the test tube.
 * 3) Prepare the ice bath. Put a 3.5 cm layer of crushed ice on the base of the 1.0 L beaker. Sprinkle a thin layer of rock salt on top of the ice. Add another 3.5 cm layer of crushed ice followed by another thin layer of rock salt. Repeat this pattern of layering until the beaker is nearly full and then stir the contents of the beaker with the stir stick.
 * 4) Immerse the test tube assembly in the ice bath. You may want to fasten the test tube to the ring stand with the buret clamp. If not, you can hold the assembly in your hand ensuring that the mouth of the test tube does not fall below the surface of the ice bath.
 * 5) Stir the distilled water inside the test tube continuously, and stir the water in the ice bath once every minute. When the thermometer reaches 4.0°C, start recording the temperature to at least one decimal place every 20 seconds until the temperature remains constant for eight consecutive readings.
 * 1) Measure and record the mass of approximately 1.1 grams of the unknown solute to as many significant digits as you can.
 * 2) Measure exactly 20.0 mL of distilled water in your graduated cylinder, and transfer it into a clean and dry test tube.
 * 1) Add the unknown solute to the test tube using a dry funnel. Seal the test tube with the solid test tube stopper and move it carefully until all of the solute dissolves.
 * 2) Remove the solid test tube stopper and insert the stopper with thermometer and stir stick. Make sure that the thermometer is immersed in the solution but is not touching the sides or bottom of the test tube.
 * 3) Replenish the ice bath. Pour off most of the salt water. Sprinkle a thin layer of rock salt over the remaining ice. Cover that with 3.5 cm layer of crushed ice. Sprinkle a thin layer of rock salt on top of the ice. Add another 3.5 cm layer of crushed ice followed by another thin layer of rock salt. Repeat this patter of layering until the beaker is nearly full and then stir the contents of the beaker with the stir stick.
 * 4) Immerse the test tube assembly in the ice bath. You may want to fasten the test tube to the ring stand with the clamp. If not, you can hold the assembly in your hand ensuring that the mouth of the test tube does not fall below the surface of the ice bath.
 * 5) Stir the solution inside the test tube continuously, and stir the water in the ice bath once every minute. When the thermometer reaches 4.0°C, start recording the temperature to at least one decimal place every 20 seconds until the temperature remains constant or decreases only slightly for eight consecutive readings.
 * __Pre-Lab Questions:__**
 * 1) Why do you need to use the same thermometer to find the freezing point of water and the freezing point of the solution?
 * 2) Why is it important that the thermometer does not touch the sides or bottom of the test tube while you are recording temperatures?
 * 3) Why is it necessary to add salt to the ice bath in this experiment?


 * Number of ions in one formula unit of the unknown solute ||  ||
 * Mass of unknown solute (grams) ||  ||
 * Volume of solvent (water mL) ||  ||


 * Water ||||  Solution  ||
 * Time ||  Temperature  ||  Time  ||  Temperature  ||
 * __Graphs and Calculations:__**
 * 1) Plot cooling curves for pure water and the solution.
 * 2) Find the freezing point depression, ΔTf
 * 3) Calculate the number of moles of the unknown compound that were added to the distilled water.
 * 4) Calculate the molar mass of the unknown compound.
 * 5) Calculate the percentage error if the actual molar mass of the compound is 58.5 g/mol.
 * __Post-Lab Questions:__**
 * 1) What are the measurements that must be taken to find the molar mass of an unknown solid using the freezing point depression method?
 * 2) When a student finished mixing the solution in the test tube (step 8), the fire alarm went off and the building was evacuated for a period of time. After returning to class, the student finished the experiment. How would the student’s calculated value for the number of moles of unknown solid be affected if several milliliters of water evaporated from the test tube while the student was out of the room? Justify your answer.
 * 3) The van’t Hoff factor is often a little less than the number of ions in one formula unit of an ionic compound due to ion pairing and ion clusters. Explain why this is.
 * 4) How would ion pairing and ion clusters in the solution affect the calculated value of the molar mass in this experiment?
 * 5) Suppose that the thermometer you used in this experiment consistently read 0.9°C below the actual temperature. What affect would this have on the experimentally determined value for the freezing point depression, ΔTf? Explain.
 * 6) Suppose that some of the unknown solute did not dissolve in the 20 mL of water due to saturation. How would this affect the calculated value of the molar mass? Explain.
 * 1) What are the measurements that must be taken to find the molar mass of an unknown solid using the freezing point depression method?
 * 2) When a student finished mixing the solution in the test tube (step 8), the fire alarm went off and the building was evacuated for a period of time. After returning to class, the student finished the experiment. How would the student’s calculated value for the number of moles of unknown solid be affected if several milliliters of water evaporated from the test tube while the student was out of the room? Justify your answer.
 * 3) The van’t Hoff factor is often a little less than the number of ions in one formula unit of an ionic compound due to ion pairing and ion clusters. Explain why this is.
 * 4) How would ion pairing and ion clusters in the solution affect the calculated value of the molar mass in this experiment?
 * 5) Suppose that the thermometer you used in this experiment consistently read 0.9°C below the actual temperature. What affect would this have on the experimentally determined value for the freezing point depression, ΔTf? Explain.
 * 6) Suppose that some of the unknown solute did not dissolve in the 20 mL of water due to saturation. How would this affect the calculated value of the molar mass? Explain.