California State University Dominguez Hills Semester, 200X

Transcription

1 California State University Dominguez Hills Semester, 200X

2 Chemistry 103L: Chemistry for the Citizens List of Experiments Orientation: Check-In and Safety Film Experiment #1: The Bunsen Burner Experiment #2: Measurements Experiment #3: Density and the Separation of a Mixture Experiment #4: Elements and Compounds Experiment #5: An Introduction to Chemical Reactions Experiment #6: Paper Chromatography Experiment #7: The Iodine Clock Reaction Experiment #8: Isolation of Caffeine from Mountain Dew Experiment #9: A Chemistry Investigation Experiment #10: Esters and the Preparation of Soap Experiment #11: Cabbage Juice and the ph of Common Chemicals Experiment #12: Percentage of Water in Popcorn Experiment #13: Vitamin C Content of Fruit Juice/Check-Out Experiment optional #1: Carbon Dioxide and Oxygen Experiment optional #2: Synthesis of Aspirin Experiment optional #3: Investigating Polymers: Slime and Silly Putty Experiment optional #4: Water Hardness i

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4 Experiment 1: The Bunsen Burner In this brief exercise, we will study the operation of one of the most useful devices in the chemistry laboratory: the Bunsen Burner. The substance burned in the Bunsen Burner is methane (also called marsh gas or natural gas, chemical formula CH 4 ). The burner mixes methane and oxygen (from the air) before these gases reach the flame. The flame color indicates the completeness of the combustion. A blue flame results from the complete combustion of methane, forming water, carbon dioxide, and giving off heat. A yellow color indicates incomplete combustion; glassware exposed to such a flame may become coated with black carbon. PROCEDURE 1. Determine where the gas inlet at your station is and locate the gas control on your burner. Determine where the air control is for the burner. Turn on the gas at the main valve at your station, and light the burner with a match. If the flame keeps blowing itself out, then turn down the gas flow valve. 2. Close off the air vents by screwing the barrel down onto its base. A blue color in your flame indicates that you have not completely shut off the air. You should see a drastic change in the flame. Note the color of your flame. 3. Using crucible tongs, hold an evaporating dish half-filled with DI water over the very top of the flame for 20 to 30 seconds. Observe any deposit formed on the underside of the evaporating dish. Clean the evaporating dish with a paper towel. 4. Hold a wire gauze horizontally in the flame, raising and lowering it every few seconds. If your gauze has a ceramic center, heat the bare metal at the corners. Does the wire glow red hot at any particular location in the flame? 5. Open the air vent slowly. If your flame goes out, then cut back on the gas flow and air flow and re-ignite the burner. You should have a flame with two distinct coneshaped sections. 6. Repeat the heating test with an evaporating dish (3) and a wire gauze (4). Page 1-1

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6 The Bunsen Burner Report Sheet Name: 1. Observations of Bunsen Burner: Where is the gas control located on the burner How is the air controlled? 2. Air vent closed Appearance of flame Description of deposit Did the wire gauze glow red anywhere in flame? 3. Air vent open Appearance of flame Description of deposit, if any Did the wire gauze glow red hot anywhere in flame? Page 1-3

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8 Preparation for The Bunsen Burner : Please check to see that all Burners are functioning and that the rubber tubing is in good condition. Otherwise, no special preparation is required. Page 1-5

9 Experiment 2: Measurements As part of their daily work, scientists must carry out common laboratory procedures, take measurements, and report their results accurately and clearly. The system of measurement used in modern science is the metric system. The metric system is a decimal-based system, meaning that measurements of each type are related by factors of 10. The metric system has one standard unit for each type of measurement: the standard unit for length measurement is the meter (m), the unit for mass measurement is the gram (g), and the unit for volume is the liter (L). When taking measurements of quantities that are significantly larger or smaller than the standard unit, prefixes are attached in front of the standard unit. The most common prefixes are list below. PROCEDURE Prefix Symbol Meaning kilo k 1000 deci d 1/10 (0.1) centi c 1/100 (0.01) milli m 1/1000 (0.001) A. Measuring Length a. Observe the marked lines on a meter stick. Identify the lines that represent centimeters and millimeters. Record your observations. Observations (include a drawing, and label the length represented by each line): b. Estimate the length and width of this paper in centimeters, and record your estimate here: Estimated length: cm width: cm. Now, use the meter stick or a ruler to measure the length and width of this paper in cm. Estimate the last digit if it fells between the lines. Actual length: cm width: cm c. Using a piece of string and the meter stick, measure the length of your shoes in centimeters. Length of shoes: cm Page 2-1

10 B. Measuring Volume a. Remove a test tube from your drawer (do not use the wider test tubes). Fill the tube with DI water to the rim (almost overflowing). Hold your large (100 ml) graduated cylinder next to the test tube. Estimate the volume of water in this test tube in milliliters (ml) using the graduated cylinder for comparison. Estimated volume ml Carefully pour the water into the large graduated cylinder (use your dropper to adjust the volume to the last line). Record the volume. Actual volume ml (using 100 ml graduated cylinder) b. Fill the test tube with DI water to the rim again. Pour the water into a medium (50 ml) sized graduated cylinder. Record the volume reading. Actual volume ml (using 50 ml graduated cylinder) c. Repeat this process using a small (10 ml) graduated cylinder. You must do this in several steps, by first transferring no more than 10 ml of water to the graduated cylinder at a time, and then pouring the water into the sink. When you are done, obtain the volume by adding the total volume measured. Tip:pour about 9 ml of water into the cylinder, then use your dropper to transfer enough water to exactly match the 10 ml line. This will give you the most accurate results. Actual volume ml (using 10 ml graduated cylinder) (Calculations) d. Obtain a plastic gallon jug and a large 1L graduated cylinder from the cart. Fill the graduated cylinder with tap water to the 1 Liter mark. Estimate how many liters are in one gallon, and write down this value. Estimated volume L (your estimate does not need to be a whole number!) e. Carefully pour the water from the graduated cylinder into the gallon jug. Continue to refill the graduated cylinder with water and pour water into the gallon jug until a total of one gallon of water has been added. From this, determine how many liters are in a gallon. Have the instructor initial your results. Actual volume L Instructor Initial: Page 2-2

11 C. Measuring Mass a. After your instructor shows you how to use a laboratory balance, determine the mass of a 50- ml beaker, a stopper (on cart), and a metal disk (also on the cart). Copy the number or letter on the disc on your report sheet (if there is one) You must always record all of the digits from the balance reading on your report sheet. Do not round off or drop any zeros! Mass of beaker: g Mass of stopper: g Mass of disc# : g b. Pour just enough rock salt into the beaker you weighed in the previous step to cover the bottom of the beaker. Then, estimate the mass and write down your estimation. Estimated mass of salt: g Use the balance to determine the mass of your salt, and record the value. Mass of salt + beaker: g Calculations Mass of salt: g Instructor s Initial: c. Set the salt aside (perhaps on the watchglass or a piece of paper towel) for use in part D. d. Pour about 1 g of sugar into the same beaker. Estimate and measure the mass of the sugar, just as you did in part b. Estimated mass of sugar: g Mass of sugar + beaker: g Mass of sugar: g Page 2-3

12 D. Measuring Temperature a. Determine the temperature of the room by simply reading the value on your thermometer. Note that if the temperature reading appears to be on the line, you should write.0 at the end of the measurement; if it is between the lines you should write.5 Temperature of room: o C b. Fill your 250-mL beaker about half way to the top with deionized (DI) water. Then, record the temperature of the water. Temperature of water: o C c. Put a handful of ice into the beaker, and stir it for about a minute with your glass stirring rod (never stir with your thermometer!) Then, record the temperature of the slush you have made. Temperature of ice-water: o C d. Pour the salt from Section C into the ice water, and stir it for about two minutes. Record the temperature of this mixture. Temperature of ice water: with salt o C Copy all values from these sheets over to the Report Sheet. Do not forget to copy all units! You will only turn in the Report Sheet when you are done. Save the remainder to help you review for the lab final. Page 2-4

13 Report Sheet Name: MEASUREMENTS 1. Measuring Length a. What units are represented by the numbers marked on the meterstick? There are centimeters in 1 meter. There are millimeters in 1 cm. There are millimeters in 1 m. b. Estimated value Measurement Length of the paper Width of the paper c. Length of your shoe 2. Measuring Volume Estimated volume of the test tube Size of the cylinder (in ml) Volume of water from the test tube Largest cylinder Medium cylinder Smallest cylinder Estimated volume of one gallon of water Actual volume of one gallon of water C. Measuring Mass a. mass of 50-mL beaker b. mass of the stopper c. mass of the disk code # of the disk Page 2-5

14 d. Estimated value Measurement 5 g of salt 1 g of sugar D. Measuring Temperature a. Temperature of the air in the room b. Temperature of the water c. Temperature of the ice-water d. Temperature of the salt/ice-water mixture Be sure that you have included the correct units for all answers, where appropriate! Questions. 1. Balances can be used to weigh objects of many different masses and sizes. Why would a balance at a truck-stop which is used to weigh 18-wheelers and their cargo be inappropriate for determining your own weight? Be as specific as possible. Page 2-6

15 2. Devise a way to measure the approximate thickness of a single sheet of paper in millimeters, using a ruler and paper (and maybe a calculator) as your only tools. (Only suggest a procedure; you do not need to carry it out.) Then, list the steps you would use clearly and in order. Page 2-7

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17 Materials for this experiment: Meter sticks String Size 4 stoppers Unknown metal disks 5 one-gallon plastic bottle salt 5 g/ student sugar 10 g per student 5 one-liter graduated cylinder ice Page 2-9

18 Experiment 3: Density & Separation of a Mixture To determine the density of a substance, you need to measure both its mass and its volume. From the mass and volume, the density is calculated. If the mass is measured in grams and the volume is measured in ml, the density will have a unit of g/ml. PROCEDURE A. Density of a liquid a. Place about 20 ml of water in a graduated cylinder. Determine the volume exactly from the markings on the cylinder. Record the volume in ml to one decimal place. Volume of water: ml b. The mass of the liquid is determined by weighing by difference. Alternatively, you may use the zeroing feature of the balance instead. First, determine the mass of a 50-mL dry beaker. Pour the water from the cylinder into the beaker, and reweigh. Record the mass, writing down all the digits in the display. Calculate the mass of the water. Finally, determine the density of water. Mass of water: g Density of water (use formula above): Show calculations! Instructor s Initial c. Repeat the same procedure for another liquid assigned by your instructor. Put your data here. Label the information, and use the appropriate units. Page 3-1

19 B. Density of a solid a. Obtain a solid metal cylinder and record the letter on it. Determine the mass and record. Letter on metal cylinder: Mass of cylinder: b. Take out your 50 ml graduated cylinder, and fill it with water until the cylinder is about half full. Then, read the exact volume of water and record it. Volume of water in cylinder: ml (How many decimal places?) c. Slowly lower the metal object into the water. Measure and record the final volume when the object is completely submerged. Volume of water + metal cylinder: ml d. Calculate the volume of the metal. Then, determine the density of the solid. Show your calculations and data here. Density of the solid metal cylinder: (Units?) C. The Separation of a Salt-Water mixture a. Determine the mass of a piece of weighing paper on the balance. Then, weigh out about 1.5 grams of salt (NaCl) onto the paper. Determine the exact mass of salt that you used. Be sure to clean the balance pan when you are done! Mass of salt (to three decimal places): (you supply the units from now on) b. Pour the salt into a 150 ml beaker. Then, pour about 5 ml of DI water into the beaker and stir to dissolve all of the salt. If the salt does not dissolve after a few minutes of stirring, add more water, bit by bit, and stir until the salt does dissolve. You want to use as little water as possible, c. Determine the mass of your empty evaporating dish. Then, pour your salt water mixture into it. Mass of evaporating dish: (Did you write down all digits from the balance?) Page 3-2

20 d. Set-up a steam bath like that in figure 3-1. Set the evaporating dish on top of the steam bath, making sure that the evaporating dish is completely supported by the beaker. Allow all the water in the solution to evaporate. Some tips: First, choose the size of beaker which the evaporating dish will fit best into, just like you see in the diagram. The dish should not fall in to the beaker, nor should it be too large to rest on its rim. The beaker should be filled about halfway to the top with tap water (why not DI water?). This should be enough. Do not add cold water to the beaker once it is hot or it may shatter! Figure 3-1: Steam Bath Assembly e. Remove the evaporating dish from the beaker carefully using your crucible tongs. Using the beaker tongs provided by the instructor, remove the beaker holding the boiling water. Dry the evaporating dish completely by placing it on the wire gauze and heating it directly with the flame for about 2 minutes. f. Turn off the flame and allow the dish to cool for about 10 minutes. g. Determine the mass of the evaporating dish with the salt residue in it. Calculate the percent of the salt that you were able to recover. Mass of Evaporating Dish + Salt Residue: Calculations Mass of Salt Resdidue: Percent of Salt Recovered: % Page 3-3

21 h. Wash the salt residue down the drain. D. Determining the Percent of Salt in an Unknown Solution a. Rinse the evaporating dish you used in the last section with DI water, and then dry it with a paper towel. Try your best to remove all visible water. b. Select an unknown sample (A, B, or C) from the cart. Your instructor may assign you one. Immediately write down the letter of the unknown you are using on the report sheet. Unknown: c. Pour about 5-7 ml of one of the three unknown salt solutions into the evaporating dish. Carefully place this on the balance and record the mass. Mass of evaporating dish + salt solution: Mass of salt solution: (You already determined the mass of the dish) d. Set the evaporating dish on top of the steam bath you assembled in the previous section. Allow all the water in the solution to evaporate. e. Dry the evaporating dish completely by placing it on a wire gauze and heating it directly with the flame for about 2 minutes. Turn off the flame and allow the dish to cool for about 10 minutes. f. Determine the mass of the evaporating dish with the salt residue in it. Calculate the percent of the salt in the unknown solution. Mass of evaporating dish + salt residue: Mass of salt residue: (Again, you should know the mass of the dish) g. Wash the salt residue down the drain. Page 3-4

22 Report Sheet Name DENSITY & SEPARATION OF A MIXTURE Always give the appropriate units! A. Density of a liquid Name of the liquid Liquid 1 Liquid 2 Volume of the liquid Mass of beaker Mass of beaker + liquid Mass of the liquid Density of the liquid Show calculations for density: B. Density of a Solid Mass of the solid Volume of the solid by displacement: Initial water level Final water level with solid Volume of the solid Density of the solid Show calculations: Page 3-5

23 C. Separation of a Mixture Mass of weighing paper g Mass of weighing paper and salt g Mass of salt (show calculation) g Mass of evaporating dish g Mass of evaporating dish and salt residue g Mass of salt residue (show calculation) g Percent of salt recovered (show calculation) % Was your salt-water mixture a homogenous mixture or a heterogeneous one? Judging by your percent recovery, how successful were you at separating all of the salt from the water? If your value was either greater or lower than 100%, explain how this could have occurred. Page 3-6

24 D. Percent of Salt in an Unknown Solution Unknown Mass of evaporating dish (from Part C) g Mass of evaporating dish and solution g Mass of solution (show calculation) g Mass of evaporating dish and salt residue obtained after heating Mass of salt residue (show calculation) g g Percent salt in solution (show calculation) % Questions Suppose you have 10. ml of water in one beaker #1 and 25. ml of the same water in beaker #2. Is the density of the water in beaker #2 greater than the density of the water in beaker #1? Explain. Page 3-7

25 Suppose that you have 25. ml of distilled water and 25. ml of sea water. Should one sample have a greater mass than the other, and, if so, which one? Explain Page 3-8

26 Materials for this lab: Liquids for density tests: Toluene, Methylene Chloride, Ethanol (95%), Methanol (two 250 ml bottles of each, labeled with name) Coke, diet coke, pepsi, diet pepsi : one can each Unknown metal cylinders A, B, C, D Sodium chloride (3 small containers, about 10+ grams each) Unknown salt solutions: Prepare 500 g of each according to directions below. Label as Unknown A, Unknown B, and Unknown C Unknown A: 10 g NaCl ml water Unknown B: 20 g NaCl ml water Unknown C: 15 g NaCl ml water Page 3-9

27 Experiment 4: Elements and Compounds There are over 100 known elements, each unique in its own way. These elements can theoretically combine together to create an infinite number of compounds. These compounds are almost always significantly different from the elements from which they were formed. Consider such toxic elements as sodium and chlorine, which combine to form sodium chloride, a compound essential for life. In this experiment you will have the opportunity to take a tour of some of the more common elements, examine their properties, and learn how to identify their presence in a solution. PROCEDURE A. Instructor Demonstration: The Electrolysis of Water a. Set-up an apparatus designed to pass electricity through water. b. Apply a small voltage across the water. Observe the results. c. Increase the voltage. Observe any changes. d. After the apparatus has run for some time, compare the amount of gas in each tube. B. Instructor Demonstration: Sodium Chloride vs. Sodium Metal a. Observe sodium chloride and sodium metal. How are they similar? How are they different? b. To a beaker of water, add a few drops of phenolthalein indicator. Observe the color. Now, add a small sample of sodium chloride. Observe any significant changes. c. Add a few drops of water to a different beaker of water. Now, carefully add a small piece of sodium metal to the water. Observe any changes. C. Sublimation of Iodine a. Set up a ring stand and wire gauze. Put a Bunsen burner under the gauze, but do not light it yet. b. Add a small amount (about 1 gram) of solid iodine (I 2 ) to a 250 ml or 400 ml beaker. This beaker must be dry or it may be ruined by this experiment! c. Take your evaporating dish and fill it with ice. Then, cover the top of the beaker with the evaporating dish. The bottom of the dish must totally cover the opening of the beaker. d. Light your Bunsen burner, and gently heat the bottom of the beaker. Do not continue heating after the iodine has sublimed. e. Record your observations on the report sheet. D. Reaction of Potassium Iodide with Lead (II) Nitrate a. Measure out between 1 and 2 ml of potassium iodide (KI) solution into one test tube; measure out between 1 and 2 ml of lead (II) nitrate solution Pb(NO 3 ) 2 into another test tube. b. Pour the contents of one test tube into the other. c. Record your observations on the report sheet. E. Flame Tests a. Fill each of seven wells in a spot plate with one of the solutions provided. Each solution contains a different ion: sodium(na), potassium(k), lithium(li), copper (Cu), calcium(ca), barium(ba), and strontium(sr). Be sure to keep track of which well contains which ion. b. Take a platinum wire, and dip into distilled water to clean it. Then, dip into one of the solutions in the spot plate. Place the wet wire in the hottest part of a Bunsen burner flame. Do not put the glass portion of the wire holder in the flame. Observe and record your observations. c. Clean the wire by rinsing it with distilled water, then dipping it in 1 M hydrochloric acid (HCl), and finally by rinsing with distilled water again. Caution: hydrochloric acid can cause severe burns! d. Repeat this procedure for the remaining solutions. F. Elements a. Samples of several elements have been placed around the room. Provide a visual inspection of each one on the report sheet. Page 4-1

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29 Report Sheet Name Elements and Compounds A. Electrolysis of Water a. What did you observe when voltage was applied to the water? What changes occurred when the voltage was increased? b. Compare the difference in the volume of gas in each tube. One of the tubes contains hydrogen gas, the other oxygen. Which tube contains which gas? [Hint: think about the chemical formula of water] B. Sodium Chloride vs. Sodium Metal a. Description of sodium chloride b. Description of sodium metal c. What happened when sodium chloride was added to water? d. What happened when sodium metal was added to water? Page 4-3

30 e. Based on these observations, should one assume that sodium and its compounds (like sodium chloride) will always have similar chemical and physical properties? C. Sublimation of Iodine a. Describe what you observed in the beaker after the iodine was heated. b. Did you observe a physical change or a chemical change? Explain how you came to this conclusion. D. Reaction of Potassium Iodide with Lead (II) Nitrate a. What did you observe when the solutions were mixed? b. Was this a chemical change or a physical change? c. Is the color of the product a physical property or a chemical property? Page 4-4

31 E. Flame Tests a. Describe the color of light given off by each of the six solutions. sodium: potassium: strontium: lithium: copper (II): calcium: barium: b. Based on your observations, can we assume that it is always possible to identify an element with a flame test, using only our eyes to assist us? Explain. F. Elements a. Briefly describe the physical properties of each of the elements below. copper: mercury: sulfur: phosphorus: oxygen: nickel: carbon: zinc: Page 4-5

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33 Materials for this lab: For the instructor: Apparatus for electrolysis of water (with acidified water) Three 800-mL beakers Small samples (< 10 g) of sodium chloride and sodium metal(in oil) Phenolthalein solution in a small dropper bottle For the class: 3 small bottles of each of the following solids: iodine potassium iodide lead (II) nitrate dropper bottles containing 0.2 M solutions of each of the following (for flame tests) NaCl, LiCl, KCl, CuCl 2, BaCl 2, CaCl 2, SrCl 2 platinum wire 1 M HCl dipping solution to clean wires ice Samples of these elements for viewing: copper, mercury, sulfur, phosphorus (both allotropes), oxygen, nickel, carbon, zinc Page 4-7

34 Experiment 5: Introduction to Chemical Reactions One of the most fascinating features of a chemistry class is the opportunity to carry out chemical reactions. Many of these produce unexpected and exciting results. This experiment will investigate several characteristic reactions. Chemists classify most reactions into five major types, based on what reacts and what is produced: Combination: Two elements and/or compounds combine to form a new compound Element or Compound + Element or Compound Compound Decomposition: A compound is broken down into two simpler substances Compound Element or Compound + Element or Compound Single Replacement: An element reacts with a compound Element + Compound Different Element + Different Compound Double Replacement: Two compounds react together, producing two new compounds Compound A+ Compound B Compound C + Compound D Combustion: A hydrocarbon (something containing carbon and hydrogen) reacts with oxygen gas, producing carbon dioxide and water Hydrocarbon + Oxygen Carbon Dioxide + Water It will be your job today to describe several different reactions, and then indicate which category best describes them. Enjoy, but be safe! Procedure Carry out each reaction in a clean test tube unless otherwise indicated. Then, record observations on your report sheet. Be as specific as possible, indicating the state(solid, liquid, gas), color, and other interesting observations of the chemical products. Reaction #1: This reaction must be carried out in the fume hood. Place a few pieces (3 to 4) of copper metal into a test tube. Slowly add a few drops of 6 M nitric acid (HNO 3 ) to this tube. You should see a quick and vigorous chemical reaction. Record your observations. Allow the reaction to go to completion and any gases to dissipate before removing your test tube from the fume hood. Any liquid may go down the drain with running water, but solids should go in the trash. R eaction #2: Place a few pieces of copper metal in a different test tube. Add about 2 ml (40 drops should be fine) of 0.1 M silver nitrate solution, (AgNO 3 ) to the tube. Do not get this on your skin, or it may cause dark spots to appear! Set the tube aside for several minutes. Record your observations when you are done. Pour any liquids in the container marked Silver waste ; solids should go in the trash. R eaction #3: Read the directions for this reaction completely before beginning! This reaction must be carried out by a group of two students. Add a piece of zinc (sometimes called mossy zinc ) to 5 ml of Page 5-1

35 6 M hydrochloric acid (HCl) in a test tube. Place your thumb over the mouth of the test tube, but do not shake it. A partner should be ready to light a match and hand it to you. You will soon (about 40 seconds into the reaction) feel pressure building up in the tube. Move the test tube near the flaming match and quickly release your thumb so that you do not burn it. Do not stick the match into the test tube, and do not point the test tube at anyone (including yourself)! Any liquid wastes may go down the drain with running water, solids into the trash. Reactions 4 & 5: Pour 5 ml of distilled water into your evaporating dish. Then, add one drop of phenolphthalein to it. Ignite a small piece of magnesium by holding it in a Bunsen burner flame with tongs. Do not look directly at the reaction. After the reaction has completed, observe the product you have formed, then drop it into the evaporating dish, where yet another reaction will occur. A change in the color of the solution indicates that you have made a basic solution (one which contains hydroxide ions). A base can be thought of as the opposite of an acid. The products may go down the drain. Reaction 6: Place 20 drops of 0.1 M copper (II) sulfate solution (CuSO 4 ) into a test tube. To this, add exactly one drop of 6 M ammonia solution (NH 3 ), and stir with a stirring rod. Record your observations, but do not dispose of the contents of this test tube. Reaction 7: Add 10 more drops of 6 M NH 3 to the test tube from Reaction 6, and stir with a glass stirring rod. Record your observations, then pour the liquids down the drain with running water. Reaction 8: Pour 2 ml (40 drops) of 0.1 M lead (II) nitrate (Pb(NO 3 ) 2 ) in one test tube, and 2 ml of 0.1 M potassium iodide (KI) in another. Record what each of these solutions looks like, then combine the contents of each. Pour the waste in the lead waste Reaction 9: Place 5 drops of 0.1 M iron (III) chloride (FeCl 3 ) in a clean test tube. Using your spatula, add one small crystal of ammonium thiocyanate (NH 4 SCN) to the solution, and stir with a clean stir bar. Dispose of the products in the container labeled Reaction 9 Waste. Page 5-2

36 Report Sheet Name: Introduction to Chemical Reactions Record your observations below, and indicate the type of reaction which best describes each. You are not required to provide the reactions type for reactions six, seven, and nine; the equations for these are somewhat complex. Reaction One: Word equation: copper + nitric acid copper (II) nitrate + nitrogen dioxide Type of reaction: Observations: Reaction Two: Word equation: copper + silver nitrate silver + copper (II) nitrate Type of reaction: Observations: Page 5-3

37 Reaction Three: Word equation: zinc + hydrochloric acid hydrogen + zinc chloride Type of reaction: Observations: Reaction Four: Word equation: magnesium + oxygen magnesium oxide Type of reaction: Observations: Reaction Five: Word equation: magnesium oxide + water magnesium hydroxide Type of reaction: Observations: Page 5-4

38 Reaction Six: Observations: Reaction Seven: Observations: Reaction Eight: Word equation: lead (II) nitrate + potassium iodide lead (II) iodide + potassium nitrate Type of reaction: Observations: Reaction Nine: Observations: Page 5-5

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40 Preparation for this experiment: Pieces of copper foil (or thin strips) [1-2 small containers] Pieces of mossy zinc [1-2 small containers] Pieces of magnesium ribbon [1-2 small containers] Ammonium thiocyanate [1-2 small containers] Phenolphthalein solution in dropper bottles 2-3 dropper bottles ( ml) containing each of the following aqueous solutions: 6 M HNO M AgNO 3 6 M HCl 0.1 M CuSO 4 6 M NH 4 OH (label as 6 M NH 3 ) 0.1 M Pb(NO 3 ) M KI 0.1 M FeCl 3 Waste containers: Lead Waste, Silver Waste, Reaction 9 Waste Page 5-7

41 Experiment 6: Paper Chromatography In this experiment you will explore the technique of chromatography, a very important method for the separation and identification of substances. There are many different forms of chromatography, but they are all based on the same principles. A mixture is generally dissolved in a solution or absorbed onto a solid surface. The mixture is then pushed across a stationary surface, causing it to separate into its components. The different components demonstrate a differing level of attraction for the stationary surface; those that are more attracted to the surface tend to be pulled more slowly across it. In this experiment you will carry out paper chromatography, which uses paper as the stationary surface and a solvent as the mobile phase. The compounds you will be separating are the colored dies commonly used in foods and in inks. PROCEDURE: Part A: Chromatography of Ink from Felt-Tip Markers 1. Cut a strip of chromatography paper from the roll. It should be approximately 15 cm long. Fold the paper about 3 to 4 cm from the bottom into an L shape. The paper should be able to stand up on the short side. 2. Draw a line (in pencil) about 1 cm from the fold. On the line, make three small dots with felt-tipped pens of different colors. These dots should be about equally-spaced from each other, with none of them too close to the edge of the paper. 3. Pour a small amount of water into your 150 ml beaker. You should add only as much water as is necessary to completely cover the glass at the bottom of the beaker. 4. Place the bottom of the chromatography paper into the beaker. The water-level must be lower than the dots on the paper. The water will push the ink up the paper. Try to keep the paper from touching the sides of the beaker. 5. Remove the paper from the water once the water level has reached about 1 cm from the top of the paper. Do not allow the water to rise all the way to the top! Sit the paper aside and allow it to dry. 6. Answer the questions on the report sheet. Part B: Separation of Colored Mixtures (Read all the way through before beginning!) 1. Using a scrap piece of chromatography paper and a capillary tube, practice making small spots of dye solution on the paper. A good spot should be round and no more than 2 mm wide. You may use any of the five dye compounds to practice this. 2. On the 14 cm 10 cm chromatography paper, draw a straight baseline in pencil--do not write on the paper with ink-approximately 1.5 cm up from the bottom of the sheet. (Consider the 14- cm edges to be the top and bottom of the sheet.). Refer to Figure 6-1 for an example. 3. Using separate capillary tubes for each solution, place one spot of each of the following solutions along your baseline: tartrazine (label as T ) erioglaucine (label as E ) allura red (label as A.R. ) erythrosin B (label as E.B. ) sunset yellow (label as S.Y. ) Page 6-1

42 In applying the spots you should follow these guidelines: Use a fresh capillary tube for each solution. Place the first spot at least 2 cm in from the left edge of the paper and the last spot at least 2 cm in from the right edge of the paper. Leave about 2 cm of space between neighboring spots. Write the labels for each spot along the bottom edge of the paper Figure 6-1: Preparing the Chromatograph 5. Holding the paper upright with the spots facing you, curve the two vertical edges away from you and toward each other until they just touch each other (make sure they do not overlap ). Staple the touching sides together at the top and bottom of the cylinder you ve created. See the figure below. Figure 6-2: Chromatograph Folded and Stapled 6. Add a small amount of the ammonia solution to an 800 ml or a l-l beaker (the liquid should not be more than 0.5 cm deep). Stand your paper cylinder alongside the beaker and check to see that the level of the ammonia in the beaker is below the baseline on the cylinder. If the ammonia level is above the baseline, pour out some of the solution and check again. Refer to Figure 6-3 for this and the next two steps. Page 6-2

43 Figure 6-3: Developing the Chromatograph 7. Once you are sure your spots won't be submerged, place the cylinder in the beaker, baseline towards the bottom. You may ask your instructor for help if you are unclear about what to do. 8. Cover the beaker with plastic wrap or aluminum foil. Leave the setup undisturbed once the solvent has begun moving up the paper. Excessive motion can ruin the chromatograph. 9. When the solvent has traveled to just below the top of the paper cylinder (about 2 cm from the top) remove the cylinder from the beaker and allow it to dry. 10. Repeat this experiment, using the four unknown mixtures (A, B, C, and D) instead of the pure colorings. Label each spot with the letter of the unknown. 11. Once you have finished developing the chromatograph, it should be possible to identify which of the five colorings were present in each mixture. Answer the questions on the report sheet. Page 6-3

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45 Report Sheet Name: Paper Chromatography Attach your dry chromatograms to this page with a stapler or paper clip. 1. Which of the inks (blue, green, and black) appear to be mixture of other colors? Which individual colors appear to be in each ink? 2. Which of the colorings were present in each of the four mixtures? Mixture A Mixture B Mixture C Mixture D 3. Why did you use pencil instead of pen when writing on the chromatography paper? 4. All of these colorings were made and purified in a chemical laboratory. Alternatively, they could have been extracted from natural sources such as plants and then purified. The Food and Drug Administration directs food suppliers to label color additives as natural colors if they are obtained from natural sources, and artificial colors if they are man-made. Which (if either) is safer for you to consume? Explain your reasoning. Page 6-5

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47 Preparation for this Experiment Felt-tip pens (blue, black, and green) mL or 1-L beaker per student Roll of chromatography paper and scissors Capillary Tubes (preferably mini or prepared by me) 1.0 M NH 4 OH (about 500 ml-1 L) Label this as ammonia solution Solutions of each of the dyes in small bottles and unknowns A, B, C, and D (should already be prepared see me if you need to make new solutions) Plastic wrap (or aluminum foil) to cover beakers 2 pieces of chromatography paper per student plus several extra for those who must redo experiment(each cut to 14 cm 10 cm) Page 6-7

48 Experiment 7: The Clock Reaction In Experiment Five you observed several fascinating chemical reactions, most of which seemed to occur almost instantaneously. In today s experiment, you will carry out the socalled Iodine Clock reaction, which relies on a color change to indicate that the reaction has completed. Several different factors have can have an effect on the speed of this reaction; we will attempt to test each of these factors to determine which speed up and which slow down a given reaction. The chemical reaction is rather complex; it can be summarized (in words) as iodate ion + bisulfite ion iodine + sulfate ion + water + hydronium ion. Although you may not understand the details of this reaction, you will be able to observe how changing aspects of the reaction might cause the reaction to go more quickly or more slowly. We will be using a starch solution as an indicator in this reaction; its presence causes a color change at the end of the reaction which would not be observed otherwise. We will study indicators more in depth in later experiments. Procedure Note: You do not have to exclusively use 50 ml beakers throughout this experiment; you may substitute other sizes if necessary. A. The Effect of a Catalyst on Reaction Rate Test One With Catalyst 1. Pour 20 ml of 0.5% potassium iodate (KIO 3 ) solution into a 50-mL beaker, and add 8 drops of copper (II) sulfate (CuSO 4 ) solution (the catalyst) to the beaker. 2. Pour 20 ml of 0.2% sodium bisulfite (NaHSO 3 ) solution into a second 50-mL beaker, and add 5 drops of starch solution to the beaker. 3. Record the start time shown on your stopwatch. 4. Start the stopwatch as you pour the contents of the first beaker into the second beaker, stirring the mixed solutions continuously. Stop the stopwatch when the reaction mixture turns deep blue. 5. Record the stop time and calculate the amount of time required for the color change to occur. Test Two Without Catalyst Repeat the steps of the previous part; however, on step one, do not add the copper (II) sulfate solution. The reaction should proceed more slowly; however, we have not observed this to consistently occur. Page 7-1

49 B. Effect of Concentration on Reaction Rate Test 1 1. Pour 20 ml of 1.0% potassium iodate solution into a 50-mL beaker, and add 8 drops of copper (II) sulfate solution to the beaker. 2. Pour 20 ml of 0.2% sodium bisulfite solution into a second 50-mL beaker, and add 5 drops of starch solution to the beaker. 3. Record the start time shown on the stopwatch. 4. Start the stopwatch as you pour the contents of the first beaker into the second beaker, stirring the mixed solutions continuously. Stop the stopwatch when the reaction mixture turns deep blue. 5. Record the stop time and calculate the amount of time required for the color change to occur. Test 2 1. Pour 20 ml of 0.5% potassium iodate solution into a 50-mL beaker, and add 8 drops of copper (II) sulfate solution to the beaker. 2. Pour 20 ml of 0.4% sodium bisulfite solution into a second 50-mL beaker, and add 5 drops of starch solution to the beaker. 3. Record the start time shown on your stopwatch. 4. Start the stopwatch as you pour the contents of the first beaker into the second beaker, stirring the mixed solutions continuously. Stop the stopwatch when the reaction mixture turns deep blue. 5. Record the stop time and calculate the amount of time required for the color change to occur. Test 3 1. Pour 20 ml of 1.0% potassium iodate solution into a 50-mL beaker, and add 8 drops of copper (II) sulfate solution to the beaker. 2. Pour 20 ml of 0.4% M sodium bisulfite solution into to a second 50-mL beaker, and add 5 drops of starch solution to the beaker. 3. Record the start time shown on your stopwatch. 4. Start the stopwatch as you pour the contents of the first beaker into the second beaker, stirring the mixed solutions continuously. Stop the stopwatch when the reaction mixture turns deep blue. 5. Record the stop time and calculate the amount of time required for the color change to occur. Page 7-2

50 C. Effect of Temperature on Reaction Rate Test 1: Below Room Temperature 1. Determine the ambient air temperature in the laboratory and record it in column 2, row 2, of the Part C data table. 2. Fill a 500-mL beaker about one-fourth full with crushed ice, and add about 300 ml of water. (It is fine to "eyeball" how much water you add. The beaker should be about threefourths full after you add the water.) 3. Pour 20 ml of 0.5% potassium iodate solution into a 50-mL beaker, and add 8 drops of copper (II) sulfate to the beaker. 4. Pour 20 ml of 0.2% sodium bisulfite solution into a second 50-mL beaker, and add 5 drops of starch solution to the beaker. 5. Place both beakers in the shallow tray, and carefully pour ice water into the tray until the water level in the tray is about even with the liquid level in the beakers. Leave the beakers in this ice bath until the solutions reach a temperature of about 10 C. Record the temperature of the solutions in column 2, row 3, of the Part C data table. 6. Record the start time shown on your stopwatch. 6. Start the stopwatch as you pour the contents of the first beaker into the second beaker, stirring the mixed solutions continuously. Stop the stopwatch when the reaction mixture turns deep blue. 7. Record the stop time and calculate the amount of time required for the color change to occur. Test 2 Above Room Temperature 1. In a 500-mL beaker, heat about 300 ml of water to about 65 C. 2. Repeat steps 3 and 4 of Part C Test Place both beakers in the shallow tray, and carefully pour the hot water into the tray until the water level in the tray is about even with the liquid level in the beakers. Leave the beakers in this water bath until the solutions reach a temperature of about 40 C. Record the temperature of the solutions in column 2, row 4, of the Part C data table. 4. Record the start time shown on your stopwatch. 4. Start the stopwatch as you pour the contents of the first beaker into the second beaker, stirring the mixed solutions continuously. Stop the stopwatch when the reaction mixture turns deep blue. 5. Record the stop time and calculate the amount of time required for the color change to occur. Page 7-3

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52 Report Sheet The Clock Reaction Name: Data Tables Part A: Effect of Catalyst on Reaction Rate Test Catalyst Added? Amount of Time Required (sec) 1 Yes 2 No Part B: Effect of Concentration on Reaction Rate Test % KIO 3 % NaHSO 3 Amount of Time Required (sec) Part A, Test 1 0.5% 0.2% % 0.2 % % 0.4 % 3 1.0% 0.4 % Part C: Effect of Temperature on Reaction Rate Test Temperature (ºC) Amount of Time Required (sec) Part A, Test 1 Room: Below Room Temp Above Room Temp Conclusions How is the rate of the reaction affected by the presence of copper (II) sulfate? Page 7-5

53 In many reactions, the use of a catalyst speeds the reaction up by a tremendous amount. Based on your data, do you think it was essential to use the catalyst in Parts B and C? Explain Briefly. How is the rate affected by doubling the concentration of KIO 3 NaHSO 3 What appears to happen to the reaction rate when the concentration of both of these chemicals is doubled? How is the reaction rate affected by lowering and raising the temperature? Try to explain these results based on differences at the molecular level. Page 7-6

54 Preparation Required for This Lab stopwatches ice Each in three regular bottles without droppers: 1.0% KIO 3 solution (about grams KIO 3 per 1 L water) about 50 ml/student 0.5% KIO 3 solution (dilute 1.0% solution by a factor of 2) at least 100 ml/ student 0.4% NaHSO 3 solution (about 3.33 grams NaHSO 3 per 1 L water ) about 50 ml/student 0.2% NaHSO 3 solution (dilute 0.4% solution by a factor of 2) at least 100 ml/student In three dropper bottles, about ml per bottle: M CuSO 4 solution 1% starch solution Page 7-7

55 Experiment 8: Isolation of Caffeine from Mountain Dew The soft drink Mountain Dew is a solution of several compounds in water, including carbon dioxide, food colorings, and the common stimulant caffeine. The combination of caffeine, sugar and carbonation give the beverage its kick, making it a popular drink. Our goal in this experiment will be to extract the caffeine, separating it from the remainder of the soft drink solution. Caffeine cannot be separated from Mountain Dew by the separation methods we have used thus far (evaporation and filtration), so a new method is introduced. We treat the soda with methylene chloride (CH 2 Cl 2 ), an organic solvent which is immiscible with water. While caffeine is soluble in water, it is even more soluble in methylene chloride. Agitating the mixture will encourage the caffeine to migrate to the organic solvent, leaving the other dissolved compounds in the water layer. We then separate the organic layer from the water layer and evaporate the solvent, leaving solid crystals of caffeine behind. 1. Obtain a can of Mountain Dew (opened or unopened). Measure 75 ml of the soda in your large graduated cylinder. Give the remaining soda to another student. Also, check to make sure that your 50 ml beaker is completely dry. If not, use paper towels to remove any water and allow it to air dry completely. You will not need it until step Pour the soda from your graduated cylinder into your 250 ml Erlenmeyer flask. Swirl and shake the flask for about 5 minutes, until foaming is significantly reduced. Do this over a sink and try not to spill any! 3. Use the electronic balances to measure out approximately 1.5 grams of sodium carbonate (Na 2 CO 3 ) onto a piece of weighing paper. Then add it to your 250 ml Erlenmeyer flask. Swirl the soda mixture until the sodium carbonate dissolves. 4. Pour 20 ml of methylene chloride into a graduated cylinder. Caution: Methylene chloride is a suspected carcinogen. Wash your skin immediately if any should splash on you. Unlike most organic solvents, methylene chloride is not flammable. 5. Pour the methylene chloride into your 250 ml flask. The liquids in the flask should separate into two distinct layers, with the more dense methylene chloride on the bottom. Swirl the flask for about 10 minutes to thoroughly mix the liquids. Do not shake, or you may form an emulsion which is difficult to remove at a later stage. 6. Set the flask aside and allow it to sit for about 10 minutes. In the meantime, prepare a gravity filtration set-up as follows: a. Place an iron ring on a ring stand. The ring should be about 8 inches above the bottom of the stand. b. From the cart, obtain a clay triangle. Place the triangle on the iron ring. c. Insert the stem of your funnel through the clay triangle. Place a 250 ml beaker below the stem. This beaker does not need to be dry. d. Obtain a piece of filter paper. Fold it into quarters, then open it into a cone (your instructor will show you how to do this) and place it into the funnel. e. Use your wash bottle to thoroughly wet the filter paper with distilled water. The paper should cling to the side of the funnel. This step is essential for success! Page 8-1

56 7. If an emulsion (a thick layer of small bubbles) has formed in the flask, carry out the following steps. Otherwise, skip to step 8. Add about 1 gram of sodium chloride to the flask (NaCl), swirl for about 2 minutes, and let the flask sit for another 5 minutes. Repeat this process a second time if necessary. It is alright to proceed even if the emulsion has not completely broken up. 8. Carefully try to separate the two layers by pouring only the top layer into a large beaker (any beaker will do, except for your 50 ml beaker). It is important that the bottom layer remain in the Erlenmeyer flask. It is alright if a small amount of the top layer remains in the Erlenmeyer flask as it will be removed in the next step. 9. Gently pour the liquid remaining in the Erlenmeyer flask into the filter you setup in step 6. Any remaining water should migrate to the filter paper and separate from the methylene chloride. However, the methylene chloride should not pass through the wet filter paper. You may need to carefully stir the liquid in the filter with your stirring rod to help separate any water trapped in the methylene chloride. 10. Write your initials on your dry 50 ml beaker, and weigh it on the balance. Record the mass on your report sheet. 11. Using your dropper, transfer the methylene chloride solution remaining in the filter cone to the 50 ml beaker. Leave behind any emulsified portions which did not separate. 12. Set the beaker on a warm hot plate in a fume hood to evaporate the methylene chloride. The hot plate should be at a low-to-middle setting. Keep your eye on the beaker, especially after more than half of the solvent has evaporated Once all the solvent has been evaporated you should immediately remove the beaker from the hot plate. It should not be too hot for you to pick up with your fingers. The beaker should contain a residue of caffeine powder. Do not allow the beaker to sit too long or some of your caffeine may sublimate. 14. Allow the beaker to cool for about 3 minutes. Then, weigh the beaker on the same balance you used in step 10. Record the mass on the report sheet. 15. Have your instructor take a look at the crystals in your beaker and initial your report sheet. Then, clean up your lab area. Any left over solution (such as those collected in steps 8 and 9) are mostly water and should go down the drain. The caffeine may also be washed down the drain. Any left over methlyene chloride (if you took too much in step 4) should go in a waste container in the hood marked halogenated organic waste. Page 8-2

57 Report Sheet Name: ISOLATION OF CAFFEINE FROM MOUNTAIN DEW Volume of your Mountain Dew sample: Mass of dry 50 ml beaker: Mass of 50 ml beaker + caffeine residue: Mass of caffeine residue: ml g g g Divide the mass of the caffeine residue by 1,000 to convert to milligrams of caffeine mg Total volume of soda in a 12 oz. can in milliliters (read off of the can) ml Determine the number of milligrams of caffeine in one can of Mountain Dew. To do this, set-up the ratio below and solve for x. milligrams of caffeine residue milliliters of soda in your sample = x milligrams of caffeine in soda can milliliters of soda in one can Show your calculation mg caffeine/can Instructor s Initials: Page 8-3

58 Questions 1. Pure caffeine is a white, powdery solid. Based on this information, how pure would you say your sample is? 2. The caffeine collected in this experiment is often not white. What is the source of the impurity that gives the impure caffeine its coloring? 3. In step 13, you were warned not to overheat your sample on the hot plate. What does the term sublimate mean in this instruction? 4. It is possible to purify the caffeine by sublimation. Write down a short procedure demonstrating how this could be accomplished. Hint: Review the procedure for subliming iodine in Experiment 4: Elements and Compounds. Page 8-4

59 Materials Required for This Experiment Cans of Mountain Dew (1 can for every 3 students) Methylene Chloride (about 25 ml per student) in 3 or 4 bottles, no droppers attached 4 small containers of Na 2 CO 3 (about 2 grams/student) 4 small containers of NaCl (table salt, not Rock Salt) (about 2 grams/student) Clay triangles (1 per student) Small waste container in hood labeled as Halogenated Organic Waste for leftover CH 2 Cl 2 ALSO: Please place 3-4 hot plates in the hoods. Turn them on to a low-medium setting before the class begins. Page 8-5

60 Experiment 9: A Chemistry Investigation Qualitative analysis in chemistry is very much like a detective game. The characters in a detective story have methods of operation and other distinguishing features. These characteristics make it possible to identify individuals as having been responsible for certain acts. The clues that one observes are evidence of some kind of interaction. In qualitative analysis you will make use of clues, evidence of chemical interaction, to help you identify the presence of specific ions in solution. However, before you can expect to identify the presence of ions, you must first become familiar with their characteristic behavior (method of operation). In this experiment you will be given four solutions, labeled 1, 2, 3, and 4. You will discover how they behave when they are mixed with four reagents labeled A, B, C and D. By making careful observations you will detect evidence of chemical reaction that will be characteristic for each of the solutions. These clues may be the formation of precipitates, change in color, production of a gas, production of heat or other evidence of chemical reaction. Recording your observations in tabular form will help recognize characteristic behavior for each of the four solutions. The solutions and reagents used in this experiment are deliberately identified by letter or number. This will permit you to focus your attention on the reasoning involved in developing an analytical scheme rather than on the actual reaction taking place. The scheme will permit you to identify unknown solutions according to their differentiating properties. Consider the following hypothetical case, reagents X, Y and Z, were allowed to react with a few drops of each of the solutions I, II, III and IV. The observations were recorded in the table as shown. SOLUTIONS/REAGENTS X Y Z I - + WHITE + YELLOW II + WHITE - + YELLOW III + GASES - + YELLOW IV - + WHITE - The + means distinctive evidence of reactions was observed and a - indicates no observable evidence or reaction. The color of precipitates or changes in color of solutions was also noted. If you were given an unknown solution which you were told was like one the four solutions tested, what method would you outline to identify the unknown? How many tests would you have to use? If another unknown gives a + test with X, is this sufficient evidence to identify it as one of the four solutions? Procedure 1. Obtain a set of solutions and reagents. Set aside a 400 ml beaker as a waste beaker. Page 9-1

61 2. Using the height of a 2mL sample of water in a test tube as a reference, place about 2mL of solutions 1, 2, 3 and 4 into a set of four clean test tubes and label them as 1, 2, 3 and 4. Add about 2 ml of reagent A to each test tube. Observe and record on the report sheet your results under the first column as + or and include a brief description of any characteristic property of the + test. Pay attention to the formation of gases, precipitates, color changes or temperature changes of the mixtures. 3. Dispose the waste in the waste beaker, clean the test tubes. Place about 2mL of the solutions 1, 2, 3 and 4 to each test tube. Add about 2mL of reagent B to each test tube. Observe and record your results under the second column as + or and include a description of any characteristic property of the + test. 4. Continue performing tests on each of the solutions 1, 2, 3, and 4 using reagents A, B, C and D until you have tested all possible combinations. Record your results as + or and include a description of any characteristic property for the + tests under the appropriate column of your table. 5. Study the data carefully and note the identifying clues. Obtain an unknown from your instructor. Record your unknown number and test the unknown with reagents A, B, C and D to determine which solution it is like. Report your analysis together with a summary of the supporting evidence. Page 9-2

62 Chemistry Investigation Report Sheet Name: SOLUTIONS/REAGENTS A B C D SOLUTIONS/REAGENTS A B C D UNKNOWN# Which solution is your unknown? Explain how you came to this conclusion. Page 9-3

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64 Chemicals for this lab: 15mL /student in dropper bottles (at least 3 of each) Label the following solutions as Solution 1 for 1.0 M sodium carbonate Solution 2 for 0.1M NaCl Solution 3 for 0.1 M barium chloride Solution 4 for 0.1 M lead(ii) nitrate 15 ml/student in dropper bottles (at least 3 of each) Label the following reagents as Reagent A for 0.1M HCl Reagent B for 0.1 M silver nitrate Reagent C for 0.1 M sodium sulfate Reagent D for 0.1 M potassium iodate Unknowns: 10ml/ student from the solution 1,2,3 or 4. Please label them with unknown numbers. Page 9-5

65 Experiment 10: Esters & the Preparation of Soap Esters are often noted for the pleasant aroma they produce. They are often prepared by reacting an alcohol with a carboxylic acid; a small amount of acid is added to catalyze the reaction. In this experiment you will prepare several esters and describe their properties. In addition, you will prepare soap, through what is called a saponification reaction. This is accomplished by adding a strong base to fat or oil, which contain structural elements similar to esters. This method of soap production has been practiced for years; it was not unusual for the pioneers to prepare their own soap from animal fat to clean their pans. Procedure Part A: Ester Synthesis 1. Heat about 400 ml of water in a 600-mL beaker over a Bunsen burner. The water should not be boiled; a temperature around 85 C should suffice. 2. Label four clean test tubes with letters A-D, and obtain stoppers or corks for each. 3. Add the following chemicals to each of the four test tubes: Test tube A: Test tube B: Test tube C: Test tube D: 20 drops of ethanol and 20 drops of glacial acetic acid 20 drops of methanol and a small amount of salicyclic acid (about the size of a match head) 20 drops of 1-pentanol and 20 drops of glacial acetic acid 20 drops of ethanol and a small amount of benzoic acid (about the size of a match head) 4. Add the indicated number of drops of 85% phosphoric acid to each test tube. Perform this slowly in the hood. Phosphoric acid can burn the skin, so use it carefully. Notify your instructor if any should spill. A: 10 drops B: 3 drops C: 5 drops D: 10 drops 5. Gently stir each tube with a stirring rod. Turn off the Bunsen burner once the temperature reaches 85 C. 6. Loosely stopper each test tube and put them into the hot water bath. 7. After 8 to 10 minutes, remove each tube from the water bath. Wave the fumes from each test tube to you and describe the aroma in your report sheet. 8. Draw the structure of each ester produced. Page 10-1

66 Part B: Soap Preparation and Properties 1. Determine the mass of a dry evaporating dish. 2. Add about 3 grams of lard to this evaporating dish. 3. Pour 10 ml of ethanol into the evaporating dish. Carefully add about 1 ml of 50% sodium hydroxide solution (be sure to wear your goggles!) 4. Gently heat the mixture on a hot plate until it becomes a thick paste. You should occasionally stir the mixture with your stirring rod. Do not overheat. Then, allow the evaporating dish and its contents to cool. Describe the appearance of this product on your report sheet. 5. Make a soap solution by combining the paste with 50 ml of deionized water in the beaker provided for this experiment. Heat the mixture on a hot plate for about 5 minutes (or more), stirring the contents often. The mixture will foam, but most of the paste should dissolve. Remove the beaker from the hot plate and allow the mixture to cool. You may also stir in a small amount (about 1-3 ml) of a perfume if you wish to provide your soap with a pleasant smell. 6. You will now extract soap through a procedure called salting out. Slowly add solid sodium chloride (NaCl) to the solution, a little at a time, stirring as you go. Keep adding the salt until no more appears to dissolve (this is tricky, as you may initially have a hard time telling the soap and the salt apart.) 7. Setup a vacuum filtration according to your instructor s directions. Attach the side arm of the filter flask to the water aspirator (or an alternative vacuum source if your instructor suggests one). Refer to Figure 8-1 for the setup (of course, you will not have any liquid in your filter flask when you begin!). Figure 10-1: Vacuum Filtration 8. Collect a small piece of the soap on the tip of your stirring bar. Add this to a beaker and pour in some water. Compare what you see with familiar household soap. Page 10-2

67 Esters and Soaps Report Sheet Name: Part A: Ester Synthesis Test Tube #1: Describe the aroma of the ester produced: What is the structure of the ester produced? Test Tube #2: Describe the aroma of the ester produced: What is the structure of the ester produced? Page 10-3

68 Test Tube #3: Describe the aroma of the ester produced: What is the structure of the ester produced? Test Tube #4: Describe the aroma of the ester produced: What is the structure of the ester produced? Page 10-4

69 Part B: Soap Preparation and Properties Describe the appearance of the soap paste you made in Step 4. Describe the appearance of the soap formed after salting out In what ways does this product behave like household soap products? In what ways is it different? Would it be a good idea to use the soap you just formed to wash your face? Why or why not? Be sure to think about the chemicals you used to produce it and what may remain in the left-over residue. Page 10-5

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71 Preparation for this lab: one vacuum filtration set-up per student with filter paper 3 dropper bottles of each of the following (about 50 ml/bottle) 85% phosphoric acid in hood ethanol (95%) methanol 1-pentanol 3 bottles (about 250 ml each) of 95% ethanol 3 bottles (about 100+ ml each) of 50% NaOH solution Several (at least 5) medium-sized containers of NaCl (about 100 grams in each) lard (3 grams/student) 2 dropper botters of glacial acetic acid in hood 3 small (10-25 gram) containers of salicylic acid 3 small (10-25 gram) containers of benzoic acid organic waste container Page 10-7

72 Experiment 11: Measuring Acidity with Red Cabbage Juice Many common household chemicals can be described as acids or bases. For example, many cleaning supplies are basic (sometimes called alkaline). Vinegar is a 5% solution of acetic acid (the remainder is water). Shampoos and cosmetic products are usually ph balanced to a level which minimizes irritation of the skin. In this experiment, you will test various substances and designate them as an acid, a base, or neither (neutral). Acids are substances that can donate a hydrogen ion in water; bases are substances that can accept a hydrogen ion or produce hydroxide ion in water. The amount of hydrogen ion donated or accepted determines the strength of the acid or base. The concentration of hydrogen ion in a solution is described by its ph value. The ph of a solution is mathematically defined as where, [H + ] is the concentration of the hydrogen ion (i.e. the amount of hydrogen ion per liter). The negative logarithmic relationship means that every time the ph increases by 1, the hydrogen ion concentration decreases by a factor of 10. A lower ph value corresponds to a higher acidity of the solution. A ph above 7 is basic, and ph below 7 is acidic. For a neutral solution, the ph is 7. Many naturally occurring molecules are influenced by the level of the acidity around them. The hydrangea flower has a red color in acidic soil, but blue in basic soil. Molecules that change color in response to changes in acidity level are called indicators. Some indicators change color only at one particular ph, whereas others exhibit numerous numbers of colors over a broad ph range. Red cabbage leaves contain a compound that exhibits different colors at different ph values. In this experiment, you will examine the color changes which Procedure: accompany the ph changes. Note that your instructor may recommend that you work in pairs for some or all of this experiment. A. Determine the indicator s color at various ph values 1. Obtain about 5 large red cabbage leaves and tear them into small pieces. Place them in a 400mL beaker and cover them with DI water. 2. Heat the cabbage water until the water boils. Boil it for 5 to 7 minutes, then turn off the burner. Page 11-1

73 3. Pour the cabbage juice into a 250-mL beaker and discard the leaves. Allow the juice to cool to room temperature. 4. Obtain 12 clean test tubes and a test tube rack. Label the test tubes from 1 to 12 to correspond with the ph values of the solutions you will be testing. 5. Add about 2 ml of the cabbage juice to each of the test tubes. Then, add about 2mL of the standard solution labeled ph=1 to test tube 1, and ph=2 to test tube 2, and so on. Think of a way to do this step without measuring the volume of solutions repeatedly. 6. Record the colors of the solutions in the test tubes in the data table on the report sheet Save these test tubes to use as standards for comparison to the household substances you will be testing in Part B. B. Determining the ph values of household substances 1. Obtain 10 additional test tubes and label them to indicate which household substance is in each. 2. Add 2mL of cabbage juice to each of the 10 test tubes. 3. If the substance is a liquid, add 2ml of it to a test tube. If the substance is viscous (i.e. syrupy, like baby oil), add only few drops to a test tube and use a stirring rod to swirl the substance in the test tube. If the substance is a solid, add a small spatula tip of it (about the size of a pea) to the test tube containing the cabbage juice. 4. Record the color of each household substance with cabbage juice on the report sheet. Compare these solutions to the standard solutions from part A and estimate the ph based on the color. 5. Categorize your household substances as acids, bases, or neutral. Page 11-2

74 Report Sheet Name: Measuring Acidity with Red Cabbage Juice A. Determine Indicator s Color at Various ph B. ph of Household substances ph Color ph Color substance Color in Cabbage Juice ph Vinegar Ammonia Lemon Juice Apple Juice 7-UP Ivory Liquid Detergent Shampoo Hair Conditioner Antacid Aspirin Mouthwash Windex Wisk Cleaners Ajax Corn Syrup Acid. Base or Neutral? Page 11-3

75 Questions Suppose that you have two samples of vinegar, although each contains a different volume of it. Sample A contains 100 ml, while Sample B contains 500 ml. Other than the difference in volume, the vinegar sample are identical. Is one sample more acidic than the other? If so, then identify the more acidic one. Give a brief explanation of your answer. Five drops of an indicator is added to each of the samples in the previous question, and both solutions turn orange. The color in Sample B is fainter than that of Sample A, although the actual shade of the color is the same. Explain these two observations. Does the information given in the second question seem to be consistent with what you would expect from your answer to the first question? Explain briefly. Page 11-4

76 Materials for this experiment: Please refer to chem 110 lab -- We will use the same standard buffers and the same household good for this experiment. Cabbage We will not use ph paper nor ph meters. Page 11-5

77 Experiment 12: Percentage of Water in Popcorn The science behind how popcorn kernals pop is quite simple. Each kernal contains a small amount of water trapped inside of it. Heating the kernals causes the water to become high-energy molecules of steam. As we continue to heat the kernals, the pressure exerted by the trapped steam on the inside of the kernal becomes so great that it eventually causes it to burst. What remains is, of course, a popular snack. Unfortunately, not all kernals in a given sample of popcorn will pop, regardless of how much heat we apply. These duds may contain very small cracks, through which any water already may have escaped. Any water remaining in these kernals easily escapes through the cracks as steam when we heat the popcorn. Without this highpressure build up of steam, the kernal remains unpopped. In today s experiment we will determine the percentage (by mass) of water in a sample of popcorn by comparing the mass of kernals before and after popping. We will also determine the average number of duds in a given sample of popcorn. PROCEDURE 1. Using an electronic balance, determine the mass of your dry evaporating dish and record it on the report sheet. 2. Obtain exactly ten kernels of popcorn and place them in the evaporating dish. Weigh the dish with the kernels inside it, and record the mass on the report sheet. 3. You will now pop your kernels by constructing a hot air popper using a ring stand assembly, a beaker, the evaporating dish, and a watch glass. It is a good idea to put an additional iron ring around the beaker to stabilize it if the assembly is inadvertently shaken. Be sure to choose the size of beaker which the evaporating dish will fit best into, as you see in the diagram. The dish should not fall in to the beaker, nor should it be too large to rest on its rim. Do not pour water into the beaker! Page 12-1

78 4. Place something fairly heavy on the watch glass; otherwise, the popping kernels may knock it off and shatter it. Good choices include metal cylinders, very-large rubber stoppers, etc. 5. Turn on the Bunsen burner, and adjust the flame so that the tip of the inner cone is at the level of the wire gauze. The air in the beaker will gradually become hot enough to pop the popcorn. Note that this may require ten minutes or more. 6. Continue heating until all kernels have popped, or until the popcorn begins to brown. Do not burn the popcorn to a black crisp. Not only will that ruin your experiment, it will produce a terrible odor in the lab room! It is likely that a few duds will remain unpopped. 7. Allow the apparatus to cool for at least 10 minutes. Remove the evaporating dish from the assembly and weigh it with the popped corn and any duds inside. Be sure that you do not inadvertently include the watch glass in your weighing. Record this mass on the report sheet. Additionally, you should record the number of duds in your sample. 8. Repeat this process a second time, using a different brand of popcorn if one is available. You should clean the evaporating dish as best as you can with a dry paper towel. a. Popped popcorn may go in the trash. Do not eat it! Page 12-2

79 Report Sheet Name Percentage of Water in Popcorn Always give the appropriate units! Mass of empty evaporating dish Trial 1 Trial 2 Mass of evaporating dish + kernels Mass of kernels Calculation: Mass of dish, popped kernels, and duds (after heating) Mass of popped kernels and duds Calculation: Mass of water lost Calculation: % water in popcorn Calculation Average % water in popcorn: Page 12-3

80 Number of kernels used Trial One Trial Two Number of duds remaining Trial One Trial Two Percentage of kernels which were duds Trial One Trial Two Questions Suppose that you have two samples of the same type of popcorn. One sample contains 100. grams of kernels, and the other contains 200. grams. Which contains (i) the greatest mass of water, and (ii) the greatest percentage of water? Is the percentage of water an intensive or extensive property? Why should you not eat any of the popcorn you produced in this experiment. Be as specific as possible. Page 12-4

81 Materials for this lab: Popcorn (preferably two or more different brands): 25 kernels per student Very large rubber stoppers, metal cylinders, or any other heavy objects which will help to weigh down the watch glass (see step 4). Page 12-5

82 Experiment 13: Vitamin C Content of Beverages Vitamins are a class of structurally-diverse chemicals which are an essential component of our diet. Each vitamin has its own particular function and properties. For example, some vitamins are soluble in fats and oils (including Vitamins A and E), while others will dissolve in water. Vitamin C, also known as asorbic acid, is among the latter. This vitamin has been celebrated for years for its ability to prevent colds and deter infection. Linus Pauling, the greatest chemist of the 20 th century, devoted many of the final years of his life to studying and promoting it. As you probably already know, citrus fruits (and their juices) are the main source of Vitamin C for most people. In addition, many alternative beverages, including Tang and Kool-Aid contain Vitamin C. In today s experiment, we will attempt to determine the number of milligrams of Vitamin C present in a typical 8-ounce serving of a fruit juice or beverage. We will also analyze a standard Vitamin C tablet. For reasons which will become clear as we continue, this experiment tends to work best on liquids with little or no color. You will observe the reaction of Vitamin C with iodine (I 2 ) in the presence of starch. As you add iodine to the juice, you gradually transform all of the Vitamin C. When none of the vitamin is left, the iodine will interact with the starch, causing the solution to change color. This will be the signal that the reaction is complete. The Reaction: Procedure Part A Vitamin C in a Commercial Vitamin C pill 1. Obtain a Vitamin C pill. Note the number of milligrams of vitamin C each pill should have according to the container. Crush the pill into a fine powder with in a mortar. Carefully scrape all of this powder onto a piece of weighing paper. 2. Transfer the powder to a 250 ml or 125 ml Erlenmeyer flask. Add about 50 ml of distilled water to the powder and swirl for about 2 minutes to dissolve as much of the powder as possible. Some of the material may not dissolve; this is probably a filler material used to hold the tablet itself together. Finally add about 1 ml of the 1% starch solution (the indicator). Page 13-1

83 3. Set up a buret according to your instructor s example. Add the iodine solution to the buret; the solution should fill the buret to a level between the 0 and 1 ml markings. Be careful: the iodine solution stains clothing and skin! You should use a glass funnel to help you in this process.record the value of this starting point on the report sheet. 4. Add iodine solution to the flask, carefully controlling the rate with the stopcock (the small handle on the buret). Swirl the flask gently as you add the solution. At first, you will see color changes which quickly fade. As you get closer to the endpoint of the titration, the color will seem to persist for a longer period of time. You should slow down the rate at which you are adding the iodine solution. The titration is complete when the solution turns a deep blue; it must remain blue for at least 30 seconds, during which time you should gently swirl the flask. Record the ending point reading on the buret on your report sheet. 5. Thoroughly rinse the flask and the buret with water. Part B Vitamin C in Fruit Juice or Sports Drink 1. Using a graduated cylinder, measure out exactly 50.0 ml of your sample. Pour this into the Erlenmeyer flask you used in Part A. Add 1 ml of 1% starch solution and swirl. Note: Depending on your juice, it may be necessary for you to add 2 ml of 0.1 M acetic acid. Only add it if the experiment does not seem to work and you are running it over again. 2. The iodine solution you used to titrate the Vitamin C pill is too concentrated for use in this part of the experiment; you must first dilute (water down) the solution. Measure out exactly 10.0 ml of the iodine solution you used in Part A with a the buret, and transfer this to your 150 ml beaker. Add exactly 90.0 ml of water to this (measuring the volume with your 100 ml graduated cylinder) and stir with a glass stirring rod. We will call this the diluted solution ; it is 10 times less concentrated than the solution you used in Part A. Pour the diluted solution into the buret as you did in step 3 of the previous part and record your starting point. 3. Using the buret, add the iodine solution to the juice solution you prepared in step one. While controlling the flow from the buret with one hand, you should use the other to gently swirl the flask as you did in step 4 of Part A. 4. When the solution changes color and does not return to the original color, you have reached the end point of the titration. You should swirl the solution for another thirty seconds. If the solution in the flask does not return to the original color, then you are done. Otherwise, it will be necessary to add a few more drops of the iodine solution from the buret. The color maybe different from that of Part B if your juice or sports drink has any color to it. However, the color change should be very distinct. Record the ending point of the titration. Page 13-2

84 5. Repeat this procedure at least once (twice if possible). You should rely on the information from the first trial to predict how much iodine you will need to add. You may use the same Erlenmeyer-flask, but be sure to pour the contents down the drain and rinse it several times with distilled water first. 6. Follow the procedures indicated in the report sheet to determine the number of milligrams of Vitamin C present in one 8 oz. serving of the drink. Page 13-3

85 Page 13-4

86 How Much Vitamin C? Report Sheet Name: Part A Vitamin C Pill 1. Milligrams of Vitamin C in each pill (according to label) mg 2. Final Buret Reading(mL) 3. Starting Buret Reading (ml) 4. Total Volume of Iodine Solution Added (#2 minus #3) 5. Each milliliter of the iodine solution should react with 14 mg of Vitamin C. Find the amount of Vitamin C in the pill by taking the value you got in #4 and multiplying it by 14. Record this value on the right. mg Vit. C/pill Part B Fruit Juice or Sports Drink 1. What juice or other product did you use? 2. After diluting the iodine solution, figure out how many milligrams of vitamin C should react with 1 ml of the diluted solution. (Think: How much did I dilute the solution by? ) mg Vit C/mL diluted Iodine solution Trial One Trial Two 3. Final Buret Reading(mL) 4. Initial Buret Reading (ml) 5. Total Volume of Iodine Solution Added (#3 minus #4) 6. Milligrams of Vitamin C in 50. ml of your juice (#5 divided by #2) Page 13-5

87 7. Average number of milligrams of Vitamin C in 50 ml of your juice (Add up the values for both trials in #6. Then, divide the result by 2) 8. Milligrams of Vitamin C in one 8 fluid ounce serving of this juice (Multiply the value in #7 by 4.73) 9. Percent of USDA-recommended Vitamin C requirement in 1 serving of juice (Divide the value in step 8 by 60. Then, multiply by 100) mg mg % According to your data, does this juice appear to be a good source of Vitamin C? Explain. This method would not work well for purple grape juice, which also contains Vitamin C. Explain why. Is it likely that the same procedure as the one we used today could be used to analyze all vitamins? Briefly explain your answer, using the material you have read about vitamins to assist you. Page 13-6

88 Preparation Instructions: Prepare at least 1 liter of iodine solution for every 10 students in class. The iodine solution is prepared by dissolving 20. grams of iodine and 40. grams of potassium iodide in 1 liter of water. Stir the solution, then gravity filter to prevent burets from clogging up. In addition students will need: 1% starch solution (at least 50 ml per 10 students) 1 Vitamin C pill per student 6 mortar and pestles for class use Students might also need to add a few drops of 0.1 M acetic acid (unlikely) Also, please purchase the following from the supermarket: Any colorless (or clear and very slightly colored) sports drink which contains at least 50% of the USDA Vitamin C requirement (30 mg) per 8 fl. ounce serving White grape fruit juice (Welch s works well) Apple Juice which is fortified with Vitamin C (Tropicana works well, but check the label for Vitamin C) There should be enough juice/sports drink available for each student to have 150 ml of a given liquid (they only need 150 ml of one of these, not 150 ml of each) Page 13-7

89 Experiment Optional #1: Respiratory Gases In this experiment you will generate the respiratory gases carbon dioxide, CO 2, and oxygen O 2. Then, you will examine their properties by carrying out chemical tests on each of the gases. Cellular respiration is the oxidation of fuel molecules to produce energy. Except under anaerobic conditions, oxygen is a required reactant for oxidation, while carbon dioxide and water are the main by-products. Oxygen is transported in the blood of mammals by forming a loose chemical combination with hemoglobin called oxyhemoglobin. Hemoglobin can also combine with carbon dioxide, but only about ten percent of the gas is carried in this manner. Carbon monoxide, CO, also combines with hemoglobin; in fact, it binds so strongly (CO bonds about 200 times stronger than O 2 ) that the attachment of oxygen to the hemoglobin is prevented and oxygen cannot reach the cells. Overexposure to carbon monoxide eventually results in death (asphyxiation). PROCEDURE: Part One: Carbon Dioxide 1. Set up the generation apparatus as shown below. Your instructor will likely demonstrate the generation of gas and call your attention to various important details. Be sure to pay close attention to what they say and ask for help if you are unsure as to what to do. CAUTION: The Thistle tube is easily broken if not handled with care; you may become seriously cut in the process! When adjusting the thistle tube in the stopper be sure to lubricate the stem with water or glycerin. Protect your hand with a towel and slowly twist the thistle tube as you push it in or pull it out very gently. Let your instructor know if your tube is particularly stiff. The bottom of the tube must be below the surface of the liquid but should not touch the bottom of the flask, but be close to it when the stopper is in place. Connect one end of the rubber tubing to the side arm of the flask and immerse the other in tap water in the pneumatic trough 2. Remove the stopper and place about 15g of marble chips in the flask. Add about 25 ml of deionized water. Replace the stopper and check to see that the bottom of the thistle tube is below the water level. This is a very important safety precaution to keep gas and acid from backing up. 3. Fill a wide-mouth bottle with water, cover it with a glass square and invert it in the pneumatic trough. Remove the glass. Do the same thing with another bottle and also prepare two test tubes for gas collection in the same manner. Page Optional#1-1

90 4. Pour about 15 ml of dilute (6 M) hydrochloric acid through the thistle tube. Caution: hydrochloric acid causes burns; wear your goggles at all times. Gas will begin to bubble into the pneumatic trough. Place one of the bottles over the bubbles until it is full of gas. Discard this first bottle as it is mostly air. 5. Collect a second bottle of gas. Cover it with the glass plate under water, then remove and turn it right side up while still covered with the glass plate. Also collect gas in two test tubes and stopper them. If at any time the reaction should cease, slowly add 5 or 10 more milliliters of acid through the thistle tube. 6. Pour 5 ml of lime water in one empty test tube (one that you did not fill with carbon dioxide), shake and observe. Combine the contents of one of the above test tubes that has the collected carbon dioxide with 5 ml of lime water, shake, and observe. Pour the contents into the sink. 7. Place another 5 ml of lime water into a clean test tube. Observe the appearance. With a straw gently blow into the lime water. Again observe the appearance. Clean the test tube. 8. Obtain a splint and ignite it with a match. Carefully lower it into a wide mouth bottle of air until it goes out. Add 2-3 ml of lime water. Cover the bottle with a glass plate and shake the bottle to mix the gas with the lime water. Observe the appearance. 9. Invert a bottle of carbon dioxide on top of a bottle of air. Slide the glass plate out of the way so that the two bottle mouths are joined. After 30 seconds remove the upper bottle and cover both bottles with glass plates. Test each bottle with 2-3 ml of lime water and shake, as above. 10. Thrust a glowing splint into the second test tube and observe the results. Part Two: Oxygen 1. Clean the gas-generation apparatus from Part One and rinse the flask with fresh DI water. 2. Pour about 25 ml of DI water into the side-arm flask. Then add about 1 gram of manganese dioxide powder (MnO 2 ) to the flask. Note that essentially none of it will dissolve in the water. Stopper the flask as you did in Part One. 3. Pour about 10 ml of 9% hydrogen peroxide (H 2 O 2 ) solution into the thistle tube. The hydrogen peroxide will decompose when it comes into contact with the manganese dioxide, forming water and bubbles of oxygen gas. This gas will travel through the tube attached to the side-arm into the collection bottle. Collect one bottle full of gas and, as before, discard it. 4. Collect two bottles of oxygen gas. You may add more hydrogen peroxide solution (5-10 ml) if the reaction stops to start the generation again. Please do not waste hydrogen peroxide solution! 5. Read the directions for this step completely before beginning work. Quickly take the glass plate Page Optional#1-2

91 off of one of the bottles of oxygen. Then, holding the bottle sideways, insert a burning splint into the bottle. You may be seriously burned if you thrust the burning splint down into an upwards-pointing bottle! 6. Obtain a stopwatch and a small candle. Light the candle. Cover the candle with a bottle containing only air. Use the stopwatch to measure how long it takes for the candle to burn out from the moment you cover it with the bottle. Then, repeat this step using a bottle of oxygen. 7. Your instructor may also choose to demonstrate the burning of sulfur in a bottle of pure oxygen. Page Optional#1-3

92 Page Optional#1-4

93 Laboratory Report Respiratory Gases Name Part One: Carbon Dioxide A.1 Write the balanced chemical equation for the reaction of marble chips with hydrochloric acid. A.2 From the result of this experiment, would you think carbon dioxide is soluble in water or not? Explain. B.1 Observation Write the balanced chemical equation for the reaction of carbon dioxide with lime water. B.2 What do you observe after you blow air into the lime water with a straw? Explain what is the component in the exhaled gases that causes the reaction? B.3 Observation: Is carbon dioxide produced from the burning of wood? Is carbon dioxide produced by oxidation outside the body (wood) as well as inside the body (breath)? Explain. B.4 Was there carbon dioxide in both bottles? B.5 Observation: Explain why fire extinguishers utilize carbon dioxide. Page Optional#1-5

94 Part Two: Oxygen Write a balanced chemical equation for the decomposition of hydrogen peroxide into water and oxygen gas. What was the purpose of adding MnO 2 to the flask? What term best describes this compound s role in the reaction above. Describe what happened when you inserted a burning splint into the bottle of oxygen. How long did it take for the candle to be extinguished in each bottle? Explain why these values differ. In 1996, an airplane operated by the carrier ValuJet plunged into the Florida Everglades, killing all 110 passengers and crew aboard. After a thorough investigation, the incident was blamed on improperly stored oxygen generators (containers which rapidly produce oxygen gas when activated) in the cargo area. Explain how these generators could contribute to such a disaster. Page Optional#1-6

95 MATERIALS Required for This Experiment: Thistle tube and 1-hole stopper, suction flask, pneumatic trough, wide mouth bottles, glass plates, test tubes, stoppers for test tubes, marble chips (CaCO 3 ), wooden splints, limewater, plastic straws, steel wool (?), rubber tubing, 6M hydrochloric acid (HCl) Irene put steel wool on this list; I m not sure that I ll need it but please have some available so that I can use it for a demonstration. Thanks! I have added a few components from Experiment 3 in the CHEM 120 lab manual to this experiment: sulfur (one container for instructor use only) and one deflagrating spoon. Containers of MnO 2 and 9% H 2 O 2 solution (about 20 ml per student should be sufficient) small candles Stopwatches Page Optional#1-7

96 Experiment Optional #2: The Synthesis of Aspirin The natural world provides us with many of the medications in common use today. Taxol is the common name of a medication used in treating certain cancers; it is derived from the Yew tree. Most controlled substances, including heroin and cocaine, are extracted from plants. The deadly toxin associated with botulism, produced by bacteria, has even found a use in cosmetic surgery, although in a very diluted form. Salicylic acid, a compound found in willow bark, has long been known for its ability to reduce fever and suppress pain. Unfortunately, it is also tends to irritate and damage the upper digestive tract. A slight modification to its structure produces acetylsalicylic acid, which retains the benefits of the original compound with significantly less irritation. This popular medication goes by the common name Aspirin. Aspirin becomes salicylic acid in the lower digestive tract, where it is absorbed into the blood stream. Procedure 1. Fill a large beaker with tap water. Place the beaker on a hot plate and heat the water inside to boiling. You should carry out steps 2 to 5 while it is heating. At some point you should also cool about 75 ml of DI water with ice for use in steps 8 and Weigh a 125-mL Erlenmeyer flask on the balance. Record this value on the report sheet. Remember, you should always write down all digits! 3. Add about 2 grams of salicylic acid to this flask, and reweigh. You can find the total mass of salicylic acid by subtraction. 4. In the hood, measure 5 ml of acetic anhydride into your small graduated cylinder. Then, slowly pour this into the Erlenmeyer flask from steps 2 and While still in the hood, add 10 drops of 85% phosphoric acid to this mixture (do not worry if you should accidentally add a few extra drops). Stir the mixture with a stirring rod. Page Optional #2-1

97 6. Place the flask in the hot water bath and stir the contents until all solids dissolve. Then, remove the flask from the bath and allow it to cool. 7. In the hood, slowly add 20 drops of DI water to the flask. 8. When the reaction is complete, pour in an additional 50 ml of ice-cold DI water. Then, place the flask in an ice bath (beaker with ice and some water) to cool for 10 minutes. 9. Slowly stir the contents. You should notice the formation of white, aspirin crystals. If not, gently scratch the inside of the flask with the glass stirring rod. This should induce crystals to grow. If this still does not work, ask your instructor to help you by seeding the mixture. 10. Setup a vacuum filtration according to your instructor s directions. Attach the side arm of the filter flask to the water aspirator (or an alternative vacuum source if your instructor suggests one). Refer to Figure 12-1 for the setup (of course, you will not have any liquid in your filter flask when you begin!). Figure Optional 2-1: Vacuum Filtration 11. Wash the crystals with enough ice-cold water to just completely cover the crystals. Allow the water to drain into the same flask you used in the previous step. 12. Carefully scoop the crystals onto your watch glass. These crystals are primarily composed of aspirin. Page Optional #2-2