PREPARATION FOR CHEMISTRY LAB: SOLUTIONS 1 1. Define the terms solvent, solute, and solution. solvent: solute: solution: 2. In this week s lab you will be working with solutions containing a variety of solutes. Write the formula if the name is given and the name if the formula is given for each of the following: a) CuSO 4 b) NaOH c) ammonia d) magnesium sulfate heptahydrate e) (anhydrous) magnesium sulfate f) NaCl g) BaCl 2 h) sucrose i) calcium chloride j) cobalt(ii) chloride k) Ni(NO 3 ) 2 3. What is an endothermic reaction? If you were holding a beaker in which an endothermic reaction just took place, would the beaker feel hot or cold? What is an exothermic reaction? If you were holding a beaker in which an exothermic reaction just took place, would the beaker feel hot or cold? 4. Calculate the molar mass of NaCl: Calculate the molar mass of sucrose: 5. Calculate the molality of a solution produced by dissolving 81.4 grams of sucrose in 497 ml of water. You may assume the density of water is 1.00 g/ml.
SOLUTIONS 2 copyright: Department of Chemistry, University of Idaho, Moscow, ID 2002. INTRODUCTION This week in lab you will be looking at several solution-based chemical reactions. You will work with invisible inks, produce solutions that get hot or cold, observe and compare the freezing points of water, a sugar solution, and a salt solution, and make colors appear or disappear. Review Chapter 4 on naming compounds, use your Index and Glossary when needed, and read Chapter 11 in your textbook. Four activity sites are set up in the laboratory corresponding to the four parts to the lab. You don t have to do the experiments in the order in which they are listed, but you must do them all. Make careful observations as you go along and RECORD all of your observations in report style. Part 1: Invisible Ink 1. Label (pen or pencil) three filter paper circles from 1 to 3. Write your name on circle #1 with phenolphthalein indicator, on circle #2 with starch indicator solution, and on circle #3 with copper(ii) sulfate solution. The solutions are in TT. Use the swab in each solution as your pen. 2. Let the papers dry. Speed up drying by fanning the air with the paper. 3. Swab circle #1 with a solution of sodium hydroxide. Use the swab in the TT. 4. Set circle #2 over one of the glass containers containing iodine crystals which is on a hot plate in the hood. Replace the container cover after your paper has developed. 5. Set circle #3 over one of the glass containers containing concentrated NH 3 located in the hood. Replace the container cover after your paper has developed. Part 2: Energy Changes 1. Sprinkle a thin layer of magnesium sulfate heptahydrate into a 50 ml beaker (beaker #1). Add 5 ml of water to dissolve the salt. 2. Repeat, in a second beaker (beaker #2), using anhydrous magnesium sulfate. (anhydrous: without associated water)
Part 3: Freezing Points 3 1. Make a freezer containing about 200 ml of saturated common salt solution in a Styrofoam-jacketed beaker. Add some ice and sprinkle a spoonful of salt on top of the ice. Add more ice, layering salt and ice to within an inch of the top. Push the computer temperature probe into the salt solution. Remove the probe when the reading falls below -10 C. Insert a glass thermometer in the ice bath to monitor its temperature during the rest of the experiment. 2. Put 10 ml of water in a TT. Place a stirring loop inside the TT and lower the computer temperature probe through the loop. Place the TT in the ice bath. Hold the probe in place with either hand and stir the liquid occasionally with a slow up-and-down motion. Do not stir vigorously. Start the computer program. When the water freezes stop the experiment. PRINT YOUR GRAPH. 3. Dissolve between 4.05 and 4.15 g of sucrose in 10 ml of water in a TT. Using the same technique as you used for water, track the temperature change in this solution as it cools to freezing. PRINT YOUR GRAPH. 4. Dissolve between 0.62 g and 0.65 g of NaCl in 10 ml of water in a TT. Using the same technique as before, track the temperature change in this solution as it cools to freezing. PRINT YOUR GRAPH. Part 4: Color Changes (Observe the test tubes against a white background.) 1. Put 20 drops of phenolphthalein indicator in a tt. Add 20 drops of 6 M NaOH. 2. Put 20 drops of 0.1 M CaCl 2 in a tt. Add 20 drops of 6 M NaOH. 3. Put 20 drops of 0.1 M CuSO 4 in a tt. Add 20 drops of 6 M NH 3. 4. a) Begin with 10 drops of 0.05 M CuSO 4 in a TT. b) Add 10 drops H 2 O to another TT. c) Add 10 drops 0.05 M CoCl 2 to each TT. Mix. d) Add 10 drops 0.5 M Ni(NO 3 ) 2 to each TT, mixing and making observations after each drop.
DATA AND ANALYSIS SHEET: SOLUTIONS 4 Name: Date Lab Partner Explain each of your observations in a sentence or two. Identify those changes that are chemical and those that are physical in nature. Give your reasons for classifying the changes as one or the other. Part 1: Invisible Ink Circle #1: Circle #2: Circle #3: Part 2: Energy Changes 2.1: 2.2:
Name: 5 Part 3: Freezing Points; ATTACH GRAPHS 3.1: 3.2: 3.3: Exact mass of sucrose used: 3.4: Exact mass of NaCl used: Part 4: Color Changes 4.1: 4.2: 4.3: 4.4.a: 4.4.c: 4.4.d:
Name: 6 QUESTIONS ABOUT THIS LAB 1. Assuming the density of water is 1.00 g/ml: a) Calculate the molality of the sucrose solution that you prepared. b) Calculate the molality of the NaCl solution that you prepared. 2. The freezing point of an aqueous solution is lower than the freezing point of pure water. Furthermore, the amount the freezing point is lowered is related to the molality of the solution. A 0.2 m solution of a non-dissociating solute will lower the freezing point twice as much as a 0.1 m solution of a non-dissociating solute. Based only on the molality of the two solutions you prepared (Question 1 above), which solution should have the lowest freezing point? Which one of your solutions actually has the lowest freezing point?
Name: 7 3. The freezing point of an aqueous solution is lower than the freezing point of pure water. Furthermore the amount the freezing point is lowered is related to the number of solute particles (moles) in solution. Thus, when a solution is prepared from an ionic compound, the number of ions resulting from each formula unit of the compound must be taken into consideration A 0.1 m solution of CaCl 2 will have a lower freezing point than a 0.1 m solution of KBr since three moles of ions would be present in solution for each mole of CaCl 2 that dissolves and two moles of ions will be present in solution for each mole of KBr that dissolves. One mole of molecules would be present in solution for each mole of a molecular compound (not an acid or a base) that dissolves. Of the two solutions you prepared, which one has more (solute) particles in solution? Why? Based on the number of particles present in solution, which one of the two solutions should have the lower freezing point? Which one of your solutions actually has the lowest freezing point? 4. 0.010 m solutions of FeCl 3, Ca(NO 3 ) 2, and glycerol (C 3 H 8 O 3 ) are prepared. Arrange the solutions in order of lowest to highest freezing point. Explain why you arranged the solutions as you did. 5. Depending on the compound, energy can be released or absorbed when the compound dissolves in water. Based on your observations in Part 2: a) In which beaker(s), if any, did an endothermic reaction occur? How can you tell? b) In which beaker(s), if any, did an exothermic reaction occur? How can you tell?
Name: 8 6. Energy is required to break attractive interactions between solute particles and between solvent particles; energy is released when attractive interactions are formed between solute and solvent particles. Using this information, explain the results from Part 2 and Question 5 above. 7. What is the color of: CoCl 2 (aq) Ni(NO 3 ) 2 (aq) CuSO 4 (aq) Co 2+ /Cu 2+ aqueous mixture Co 2+ /Cu 2+ /Ni 2+ aqueous mixture