PREDICTIONS: Make sure you read through the lab and understand the objectives before making your predictions.

Transcription

1 1 LAB EXPERIMENT 1 MAKING OBSERVATIONS OBJECTIVES: 1. To make observations while watching materials interact and undergo change. 2. To understand, recognize and record qualitative and quantitative observations. 3. To make interpretations based upon observations and data. PREDICTIONS: Make sure you read through the lab and understand the objectives before making your predictions. MATERIALS: Apparatus Reagents 250 ml beaker aluminum foil (about a 10 cm X 10 cm sheet) Thermometer 1 M copper (II) chloride solution Mass balance PROCEDURES: 1. Get a piece of aluminum foil and record the mass. 2. Place about 75 ml of copper (II) chloride solution into a clean 250 ml beaker, and record the temperature of the solution. Make and record observations about the chemicals that you are using before you start step Crumple the aluminum foil into a loose ball and place it into the beaker containing the copper (II) chloride. 4. Immediately start recording your observations in Table 1. As well, make and record observations about the chemicals as the reaction is taking place. Make sure you record the highest temperature achieved. 5. When no further changes appear to be happening, make and record observations about the chemicals once the reaction is complete. 6. Place all the contents of the beaker in the waste disposal jar provided by your instructor. 7. Before you leave the lab, wash your hands thoroughly with soap and water from the back sink in the room.

2 2 OBSERVATIONS AND DATA: Volume of copper (II) chloride used: Mass of aluminum foil used: TABLE 1: Temperature of the solution before, during and after the reaction TIME (minutes) TEMPERATURE ( C) Maximum temperature achieved: QUESTIONS AND CALCULATIONS: 1) Knowing the temperature change that occurred over a four minute span, calculate the rate of temperature change per minute? 2) Where should the chemicals from the beaker be placed after the reaction is complete? Why? 3) If you used a stronger (more concentrated) solution of copper (II) chloride with the same piece of aluminum foil, what changes in observations might you observe? CONCLUSIONS: 1) Write out one significant qualitative and one quantitative observation from this lab. 2) What is the difference between these two types of observations? 3) Write out an interpretation of the data that you could take from this experiment. 4) State any sources of error that may have affected this lab.

3 3 LAB EXPERIMENT 2 COOLING AND HEATING CURVES OF A PURE SUBSTANCE OBJECTIVES: 1. To investigate the cooling process for liquid paradichlorobenzene. 2. To investigate the heating process for solid paradichlorobenzene. 3. To determine and compare the melting and freezing temperatures of paradichlorobenzene. PREDICTIONS: Make sure you read through the lab and understand the objectives before making your predictions. MATERIALS: Apparatus Reagents Figure 1 Ring stand and support paradichlorobenzene Test tube clamp Test tube (18 mm x 150 mm) 1000 ml beaker Thermometer PROCEDURES: PART 1: THE COOLING PROCESS 1. Obtain a test tube of paradichlorobenzene from your instructor. Remove the stopper and save until the end of the experiment. 2. You will need to melt the paradichlorobenzene in order to place the thermometer into the test tube and record the temperature. To do this, place the test tube clamp around the test tube and attach it to the ring stand, immersing the test tube in the hot water bath (It should be around C) until all the substance has melted. Make sure the water level is above the level of the paradichlorobenzene. 3. Insert the thermometer into the test tube and record the temperature as the zero time of your cooling curve in your data table. Remember to hold the thermometer off of the glass. See figure Quickly remove the test tube from the hot water bath and move over to the warm water bath (It should be around C). Attach your clamp to the ring stand and immerse the test tube in the water, again making sure the water level is higher than the paradichlorobenzene. Begin taking temperature readings and make observations every 30 seconds until all of the paradichlorobenzene has solidified. You will have to hold the thermometer until the substance has solidified around it. Keep taking the readings until the temperature has reached around 49 C. (The thermometer should be embedded in the solid)

4 4 PART 2: THE HEATING PROCESS 1. Once your thermometer has been embedded in the solid, take a final temperature reading. This will be your zero temperature reading for the heating curve data. 2. Remove the test tube from the warm water bath and move back to the hot water bath. Immerse the test tube in it (making sure the water level is above the paradichlorobenzene). Begin taking temperature readings and observations every 30 seconds until all of the paradichlorobenzene has melted and a temperature of 60 C has been reached. 3. When the temperature has reached 60 C, remove the thermometer from the liquid paradichlorobenzene and immediately wipe it with a paper towel. 4. Raise the test tube out of the hot water bath and allow the paradichlorobenzene to resolidify. When it has cooled, put the stopper back on and return it to the instructor. 5. Before you leave the lab, wash your hands thoroughly with soap and water. OBSERVATIONS AND DATA: TABLE 1: Cooling and heating of paradichlorobenzene Cooling Process Time Temperature Observations (min) ( C) ANALYSIS OF DATA: PART 1: THE COOLING PROCESS Heating Process Observations Temperature ( C) 1. Using the data you obtained during the cooling process, construct a graph of temperature versus time. Use small circles for these data points, and sketch a smooth curve. 2. Indicate on the graph where solidification began and ended

5 5 PART 2: THE HEATING PROCESS 1. On the same graph as the cooling curve, plot temperature versus time for the heating process. Use small squares for these data points, and sketch a smooth curve. 2. Indicate on the graph where melting began and ended. QUESTIONS: 1. From your cooling curve, determine the freezing point of paradichlorobenzene. 2. From your heating curve, determine the melting point of paradichlorobenzene. 3. Compare your freezing and melting points with those of two other lab groups and explain any similarities or differences. 4. What can you conclude about the melting and freezing points of a pure substance? 5. How would you explain the plateaus in your heating and cooling curves? 6. Suppose more paradichlorobenzene had been used in Part 1. What would be the appearance of the new cooling curve? CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this lab. Don t forget any sources of error that may have occurred.

6 6 LAB EXPERIMENT 3 CHEMICAL AND PHYSICAL CHANGE OBJECTIVES: 1. To observe some changes in the laboratory. 2. To infer whether each is a chemical or physical change. 3. To record some recognizable characteristics of chemical changes. PREDICTIONS: Make sure you read through the lab and understand the objectives before making your predictions. MATERIALS: Figure 1 Apparatus Reagents 1 spot plate or glass square set of 4 unknown solutions PROCEDURES: 1. Obtain a clean spot plate or glass square. Place a piece of coloured paper underneath the glass square if necessary. On the paper draw a grid as seen in Figure 1. If you use the spot plate, label the chemicals you will be combining above the well of the spot plate. 2. On the glass square or spot plate, combine the solutions by mixing A with B, A with C, A with D, B with C, B with D, and C with D. Record your observations in your data table (there are only 6 observations necessary). Make sure you use separate, clean droppers for each of the solutions. 3. Once you have your data, rinse the solutions down the sink with lots of water. Before you leave the lab, make sure you wash your hands thoroughly with soap and water. OBSERVATIONS AND DATA: Table 1 Unknown A B C D A B C D

7 7 QUESTIONS: 1. State whether a physical or a chemical change occurred in each of the six combinations of solutions. 2. Describe two chemical changes that you might observe occurring in everyday life. 3. Describe two physical changes that you might observe occurring in everyday life. CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this lab. Don t forget any sources of error that may have occurred.

8 8 LAB EXPERIMENT 4 SEPARATION OF A MIXTURE BY PAPER CHROMATOGRAPHY PRE-LAB READING AND QUESTIONS Chromatography is one technique used by chemists to separate mixtures of chemical compounds in order to identify or isolate their components. In chromatography, mixtures are separated according to the different solubilities of the components in liquids, or their adsorptions on solids. Chromatography has many applications, including detection and measurement of pesticides in foods, and drugs and urine specimens. It is also used extensively in biological research to separate alcohols, amino acids, and sugars. As well, the pharmaceutical industry relies on chromatography for the production of high-purity chemicals. There are a variety of chromatographic techniques, but all share two features: a moving carrier phase, and a stationary phase. In the stationary phase of paper chromatography, the sample to be analyzed is spotted onto a piece of chromatography paper. The sample is carried along this stationary phase by a solvent which acts as the moving carrier. The components of the sample are carried different distances along the paper, depending on their individual solubilities. After a length of time, the original spot is spread out into a series of bands. These bands are then analyzed to determine their identities. In paper chromatography, one method of identifying these separated components of a mixture is to calculate the R f value of each. R f stands for ratio of fronts. An R f value is simply the ratio of the distance travelled by the solute to the distance travelled by the solvent: R f = d 1 /d 2 where d 1 = distance travelled by solute d 2 = distance travelled by solvent The R f value of a substance is a characteristic of that substance for a specific solvent. A substance having a high solubility in the moving phase will be carried further and will have a higher R f value. By definition, R f values vary from 0 to On the basis of what principle is chromatography used to separate mixtures? 2. What industry uses chromatography to produce very pure chemicals? 3. a) What constitutes the moving phase in this experiment? b) What constitutes the stationary phase in this experiment? 4. a) What does R f mean? b) Write the mathematical equation for calculating an R f value and explain all the terms. 5. What substances and mixtures will be tested in this experiment? 6. Why do you need a pencil for this experiment?

9 9 OBJECTIVES: 1. To assemble and operate a paper chromatography apparatus. 2. To study the meaning and significance of R f values. 3. To test various food colourings and to calculate their R f values. 4. To compare measured R f values with standard R f values. 5. To separate mixtures of food colourings into their components. 6. To identify the components of mixtures by means of their R f values. PREDICTIONS: Make sure you read through the lab and understand the objectives before making your predictions. MATERIALS: Apparatus Scissors Ruler Pencil Unknown mixture Chromatography paper 3 large test tubes (20mm x 200mm) 3 Erlenmeyer flasks Reagents Set of food colourings (yellow, red, blue, green) PROCEDURES: PART 1: SETTING UP 1. Obtain 3 large test tubes and 3 Erlenmeyer flasks. Place a test tube in each of the flasks and label the test tubes A, B, and C. 2. Obtain a 60 cm length of chromatography paper and cut it into 3 strips of 20 cm each. Then cut each strip so that they are 2.0 cm in width. Using a pencil, lightly draw a line across each strip 4.0 cm from one end. Use scissors to trim this end of the strip into a point, as shown in the figure below. Label these A, B, and C with a pencil cm sample dye spot 2.0 cm 4.0 cm

10 10 3. Your instructor will assign you one colour (red, yellow, or blue) for you to test. Using a capillary tube, spot your first strip of chromatography paper (labelled A) on the pencil line with the colour assigned (see the figure above). The spots should not exceed 0.5 cm in diameter. 4. For the strip labelled B, spot with the colour green. For the strip labelled C, spot with the unknown mixture. PART 2: R f VALUES OF INDIVIDUAL FOOD COLOURINGS AND SEPARATION OF MIXTURES INTO THEIR COMPONENTS 1. Add about 2 cm of water to each of the test tubes. Put strip A into test tube A so that the tip just touches the bottom of the test tube. Do not allow the dye spot to be immersed in the water. Do not allow the flat surface of the strip to rest against the walls of the test tube. 2. Do the same for strips B and C into test tubes B and C respectively. Observe what happens to your sample as the water moves slowly up the paper as a result of capillary action. 3. See whether or not your samples separate into component colours. After about 30 minutes have elapsed, remove strip A from the test tube and immediately draw a pencil line across the top edge of the solvent front (the water line) and the top of the colour line. 4. Repeat Step 3 for strip B and strip C, making lines across all the component colours and the water line. 5. Measure with a ruler the distances d 1 (the distance the colour travelled) and d 2 (the distance the water travelled) on strip A. Record these results in Table 1 and calculate the R f value. Record the class data in Table Repeat Step 5 for strip B, measuring all the component colour distances and calculating R f values for each component colour in the green food colouring. Record your results in Table Repeat Step 5 for strip C, measuring all the component colour distances and calculating R f values for each component colour in the unknown mixture. Record your results in Table Identify the dyes used in this lab by comparing your calculated R f values with those in Table 4. Record those results in Table Clean up your materials. The strips may be placed in the garbage.

11 11 OBSERVATIONS AND DATA: TABLE 1: Assigned Food Colouring Colour Tested Solute Distance (cm) Solvent Distance (cm) R f Value TABLE 2: Class Data for Individual Food Colourings Tested Station Red R f Yellow R f Blue R f Average R f TABLE 3: Separation of Mixtures Into Their Component Colours Component Colours d 1 d 2 R f Dye ID Green Colouring Unknown Mixture TABLE 4: Some Approved Dyes for Food Colouring Dye Red #2 Red #3 Red #4 Yellow #5 Yellow #6 Blue #1 Blue #2 R f

12 12 QUESTIONS: 1. a) Which, if any, of the colours you tested in this experiment appeared to contain one or more of the approved dyes listed in Table 4? b) Which, if any, of the colours you tested did not correspond to any of the approved dyes? 2. From your results, what are the components of green food colouring? Support your answer both qualitatively and quantitatively. 3. What can you conclude about the identity of the components in the unknown mixture? What qualitative and quantitative evidence supports your answer? 4. What might happen if ink, rather than pencil, were used to mark the sample line on the chromatography paper? 5. Why should green food colouring be classified as a mixture, whereas red, blue or yellow should not? 6. Identify the dyes that appear on the chromatogram in Figure 1 (see Table 4). The original sample was orange food colouring. 7. A pharmaceutical chemist runs a chromatography test on a substance and identifies two of its components by comparing their R f values. If the two components have R f values of 1.00 and 0.41, and the solvent front has travelled 12.0 cm from the sample s origin, what is the separation distance between the components on the chromatogram? 8. A chemist performs an R f calculation, obtains a value of 1.20, and decides that the answer is unacceptable. Why? Figure 1 CONCLUSION: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don t forget any sources of error that may have occurred.

13 13 LAB EXPERIMENT 5 GRAPHING AS A MEANS OF SEEKING A RELATIONSHIP OBJECTIVES: 1. To make measurements of mass and volume for three different liquids. 2. To analyze the data by means of graphing techniques. 3. To determine a mathematical relationship between mass and volume for each liquid. PREDICTIONS: Make sure you read through the lab and understand the objectives before making your predictions. MATERIALS: Apparatus Reagents 250 ml Erlenmeyer flask water mass balance methanol several burettes salt (sodium chloride) solution PROCEDURE: 1. Your instructor will assign you a volume of liquid to test. You will then use this volume for all three liquids. The data will be shared with the class and therefore you will be depending on each other for good results. 2. Determine and record the mass of a clean, dry 250 ml Erlenmeyer flask. 3. Go to one of the burettes that contains water, methanol, or salt solution (the order you do these in is not important) and get your assigned volume of liquid as accurately as possible. If you do not get the precise amount it does not matter. What does matter is that you record exactly the volume that you did get in Table Mass out the total mass of the flask and the liquid. Subtract the mass you found in Step 2 in order to determine the mass of the liquid. Record this in Table Now repeat Steps 3 and 4 for the other two liquids. Do not empty the flask each time you add a different liquid just keep determining the mass of each volume by subtracting the previous mass balance reading. Record your results in Table A data table similar to Table 2 will be on the chalkboard. Record your results from Table 1 into the chart on the board. 7. Once all groups have finished, copy the completed Table 2 into your lab report. 8. Clean up all your materials and pour the contents of the flask down the sink with plenty of water. Wash your hands with soap and water before leaving the lab.

14 14 OBSERVATIONS AND DATA: Table 1: Results for Lab Station Water Alcohol Salt Solution Volume (ml) Mass (g) Volume (ml) Mass (g) Volume (ml) Mass(g) Table 2: Class Results Water Alcohol Salt Solution Lab Station Volume (ml) Mass (g) Volume (ml) Mass (g) Volume (ml) Mass (g) ANALYSIS OF DATA: 1. Following the rules of good graphing, plot a graph showing mass vs. volume for each liquid using the class data. Plot the results for all three liquids on the same graph, making sure to differentiate between each liquid. 2. Draw best-fit lines for each graph and calculate the slope of each. Remember, slope is calculated by: m = slope = Δy (change in y values) Δx ( change in x values) Drawing deltas on your graph and subtracting your values to find your change in y and your change in x does this. 3. Determine the mathematical relationship between the mass and volume for each liquid. Remember: y = mx + b where y = the y variable m = the slope of the line x = the x variable b = the y-intercept

15 15 4. The slope of your graph is actually the density. Compare these slope values with the accepted values and do a percent error on them. Percent Error = experimental value accepted value X 100 accepted value The actual (accepted) density of water is 1.00 g/ml. The actual density of methanol is 0.78 g/ml. The actual density of salt solution is 1.10 g/ml. QUESTIONS: 1. Use your graph to predict the mass of 6.5 ml of methanol. 2. Use your mathematical relationship to calculate the mass of 6.5 ml of methanol. 3. Compare your answers to Questions 1 and 2. Explain why they might not be identical. CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don t forget any sources of error that may have occurred.

16 16 LAB EXPERIMENT 6 INVESTIGATING MASS CHANGES IN CHEMICAL REACTIONS OBJECTIVES: 1. To observe a chemical reaction in a sealed flask. 2. To determine the change in mass that occurs during a chemical reaction. PREDICTIONS: Make sure you read through the lab and understand the objectives before making your predictions. MATERIALS: Apparatus Reagents 250 mlerlenmeyer flask One of the following pairs of solutions: rubber stopper for flask 1. A. barium chloride 2 test tubes (13 mm 100 mm) B. sodium sulphate tongs 2. A. lead acetate mass balance B. potassium iodide 3. A. iron (III) nitrate B. potassium thiocyanate 4. A. calcium chloride B. sodium carbonate PROCEDURES: 1. Get 2 test tubes and label them A and B. 2. Your instructor will assign you a set of solutions to react. Half fill test tube A with solution A. Half fill test tube B with solution B. 3. Pour solution B into the Erlenmeyer flask. 4. Carefully lower test tube A into the flask using a pair of tongs so that you do not spill any into solution B. Place the rubber stopper on the flask. 5. Determine the mass of your assembled apparatus and record this value in Table After making certain that the stopper is secure, gently turn the flask upside down and mix the chemicals. 7. Determine the mass of your apparatus after you have mixed the chemicals and record this value in Table Calculate the mass gained or lost during the reaction and record this in Table 1. Record your results in Table 2 on the blackboard and into your lab report. You will be analyzing the class data as a whole when answering the questions.

17 17 9. Dispose of the reacted chemicals in the chemical waste disposal container. Make sure to wash your hands with soap and water before leaving the lab. OBSERVATIONS AND DATA: Table 1: Results for Lab Station. Identity of solution A Identity of solution B Mass of apparatus and contents before reaction (g) Mass of apparatus and contents after reaction (g) Change in mass as a result of reaction (+ or -) (g) Table 2: Class Results Lab Station Solution A Solution B Mass Change (g) QUESTIONS: 1. Why is it important that the flask be sealed for this experiment, even after the flask is returned to an upright position after mixing? 2. What observations lead you to believe that a chemical reaction occurred in the flask? 3. In general, what overall mass change results from a chemical reaction? 4. Suppose that a reaction was carried out in an open flask and the final mass was significantly greater than the initial mass. What would you conclude? 5. If a reaction was carried out in an open flask and the final mass was significantly less than the initial mass, what would you conclude? 6. Write balanced chemical equations to show what happened in each reaction. Assume all reactions are double replacement reactions. CONCLUSIONS: Make sure to answer your objectives, summarize your results, state what you have learned and discuss any sources of error which may have affected your results.

18 18 LAB EXPERIMENT 7 FINDING MASS AND COUNTING PARTICLES This lab illustrates how you can find relative masses of particles. Remember a particle can contain a single atom such as an atom of sodium or a particle can be thought of as a single molecule such as a molecule of water (made up of atoms). The particles you will use today are paper clips. Part B illustrates how the relative masses found in Part A can be used to count out equal numbers of particles. (Note: This is done without counting out paper clips in the beginning.) PREDICTIONS: Make sure you read through the lab and understand the objectives before making your predictions. MATERIALS: 1 box of small paper clips 1 box of large paper clips PROCEDURE: PART A: FINDING RELATIVE MASSES OF PAPER CLIPS 1. Obtain a palm full of large paper clips and a palm full of small paper clips. 2. Clip one large paper clip to one small paper clip and set them aside. 3. Continue pairing one large with one small paper clip, putting them with the other pairs, until all of one size of paper clip is used up. Set the unused ones aside. 4. Now separate the pairs of clips and put the large clips in one pile and the small clips in another pile. 5. Find the mass of large clips and record. Find the mass of the small clips and record. 6. Using the data from step 5, calculate the relative mass of large paper clips to one gram of small paper clips. Show this calculation as a ratio or a fraction. PART B: COUNTING PAPER CLIPS BY MEASURING MASS

19 19 7. Using the results from Part A, predict the mass of large clips that would be necessary to pair up one for one with 10 grams of small clips. 8. Weigh out 10 grams of small clips and your predicted mass of large clips and see if they pair up one for one. 9. If the number of large clips did not pair up evenly with the number of small clips in step 8 (within ± 1 clip), recalculate step 7 and repeat step 8. QUESTIONS: 1. Show the calculations for step A-6 and B What mass of large clips would be necessary to pair up one for one with 50.0 g of small clips? 3. What mass of small clips would be necessary to contain the same number of clips as in 8000 g of large clips? 4. What mass of large clips would be necessary to contain the same number of clips as in 300 kg of small clips? 5. If you were to clip two small clips to each large clip, what mass of small clips would be necessary if you had 70.0 g of large clips? 6. Based on question 5, what would the new relative mass ratio be (mass of large paper clips to one gram of small paper clips)? 7. Based on your answer to question 6, what would the mass of large clips be if you had 12.0 g of small clips? 8. If you can, mass out 12.0 g of small clips and the mass of large clips calculated in question 7. What should be true about the numbers of large clips and the number of small clips present? CONCLUSION: Answer your objectives, state what you have learned, and list any sources of error which may have affected your results.

20 20 LAB EXPERIMENT 8 MOLES OF IRON AND COPPER OBJECTIVES: 1. To determine the number of moles of copper produced in the reaction of iron and copper (II) chloride. 2. To determine the number of moles of iron used up in the reaction of iron and copper (II) chloride. 3. To determine the ratio of moles of iron to moles of copper. 4. To determine the number of atoms and formula units involved in the reaction. PREDICTIONS: Make sure you read through the lab and understand the objectives before making your predictions. MATERIALS: Apparatus Reagents 250 ml beakers copper (II) chloride wash bottle 2 iron nails stirring rod 1 M hydrochloric acid mass balance distilled water steel wool tongs PROCEDURE: DAY 1: 1. Label a clean, dry 250 ml beaker with your name. Find the mass of the beaker and record it in your data table. 2 Add approximately 8 g of copper (II) chloride crystals to the beaker. Find the mass and record it in your data table. 3 Add 50 ml of distilled water to the beaker. Stir the solution with a stir rod until all the crystals have dissolved. 4 Obtain 2 nails and clean them with a piece of steel wool until the surface of the nails is shiny. Find the mass of the clean nails and record it in your data table. 5 Place the nails into the solution and leave them undisturbed for 20 to 30 minutes. During this time copper should be forming in the solution.

21 21 6 Once the time has elapsed, use tongs to carefully pick up one of the nails and use the wash bottle to rinse off any copper from the nail. You may have to use the stir rod to scrape any excess copper off. Place the nail to dry on a piece of paper towel. Repeat this step for the other nail. 7. Carefully decant the liquid from the solid into another beaker. Decant means to pour off only the liquid from a container that contains both a solid and a liquid. Use your stir rod for this process (see figure 1). If you lose some solid, you need to start decanting it all again. Figure 1 8. Rinse the solid with about 10 ml of distilled water. Decant. Repeat this process four more times. 9. Wash the solid with 25 ml of hydrochloric acid, decant, and then wash one last time with 25 ml of distilled water. Decant one last time. 10. Place the beaker in the drying oven until next day. 11. Once the nails are completely dry, find the mass of them and record it in your data table. 12. Pour the decanted solution down the sink with lots of water and clean up. Wash your hands before leaving the lab. DAY 2: 1. Get your dry beaker with the copper in it from the drying oven. 2. Find the mass of the beaker and the copper and record it in your data table. 3. Place the copper in the garbage and clean up your beaker. Wash your hands before leaving the lab. OBSERVATIONS AND DATA: Table 1: Before the Reaction Mass of empty, dry beaker (g) Mass of beaker and copper (II) chloride (g) Mass of two iron nails

22 22 Table 2: After the Reaction Mass of two iron nails (g) Mass of beaker and dry copper (g) QUESTIONS: 1. Find the following masses: a) mass of iron used in the reaction b) mass of copper (II) chloride used c) mass of copper produced in the reaction 2. Find the number of moles of the following: a) moles of iron used b) moles of copper produced 3. Find the number of atoms of each substance involved in the reaction. a) atoms of iron used b) atoms of copper produced 4. Calculate the ratio of moles of copper produced to moles of iron used. 5. Suppose you have an unlimited supply of copper (II) chloride to react with iron. How many moles of copper would be produced by reacting 34.0 g of iron with copper (II) chloride solution? 6. How many moles of iron would have been used up if 45.0 g of copper were to be produced? 7. How many atoms of copper would be involved in question 6? 8. How many atoms of iron would be involved in question 6? 9. How many grams of copper would be produced from the reaction of 456 g of iron? CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don t forget any sources of error that may have occurred.

23 23 LAB EXPERIMENT 9 CALCULATIONS WITH A CHEMICAL REACTION OBJECTIVES: 1. To observe the reaction between solution of calcium chloride and sodium carbonate, forming insoluble calcium carbonate. 2. To calculate the number of moles of each of the starting materials present in the solution. 3. To determine the reactant that is in excess. 4. To determine the theoretical amount of calcium carbonate that could be produced. 5. To compare the theoretical amount to the actual amount of calcium carbonate and calculate the percent yield. PREDICTIONS: Make sure you read through the lab and understand the objectives before making your predictions. MATERIALS: Apparatus Reagents ml beakers 0.60 M sodium carbonate, Na 2 CO 3 wash bottle 0.40 M calcium chloride, CaCl 2 stirring rod with rubber scraper distilled water mass balance ml graduated cylinders ring stand with funnel holder funnel filter paper PROCEDURES: DAY 1: 1. Pour approximately 75 ml of sodium carbonate into a graduated cylinder and record the exact volume in your data table. Pour the sodium carbonate into a clean, dry 250 ml beaker. 2. Pour approximately 50 ml of calcium chloride into the other graduated cylinder, again recording the exact amount in your data table. 3. Pour the calcium chloride into the beaker and describe the resulting reaction in your data table. Stir the contents for about one minute.

24 24 4. Obtain a piece of filter paper and label it with your name in pencil. Find the mass of the filter paper and record it in your data table. 5. Set up your ring stand with the funnel holder. Place the funnel in the holder. Fold your filter paper into a funnel shape and place it in the funnel. Wet the filter paper with a small amount of distilled water to hold it in the funnel (your instructor may show you how). 6. Stir your solution and begin carefully pouring it into the filter paper. Be careful so that none of the solid flows out of the filter paper or funnel. Continue filtering until you get as much of the solid out of the beaker as possible (you may have to use the rubber scraper). 7. Rinse the inside of the beaker with some distilled water and filter. Repeat this two or three more times to remove as much solid as possible. 8 Once all the solid is on the filter paper and the liquid has all drained through into the beaker, carefully remove the filter paper from the funnel and unfold it onto a piece of paper towel. 9. Place the filter paper and paper towel in the drying oven until next day. 10. Clean up your apparatus and pour the filtered solution down the sink with plenty of water. Wash your hands with soap and water before leaving the lab. DAY 2: 1. Get your filter paper with your solid on it from the drying oven. 2. Find the mass of the filter paper and the solid and record it in your data table. 3. Place the solid in the garbage and clean up. Wash your hands before leaving the lab. OBSERVATIONS AND DATA: Table 1: Volume of sodium carbonate solution (ml) Volume of calcium chloride solution (ml) Describe what happens after mixing the solutions Mass of dry filter paper (g) Mass of filter + dry solid (g)

25 25 QUESTIONS: 1. Calculate the following values: d) moles of sodium carbonate used e) moles of calcium chloride used f) mass of calcium carbonate produced g) moles of calcium carbonate produced 2. Write a balanced chemical reaction for the equation that you observed in this lab. 3. Determine which of the reactants was in excess in this reaction. 4. Calculate the amount of calcium carbonate that should theoretically form from the amount of the limiting reactant. 5. Calculate the percent yield in your reaction. 6. Suppose you wanted to add just enough 0.40 M calcium chloride solution to 75.0 ml of 0.60 M sodium carbonate solution for all of the solutions to react. a) What volume of 0.40 M calcium chloride solution would be required? b) What mass of calcium carbonate would be produced assuming 100% yield? c) What mass of sodium chloride would be produced assuming 100% yield? 7. How would you be able to recover the sodium chloride from the solution? 8. A similar reaction occurs when barium chloride is mixed with sodium carbonate solution. Write a balanced chemical equation for the reaction. 9. Calculate each of the following for the reaction that occurs when 56.0 ml of 0.50 M barium chloride is mixed with 78.0 ml of 0.75 M sodium carbonate solution: a) moles of barium chloride added b) moles of sodium carbonate added c) mass of barium carbonate produced, if the yield is 78.0% CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don t forget any sources of error that may have occurred.

26 26 LAB EXPERIMENT 10 THE PERIODIC TABLE OBJECTIVES: 1. To understand the relationship between electron configuration and the location of an element within the period. 2. To examine and graph periodic trends in atomic radii and the first ionization energies. 3. To construct the periodic table published by Mendeleev in 1871 according to a list of clues and your knowledge of the modern periodic table. PREDICTIONS: Make sure you read through the lab and understand the objectives before making your predictions. MATERIALS: Apparatus Graph paper Blank periodic table sheet for electron configurations Blank 1871 version of the periodic table Reference book (Heath Chemistry) PROCEDURES: PART 1: ELECTRON CONFIGURATION AND THE PERIODIC TABLE 1. Obtain blank periodic table for the first 4 rows. 2. For each of the elements listed, write the electron configurations in the space provided. PART 2: TRENDS OF THE PERIODIC TABLE 1. Obtain 2 pieces of graph paper. 2. Using the data listed in Table 1, construct a graph of atomic radius vs. atomic number on the first piece of graph paper. 3. Using the data listed in Table 1 and the second piece of graph paper, construct a graph of first ionization energy vs. atomic number.

27 27 Table 1: ELEMENT Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminum Silicon Phosphorous Sulphur Chlorine Argon Potassium Calcium ATOMIC NUMBER ATOMIC RADIUS (nm) FIRST IONIZATION ENERGY (kj/mol) PART 3: MENDELEEV S PERIODIC TABLE 1. Obtain a blank 1871 version of the periodic table. 2. Using the clues below, place the name of the 67 known elements of 1871 in the correct space. You will need to use the reference book provided by your instructor. 1A has a single electron in the 1s sublevel 1B derived its name from the Latin word for stone, lithos 1C can be collected as a silver liquid in the electrolysis of salt 1D has a first ionization energy of 418 kj/mol 1E has a density of 8.96 g/cm 3 1F is the first alkali metal with a completed 3d sublevel 1G was originally identified by its Latin name, argentum 1H is an alkali metal located in period 6 of the modern periodic table 1I derived its name from the Latin word for dawn, aurora 2A was used as a target substance in the experiments by Irene Joliot-Curie 2B is an alkaline earth metal located in period 3 2C is the metallic component of the substance limestone 2D is a transition metal with 30 protons 2E is represented by the symbol Sr

28 28 2F possesses a nuclear charge of +48 2G is an alkaline earth metal found in period 6 of the modern periodic table 2H is a liquid metal, originally called hydrargyrum 3A has a single electron in the 2p sublevel 3B is a lightweight metal with a molar mass of 27.0 g/mol 3C is a transition metal with an atomic number of 39 3D is a metal represented by the symbol In 3E is the first metal of the lanthanide series 3F is found in group 13 and period 6 of the modern periodic table 4A is the element whose common isotopic form is the basis of the atomic mass unit 4B has a second ionization energy of 1577 kj/mol 4C is located between scandium and vanadium on the modern periodic table 4D is represented by the symbol Zr 4E derived its symbol from the Latin word stannum 4F is a very dense metal with an atomic mass of g/mol 4G is the second member of the actinide series 5A is the most abundant element in the atmosphere 5B has 3 electrons in its 3p sublevel 5C is a byproduct of fossil fuel oxidation and represented by the symbol V 5D is a period 4 nonmetal known since E is a member of both group 5 and period 5 of the modern periodic table 5F was originally called stibium 5G is located in period 6 of the modern periodic table, below niobium 5H is a metal with atomic number 83 6A is the most abundant element in the Earth s crust 6B is a member of group 16, known during the time of the Roman Empire 6C is the first member of group 6 on the modern periodic table 6D is represented by the symbol Se 6E has 42 protons within its nucleus 6F is a halogen whose crystals sublime 6G was originally called wolfram 6H is the fourth member of the actinide series 7A has an atomic radius of nm 7B is represented by the symbol Mn 7C forms a diatomic gas with the molar mass of g/mol 7D has a molar mass of g/mol 8A is a group 8 metal known during the time of the Roman Empire 8B is an element named after the German word for Satan 8C has an average atomic mass of 58.9 g/mol 8D is located between iron and osmium on the modern periodic table 8E has a nuclear charge of +45 8F has an atomic mass of g/mol 8G has 114 neutrons within its nucleus 8H is named after the Latin word for rainbow, iris 8I is an inert metal often used in electrodes and has an atomic number of 78

29 29 QUESTIONS: 1. Examine the placement of electron configurations on your periodic table. What relationship can be seen in an element s placement within a group and its electron configuration? 2. Examine the graph of atomic radius vs. atomic number. What trend do you observe as you go across a period? 3. Which group appears to have members with the largest atomic radii? Which group has the smallest? 4. Examine the graph of ionization energy vs. atomic number. What trend do you observe as you go across a period? 5. No members of group 18 of the modern periodic table can be found on Mendeleev s classification chart. Suggest a reason for their absence. 6. What factor may account for the observed trend in atomic radii as one proceeds across a period? CONCLUSIONS: Make sure to answer your objectives, summarize your results, and state what you have learned.

30 30 LAB EXPERIMENT 11 OBJECTIVES: POLAR AND NON-POLAR SOLUTES AND SOLVENTS 1. To determine the type of solvent that generally dissolves ionic compounds. 2. To determine the type of solvent that generally dissolves polar covalent compounds. 3. To determine the type of solvent that generally dissolves non-polar covalent compounds. 4. To investigate the effect of adding a polar liquid solute to a non-polar liquid solvent. PREDICTIONS: Make sure you read through the lab and understand the objectives before making your predictions. MATERIALS: Apparatus Reagents Tweezers sodium chloride crystals Test tube rack sucrose crystals 6 test tubes (13 mm x 100 mm) iodine crystals 6 stoppers for test tubes 3 unknown solid solutes paint thinner glycerol PROCEDURES: PART 1: Solubility Tests on Known Solutes 1. Obtain 6 clean, dry test tubes and place them in a test-tube rack so that you have two rows of 3 test tubes each. 2. Half fill one set of 3 test tubes with room temperature water and half fill the other 3 test tubes with paint thinner. 3. Into the first pair of test tubes (one with water and one with paint thinner) add enough salt crystals to just cover the bottom of the test tubes. 4. Stopper the test tubes and hold it on with your thumb while you shake each test tube to see if each substance dissolves. When you are sure no further changes are taking place, record your observations in Table 1. Dispose of the chemicals in the waste disposal jar. 5. Repeat Steps 2, 3 and 4 with crystals of sugar in each. Dispose of the chemicals in the waste disposal jar. 6. Repeat Steps 2, 3 and 4 with a single crystal of iodine in each. Dispose of the chemicals in the waste disposal jar.

31 31 PART 2: Solubility Tests on Unknown Solutes 1. Repeat Part 1 of this experiment, but refer to Table 2 and test the solubilities of unknown solutes A, B, and C. PART 3: Mixing Two Liquids 1. Fill a clean test tube one quarter full with paint thinner, then add twice as much water to the same test tube. 2. Stopper the test tube and shake the test tube. Examine what happens to the liquids after shaking, and record your observations in Table Add one iodine crystal to the test tube, stopper and shake. Make a sketch of the test tube and its contents in your observations. 4. In a second test tube, fill it one quarter full with glycerol. It is a polar liquid. Add twice as much water, stopper and shake. Record your observations in Table 3 and make a sketch in your observations. 5. Dispose of all chemicals in the waste disposal jar. Clean up all your apparatus and make sure you wash your hands with soap and water before leaving the lab. OBSERVATIONS AND DATA: Table 1: Known Solutes with Known Solvents Solvents Water (Polar Covalent) Paint Thinner (Non Polar Covalent) Salt (NaCl Ionic) Solutes Sugar (C 12 H 22 O 11 Polar Covalent) Iodine (Non Polar Covalent) Table 2: Unknown Solutes with Known Solvents Solutes Solvents A B C Water (Polar Covalent) Paint Thinner (Non Polar Covalent)

32 32 Table 3: Liquid Combinations Combinations of Liquids Covalent Types Results Water and Paint Thinner Water and Glycerol QUESTIONS: 1. a. What general trend appears in Table 1 with regard to which type of solute dissolves in which type of solvent? b. This general solubility trend is sometimes expressed as Like dissolves like. Explain 2. a. Attempt to classify each of the unknown solutes from Part 2 as ionic, polar covalent, or non polar covalent. b. What problem do you encounter in making this classification? c. How could you overcome this problem? 3. Explain the terms miscible and immiscible and use these terms to explain the results from Part How did the addition of an iodine crystal help in identifying the layers of liquids in the water paint thinner combination? 5. Explain how many layers you would expect to see if water, paint thinner, and glycerol were combined in one test tube. 6. Explain which solvent from this experiment you would use to remove road salt stains from a pair of jeans. CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this lab. Don t forget any sources of error that may have occurred.

33 33 LAB EXPERIMENT 12 ACID-BASE TITRATION OBJECTIVES: 1. To titrate a hydrochloric acid solution of unknown concentration with standardized sodium hydroxide solution. 2. To utilize the titration data to calculate the molarity of the hydrochloric acid solution. PREDICTIONS: Make sure you read through the lab and understand the objectives before making your predictions. MATERIALS: Apparatus Reagents 250 ml Erlenmeyer flask standardized sodium hydroxide solution (approx M) volumetric pipette (10 ml) unknown hydrochloric acid solution burette phenolphthalein solution burette clamp and stand pipette bulb PROCEDURE: 1. Obtain about 50 ml of the unknown hydrochloric acid solution and about 50 ml of the standardized sodium hydroxide solution. Your instructor will provide you with the exact molarity of the NaOH. Record this value in Table Using the pipette bulb and volumetric pipette, rinse out the pipette and then transfer 10.0 ml of the HCl solution into a 250 ml Erlenmeyer flask. 3. Add 3 drops of phenolphthalein solution. 4. Rinse out a burette with some of the standardized NaOH solution and drain it, keeping the tip filled with the NaOH. Refill the burette and record your initial volume in Table Gradually add the NaOH into the Erlenmeyer flask containing the HCl. Swirl the flask continuously. Continue adding NaOH, noting any changes in the flask. 6. As the equivalence point is reached, a pinkish colour will appear. As you approach this point, the colour will take longer to disappear as you swirl it. Add the NaOH drop by drop until the colour no longer disappears. Record the final volume of the burette in Table 1. The most accurate reading is one in which the solution is a very faint pink.

34 34 7. Repeat steps 2 through 6 after cleaning out your Erlenmeyer flask (the contents can be poured down the sink with lots of water). Continue repeating until you get two trials that are within ± 0.10 ml of one another. 8. Clean up all your apparatus and make sure you wash your hands with soap and water before leaving the lab. OBSERVATIONS AND DATA: Table 1: Volume of NaOH Needed to Neutralize ml of Unknown HCl Molarity of NaOH = Trial 1 Trial 2 Trial 3 (if necessary) Initial Volume of NaOH Final Volume of NaOH Volume of NaOH used Average Volume of NaOH QUESTIONS: 1. Calculate moles of NaOH from the average volume used, and the given molarity. 2. Calculate moles of HCl present originally. 3. Calculate the molarity of the HCl solution. 4. What was the reason for cleaning out the burette with NaOH before starting the titration? 5. What is the concentration of a NaOH solution when it requires 30.0 ml of 0.50 M HCl to neutralize 50.0 ml of the NaOH? 6. What is the concentration of acetic acid (HCH 3 COO) when 32.5 ml of 0.56 M NaOH is required to neutralize 15.0 ml of the acid? 7. When 6.25 g of oxalic acid (H 2 C 2 O 4 ) are placed in water containing phenol red, the suspension is yellow. What is the molarity of the KOH if 32.2 ml of KOH is required to change the colour to red? 8. A 5.0 g tablet of TUMS (Mg(OH) 2 ) neutralizes 450 ml of stomach acid (HCl). What is the molarity of the stomach acid? CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don t forget any sources of error that may have occurred.

35 35 LAB EXPERIMENT 13 MOLECULAR MODEL BUILDING OBJECTIVES: 1. To represent molecular structures with electron dot and structural (line) diagrams. 2. To construct molecular models of simple substances. 3. To construct molecular models illustrating the different types of isomers. PREDICTIONS: Make sure you read through the lab and understand the objectives before making your predictions. MATERIALS: Apparatus Molecular model kit PROCEDURE: 1. Construct models for each of the following alkanes. In your lab, draw the structural diagram and the electron dot diagrams for each of these molecules. a. methane, CH 4 b. ethane, C 2 H 6 c. propane, C 3 H 8 2. Construct models for all structural isomers for each of the following compounds. Draw the structural diagram, and name each one. a. butane, C 4 H 10 (there are 2 of them) b. pentane, C 5 H 12 (there are 3 of them) c. hexane, C 6 H 14 (there are 5 of them) d. cyclohexane, C 6 H 12 (there is only 1) e. methylhexane, C 7 H 16 (there are 2 of them) 3. Construct models for all structural isomers of butene, C 4 H 8. Draw the structural diagrams for each (there are 4 of them) and then name each one. Note that 2 of them are geometric isomers. 4. Construct models for all structural isomers of propyne (there is only 1), C 3 H 4, and butyne (there are 2 of them), C 4 H 6. Draw the structural diagrams for each, and then name each one. 5. Construct models for all structural isomers of n-hexanol, C 6 H 13 OH. Draw the structural diagrams for each (there are 3 of them) and then name each one.

36 36 6. Construct models for the following structural isomers. Compare their structures and the placement of the oxygen atom. Draw the structural diagram for each isomer. a. ethanol, C 2 H 5 OH b. dimethyl ether, CH 3 OCH 3 QUESTIONS: 1. Can cyclopentane (C 5 H 10 ) be considered an isomer of pentane? Explain. 2. What is the difference between a cis and trans isomer? 3. Explain why there are such a huge number of organic compounds. 4. How many isomers are possible for the following structures? a. methane b. ethane c. propane d. butane e. pentane f. hexane CONCLUSIONS: Make sure to answer your objectives, summarize your results and tell what you have learned in this experiment. Don t forget any sources of error that may have occurred.