Working with Solutions. (and why that s not always ideal)

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1 Page 1 of 13 Working with Solutions (and why that s not always ideal) Learning Objectives: Solutions are prepared by dissolving a solute into a solvent A solute is typically a solid, but may also be a liquid Certain physical properties of solutions depend on the particular solute and solvent Solutions can be ideal or non-ideal Temperature is a measure of the energy present in a solution A calorimeter is used to observe a change in solution temperature apart from the environment Background: Several definitions are needed to adequately run this experiment: Solution - a completely uniform mixture composed of two or more substances Solute the substance in a solution that dissolves Solvent the substance in a solution that the solute dissolves in Dissolution the act of dissolving and involves the making and breaking of bonds Ideal solution a solution, that when formed, does not result in a change in energy or a change in volume Non-ideal solution a solution, that when formed, results in a change in energy or a change in volume or both To investigate ideal and non-ideal solutions, we will be using the following chemicals: Calcium chloride CaCl 2 Sodium chloride NaCl Ammonium chloride NH 4 Cl 2011 All Rights Reserved.

2 Page 2 of 13 Water H 2 O Methanol CH 4 O 2-propanol (isopropanol) C 3 H 8 O Experimental Background: Experiment 1 You re Hot then You re Cold The Objective: In this experiment there are three different salts that we will dissolve in water. Using a calorimeter (an apparatus that insulates its contents from the outside environment) and electronic thermometer, we will measure the temperature change that occurs when each salt dissolves. The three salts used in this experiment are: calcium chloride CaCl 2 sodium chloride NaCl ammonium chloride NH 4 Cl Background: Each of the salts (solutes) will easily dissolve in water (solvent) to form a solution. For an ideal solution, there will be no change in energy when the salt dissolves. For a non-ideal solution, there will be a change in energy when the salt dissolves. We measure the change in energy as a change in temperature All Rights Reserved.

3 Page 3 of 13 Equipment and Materials: Anhydrous calcium chloride Anhydrous ammonium chloride (may contain an anti-caking agent) Sodium chloride 3 top-loading balances capable of weighing to ±0.1 g Plastic weigh boats 3 scoopulas mL plastic graduated cylinders 6 stir plates with stir bars 6 calorimeters 6 electronic thermometers set to read in F to ±0.1 F Stopwatch Set- Up: Place the 3 balances in different parts of the lab space with one of the 3 chemicals at each balance. Balances should have weigh boats and scoopulas present. Reagent Prep - None Lab station Set-up: Set up 6 workstations (3 on each side of the island) with the following: 100-mL plastic graduated cylinder stir plate with stir bar calorimeter electronic thermometer Instrument Parameters - None 2011 All Rights Reserved.

4 Page 4 of 13 Experimental Procedure: The experiment works best with an instructor and 2 assistants covering the 3 balances. At minimum, 1 assistant. If only 3 partner pairs are present, each pair covers a different salt. If more than 3 pairs, add calcium chloride, then ammonium chloride, and finally sodium chloride. It is best if the pairs perform all the steps together. Water for this experiment should come from the carboy near the sink (room temperature water). Procedure: Instructor - Each team will test their salt using the following procedure: 1 Record the name of your salt on the worksheet and on a weigh boat (sodium chloride, calcium chloride, and ammonium chloride) 2 Place the labeled weigh boat on an analytical balance and weigh 5 g of your salt into the weigh boat and record the weight on the worksheet. 3 Using a 100-mL graduated cylinder, accurately measure and transfer 100 ml of water to the aluminum beaker inside of the calorimeter. Add a stir bar and place on a stir plate. Stir the water at a moderate rate. (Instructor note Make sure the stir bar is spinning and not wiggling. Confirm the thermocouple probe is not touching the stir bar when the calorimeter cover is down.) 4 Place the calorimeter lid with the electronic thermometer attached and record the temperature (in degrees Fahrenheit) of the water on the worksheet. 5 When told, remove the lid and empty the contents of the weigh boat into the water, making sure the stir bar is rotating. 6 Replace the lid. 7 Record the time and temperature on the worksheet at 30 sec, 1 min, and 3 min after adding your salt. 8 Determine the maximum temperature change (difference between the initial temperature and the highest or lowest reading you recorded) on the worksheet. Student Procedure: 1 Record the name of your salt on the worksheet and on a weigh boat (sodium chloride, calcium chloride, and ammonium chloride) 2 Weigh 5 g of your salt into the weigh boat and record the weight on the worksheet All Rights Reserved.

5 Page 5 of 13 3 Transfer 100 ml of water to the aluminum beaker inside of the calorimeter. Add a stir bar and place on a stir plate. Stir the water at a moderate rate. 4 Cover with the calorimeter lid record the temperature of the water on the worksheet. 5 Add the contents of the weigh boat into the water, making sure the stir bar is rotating. 6 Start the stopwatch and replace the lid. 7 Record the time and temperature at 30 sec, 1 min, and 3 min after adding your salt. 8 Determine the maximum temperature change (difference between the initial temperature and the highest or lowest reading you recorded) on the worksheet. Analysis You re Hot then You re Cold Worksheet Salt Weight of sample (in g) = Time (minutes and seconds) Temperature (degrees Fahrenheit) Initial Reading (no salt added) Maximum Temperature Change (+ or - ) Instructor Notes: 2011 All Rights Reserved.

6 Page 6 of 13 (Instructor note Make sure the stir bar is spinning and not wiggling. Confirm the thermocouple probe is not touching the stir bar when the calorimeter cover is down.) Discussion points: The calcium chloride should cause ~10 F increase in temperature, the ammonium chloride ~5 F decrease in temperature, and the sodium chloride ~1 F temperature change. (Sodium chloride is our best attempt at an ideal solution). Use the magnetic molecular models to describe the following: Making a solution occurs in 3 steps: 1 Breaking attractive bonds between solute molecules. This requires an input of energy from the environment that results in a drop in temperature. This is called an endothermic process. 2 Breaking attractive bonds between solvent molecules. This requires an input of energy from the environment that results in a drop in temperature. This is called an endothermic process. 3 Forming attractive bonds between solute and solvent molecules (called solvation). This provides a release of energy into the environment that produces a rise in temperature. This is called an exothermic process. The overall energy change resulting from making a solution is the sum of all these 3 processes. If the bond formation energy is greater than the bond breaking energy, the formation of the solution is exothermic (non-ideal). This causes a temperature rise. If the bond formation energy is less than the bond breaking energy, the formation of the solution is endothermic (non-ideal). This causes a temperature drop. If the bond formation energy is the same as the bond breaking energy, the formation of the solution is ideal. An example of a product that uses this chemical property is the cold pack. Explain the use of a little water plus a lot of ammonium chloride causes a 30 F or more temperature change All Rights Reserved.

7 Page 7 of 13 Experiment 2 Why It Doesn t Always Add Up Background: The Objective: In this experiment there are three different clear, colorless liquids that we will mix together to make a solution. The three liquids used in this experiment are: water H 2 O methanol CH 4 O isopropanol C 3 H 8 O Experimental Background: Background: Methanol forms a solution with both water and isopropanol. For an ideal solution, the volumes are additive ( 50 ml + 50 ml = 100 ml). For a non-ideal solution, the volumes are not additive ( 50 ml + 50 ml 100 ml). Equipment and Materials: 3 squeeze bottles each of water and isopropanol 6 squeeze bottles of methanol mL Erlenmeyer flasks each of water and isopropanol (label should be different color for each solvent) mL Erlenmeyer flasks of methanol with a label that is a different color than either water or isopropanol 3 beakers containing methanol and water transfer pipets, each labeled with a band of colored tape matching the color scheme used for labeling the Erlenmeyer flasks 3 beakers for methanol and isopropanol transfer pipets, each labeled with a band of colored tape matching the color scheme used for labeling the Erlenmeyer flasks 6 10-mL graduated cylinders with red plastic base 6 10-mL stoppered graduated cylinders 6 yellow index cards 6 red index cards 2011 All Rights Reserved.

8 Page 8 of 13 Set- Up: Reagent Prep - None Lab station Set-up Set up 6 workstations (3 on each side of the island) with the following: 10-mL graduated cylinder with red plastic base 10-mL stoppered graduated cylinder Yellow and red index cards Squeeze bottle of methanol Erlenmeyer with methanol Set up 2 workstations on one side of the island and one on the other side with the following: Squeeze bottle of isopropanol Erlenmeyer with isopropanol Beaker containing methanol and isopropanol transfer pipets Set up the remaining 3 workstations with the following: Squeeze bottle of water Erlenmeyer with water Beaker containing methanol and water transfer pipets Instrument Parameters - None Experimental Procedure: Again, the experiment works best with an instructor and 2 assistants. One assistant on each side of the island and the instructor floating from side to side. At minimum, 1 assistant. If only 3 partner pairs are present, two pairs cover water and methanol. If more than 3 pairs, add isopropanol and methanol, then water and methanol. Instructor Procedure: Each partner pair will test a solution of methanol in water or methanol in isopropanol using the following procedure: 2011 All Rights Reserved.

9 Page 9 of 13 1 Record the name of your mixture on the worksheet (methanol and water or methanol and isopropanol). We will be using graduated cylinders to hold and transfer accurate amounts of our liquids. When reading the volume of a liquid in a graduated cylinder, you must read to the mark of the liquid in the middle of the cylinder, not at the wall. Some liquids will rise up the walls of the cylinder (example A) and some will move down (example B). The reading in either case is taken from the middle of the cylinder (the dotted line in both examples). (Instructor note Explain that these solutions will rise up the walls like example A) 2 Fill the 25 ml cylinder with the squeeze bottle to as near the 12 ml mark as possible with methanol. Fill to just below the mark and then add the final amount using the plastic transfer pipet. 3 Record the volume to the nearest 0.2 ml on the worksheet. (Instructor note Explain to the students that the squeeze bottle works best in an upright position. Also, the yellow index card acts as a good background against the blue grad cylinder divisions. Explain that there is no 12 ml mark, only the large line between 11 ml and 13 ml. Also each small line is 0.2 ml.) The students will need to keep their eyes level with the grad cylinder markings to get accurate readings. Make sure you personally review each final measured volume to make sure it is correct 4 Fill the 25 ml cylinder with the glass stopper to as near the 12 ml mark as possible with water or isopropanol (based on which mixture you are preparing). Again, fill to just below the mark using the squeeze bottle and then add the final amount using the plastic transfer pipet. (Instructor note Make sure the correct transfer pipet is used for each solvent. That s why they are color coded!) 5 Record the volume to the nearest 0.2 ml on the worksheet. 6 Before mixing the liquids, calculate the expected total volume (the two volumes added together, to 1 decimal place) and record on the worksheet All Rights Reserved.

10 Page 10 of 13 7 Slowly and completely transfer the 12 ml methanol into the graduated cylinder containing either water or isopropanol. 8 Place the stopper on the graduated cylinder, grasp the cylinder in your hand with your thumb over the top of the stopper, and invert the cylinder 5 times. 9 Set the graduated cylinder on the lab bench and determine the total volume present. (Instructor note Have the students view the solution and note if something happened, like a color change or BUBBLES. If they see bubbles, tell them to remember that.) 10 Record this volume on the worksheet. Student Procedure 1 Record the name of your mixture on the worksheet (methanol in water or methanol in isopropanol). We will be using graduated cylinders to hold and transfer accurate amounts of our liquids. When reading the volume of a liquid in a graduated cylinder, you must read to the mark of the liquid in the middle of the cylinder, not at the wall. 2 Fill the 25 ml cylinder with the red bottom to as near the 12 ml mark as possible with methanol. Add the final amount using the plastic transfer pipet. 3 Record the volume to the nearest 0.2 ml on the worksheet. 4 Fill the 25 ml cylinder with the glass stopper to as near the 12 ml mark as possible with your solvent (water or isopropanol). Add the final amount using the plastic transfer pipet. 5 Record the volume to the nearest 0.2 ml on the worksheet. 6 Before mixing the liquids, calculate what the expected total volume and record on the worksheet All Rights Reserved.

11 Page 11 of 13 7 Transfer the 12 ml methanol into the graduated cylinder containing either water or isopropanol. 8 Place the stopper on the graduated cylinder, grasp the cylinder in your hand with your thumb over the top of the stopper, and invert the cylinder 5 times. 9 Set the graduated cylinder on the lab bench, remove the stopper, and determine the volume. 10 Record this volume on the worksheet. Analysis Why It Doesn t Always Add Up Mixture Liquid Volume Used (to the nearest 0.2 ml) Mixture Expected Total Volume (to the nearest 0.2 ml) Actual Total Volume (to the nearest 0.2 ml) Difference (to the nearest 0.2 ml) 2011 All Rights Reserved.

12 Page 12 of 13 Instructor Notes: Explain that these solutions will rise up the walls like example A. Explain to the students that the squeeze bottle works best in an upright position. Also, the yellow index card acts as a good background against the blue grad cylinder divisions. Explain that there is no 12 ml mark, only the large line between 11 ml and 13 ml. Also each small line is 0.2 ml. Make sure the correct transfer pipet is used for each solvent. That s why they are color coded! Have the students view the solution and note if something happened, like a color change or BUBBLES. If they see bubbles, tell them to remember that. Discussion points: The water/methanol mixture will shrink ~1.0 ml. The isopropanol/methanol should be additive, or only shrink by 0.2 ml. You can use the observation to say that the isopropanol/methanol pairs knew what they were doing, but the water/methanol pairs must of screwed something up or spilled some. They will argue this point and you can use it to explain to them what happened using the molecular models. Show how 3 water molecules can fit closely together. Position 3 methanol molecules such that the OH functionalities get somewhat close without the CH3 group not touching. Ask where a water molecule could possibly fit and show them it fits between the 3 OH groups without touching anything else. This is why methanol and water don t mix additively. Then ask what did the water molecule displace in the solvent of 3 methanol molecules. The answer is air, which is observed leaving the solutions as bubbles. Then show how an isopropanol molecule cannot fit where the water molecule fit, so methanol and isopropanol is additive, air is not displaced, so no bubbles. So isopropanol/methanol is ideal, and water/methanol is non-ideal. If there is time, show that water/methanol is also non-ideal because of a large temperature change when mixed. Using the electronic thermometer, take the temperature of the methanol flask, then the isopropanol flask, and finally the water flask, making sure to dry the thermocouple between measurements. Then measure the isopropanol/methanol mixture and the temperature might be 2 to 3 F higher. (Can be explained by handling the graduate with warm hands.) Then the water/methanol, which will be ~10 F higher. The increase in water/methanol can be explained because the bond breaking of methanol/air requires little energy but the making of a methanol/water bond releases a lot of energy All Rights Reserved.

13 Page 13 of 13 NOTE: When cleaning up between teams, make sure not to mix water and isopropanol graduates. It seems that isopropanol/methanol mixtures shrink if the graduate was previously filled with water. Wrap-Up: Applying Your Knowledge Based on your observations and the results of the other partners, what might be a practical use for making a calcium chloride solution? What might be a practical use for making an ammonium chloride solution? There is a direct relationship between the amount of salt dissolved and the change in temperature. Based on your results, if you dissolved 15 g of salt instead of 5 g in the same amount of water (100 ml), what would the change in temperature be after 3 min? Would you expect a different maximum temperature change if the initial temperature of the 100 ml of water was 50 F instead of room temperature? Why? 1 If there is time, you will record the temperature of each liquid and your mixture? Temperature of methanol Temperature of water or isopropanol Temperature of mixture 2 The addition of methanol to water causes a loss of volume and an increase in temperature. If you let the methanol / water mixture sit undisturbed in the graduated cylinder for 24 hours, would you expect the temperature to still be greater than room temperature? Why? 3 Would you expect the volume to still be less than expected? Why? 4 If you added isopropanol to water, would you expect a loss in volume? Why? 5 Would you expect an increase in temperature? Why? 2011 All Rights Reserved.