Chemistry 125 Laboratory Manual

Transcription

1 Chemistry 125 Laboratory Manual This packet contains the laboratory experiments for Chemistry 125. Please refer to your class syllabus for the schedule of dates as to when the experiments will be performed. Please bring booklet with you to each laboratory class session. Table of Contents Scientific Measurement...3 Density...13 Separation of a Mixture...25 Determination of a Chemical Formula...31 Solubility Rules and Writing Formulas...37 Acid Base Titration...43 Cycle of Copper Reactions...51 Hess' Law...57 Determination of Iron by Titration with Permanganate...63 Conductivity of Solutions...71 Activity Series...77 Models...87 Boyle's Law: Pressure-Volume Relationships in Gases...95 Molar Volume of Hydrogen...99 Appendix A: Statistical Treatment of Data Appendix B: Equipment Hunt Chemistry 125 Laboratory Manual Page 1

2 Revision Dec 2003 Chemistry 125 Laboratory Manual Page 2

3 Scientific Measurement Goals: Use the balance to make quantitative measurements Use precision and accuracy in measurements Use average deviation and percent relative average deviation with measurements Read rulers, graduated cylinders, and burets correctly Background: In this experiment we will gain an understanding as to how to make measurements using a balance. The measurements made will be used to explore the limitations of measurements and then how the data obtained from measurements can be presented in a useful way. We will mass 5 pennies made before 1983 and then 5 pennies made after We will then calculate the mean, the average deviation and the relative average deviation of the two sets of pennies. If your instructor chooses, the data from the entire class can be used and the standard deviation calculated. In addition to using the balance, you will also learn how to make proper measurements using a ruler, graduated cylinder, and buret. You will use these pieces of laboratory equipment throughout the semester. Therefore, an early understanding of how to take proper measurements with this laboratory equipment is a crucial skill in the laboratory. The assessment of experimental error: Every measurement involves some measurement error. Errors can be classified in two ways: determinate error or indeterminate error. Determinate errors, also called systematic errors, are errors in your methods, equipment/materials, personal judgments, and simple mistakes. Determinate errors can be eliminated with practice and proper attention to detail. Indeterminate errors or random errors cannot be controlled or eliminated. Thus, there will always be some degree of error in measurements. Accuracy: The error in a measurement is the difference between the true value of the quantity measured and the measured value. Error calculated this way is known as absolute error. The smaller the absolute error, the closer the measured value is to the true value and the more accurate the result. Accuracy is a measure of the correctness of the measurement. Frequently, we wish to compare several measurement errors of different quantities. In these cases, it is more useful to use relative error. Relative error is defined as the error divided by the accepted value and multiplied by 100. This is usually known as percent error: (experimental - accepted) % error = accepted x 100 Sometimes we do not have or know an accepted value. In those cases, percent error cannot be calculated. We do not have an accepted value for this experiment. Therefore, we will not be calculating percent error. Chemistry 125 Laboratory Manual Page 3

4 Precision: The key to significance in experimental measurement is repetition. Only with repeated measurements can an experimenter have some confidence in the significance of the measurements. Only if a measured quantity can be reproduced repeatedly can the experimenter have that confidence. Precision is a quantitative measure of the reproducibility of experimental measurements. It is how well repeated measurements of the same quantity agree with one another. Precision is frequently measured in terms of the average deviation, which is determined by the following process: 1. From a series of measurements (three or more) determine the average value. _ x i x = mean = n 2. For each measurement determine its deviation from that average value. _ x i - x average deviation = n where, n= the number of data points and x i = the individual data points Note: the summations in the formulae above go from i = 1 to n. 3. Determine the average of the deviations without regard to sign. Example: Four measurements of the concentration of an unknown acid by titration were made. The following results were obtained: M, M, M, M. Compute the average value and the average deviation. Solution: 1. Compute the average value by summing the four measurements and dividing by four. ( )/4 = rounded off to Deviation from the average value of : ( ),.0004 ( ),.0002 ( ),.0002 ( ). Sum of the deviation = /4= Since this deviation represents an uncertainty in the measurements, the molarity of the unknown acid is not precisely M but ranges from M to M and should, therefore, be reported with the average deviation included, that is ± M. Further, to make the measurements of precision more useful, the average deviations are put on a percentage basis by determining the relative average deviation (r.a.d.). (See Appendix A, pg. 103 for more information). This is the average deviation divided by the average value and multiplied by 100%. average deviation r.a.d. = x 100 mean Chemistry 125 Laboratory Manual Page 4

5 In the above example, the relative average deviation is: x 100% = 0.3% The precision of an experiment varies with the technique and/or the apparatus used. A number of variables built into the method or design of the experiment can affect its precision. Chemistry 125 Laboratory Manual Page 5

6 Procedures: Part I. Pennies Mass Measurement Pennies minted in the United States in 1981 and earlier years are significantly different from pennies minted in 1983 or later. Is it possible to tell the difference between these two kinds of pennies by weighing them? 1. Put on your lab coat and your goggles. 2. Work in pairs. 3. One student will receive 5 pennies minted 1981 or before. The other student will receive 5 pennies minted 1983 or later. 4. Receive instructions from your instructor for the use of the pan balances in the laboratory. 5. Weigh and record the mass of each of the 5 pennies received. 6. Record the mass and the dates of the 5 pennies of your partner. 7. In your report calculate the mean of each set of pennies. 8. In your report calculate the average deviation 9. In your report calculate the relative average deviation Part II. Proper Measurements Using Lab Equipment WORK INDIVIDUALLY FOR THIS PART OF THE LAB!! Ruler 1. Obtain a ruler and examine the units and markings. Your ruler should look something like this: Suppose you wanted to measure the bold line above this ruler. Clearly the line is longer than 5 cm and shorter than 6 cm. Furthermore, the ruler is calibrated with lines each representing 0.1 cm. Counting the lines reveals the line to be somewhere between 5.8 cm and 5.9 cm. You must estimate between the lines. Always read one more digit than is marked. Final reading = 5.88 cm (The underlined 8 is the estimated digit.) 2. Measure the items listed under the Ruler section on the data page for this lab. Be sure to include proper units and make all readings to the correct number of decimals. Chemistry 125 Laboratory Manual Page 6

7 Graduated Cylinders 3. Obtain a 10 ml and 100 ml graduated cylinder. Examine the units and markings. Your graduated cylinders should look something like this (note: they are not drawn to scale): Suppose you want to measure the liquid shown in the 10 ml graduated cylinder. Notice the liquid is curved down. This is called the meniscus. Aqueous solutions will always curve down. Mercury is a liquid that actually curves upward. For purposes of this lab and this course, you will only be using aqueous solutions; therefore, always measure liquids from the bottom of the meniscus. Clearly the line is between 3 ml and 4 ml. The 10 ml graduated cylinder is calibrated with lines each representing 0.1 ml. Counting the lines reveals the line to be somewhere between 3.5 ml and 3.6 ml. You must estimate between the lines. Always read one more digit than is marked. Final reading = 3.53 ml (The underlined 3 is the estimated digit.) Reading the 100 ml graduated cylinder is accomplished in the same way; however, notice each of the markings represents 1 ml rather than 0.1 ml. A quick look at the graduated cylinder indicates the volume to be between 20 and 30 ml. Each line represents 1 ml, therefore, the line is clearly between 26 ml and 27 ml. You must estimate between the lines. Always read one more digit than is marked. Final reading = 27.2 ml (The underlined 2 is the estimated digit.) Notice, the number of decimal place you can read depends on how the lines are calibrated on a particular piece of equipment. 4. Practice using the 10 ml and 100 ml graduated cylinders located in your desk. THEN read the graduated cylinders in the back of the room to the correct number of decimal places. Chemistry 125 Laboratory Manual Page 7

8 Buret (Pronounced byur ette ) 5. Burets will be used several times during the semester. The calibrations on a buret are the same as a 10 ml graduated cylinder. Therefore, the buret will be read to the same number of decimal places. However, the numbers on the buret are reversed. The zero is at the top and 50 ml is at the bottom. Your instructor should comment briefly on the reason for this and the uses of the buret. The reason for the reversed numbers will become clearer as you use the buret during the semester. The buret should look like this: The line is clearly between 2.9 and 3.0 ml. You must estimate between the lines. Always read one more digit than is marked. Final Reading = 2.97 ml (The underlined 7 is the estimated digit.) 6. Read the burets located at the back of the room. Chemistry 125 Laboratory Manual Page 8

9 Scientific Measurements Data and Calculations Name Sec PART I. Your data from the weighing of five pennies: Date of the penny Mass of the penny (x i ) Average (mean) mass of five pennies Average deviation: Average mass ± average deviation: Percent relative average deviation: Your partner's data from the weighing of five pennies: Date of the penny Mass of the penny (x i ) Average (mean) mass of five pennies Average deviation: Average mass ± average deviation: Percent relative average deviation: Sample Calculation: Chemistry 125 Laboratory Manual Page 9

10 Part II. Proper Measurements Using Lab Equipment Ruler Measure the following items to the correct number of decimals using a ruler. Include proper units. 1. The diameter of a 250 ml beaker = 2. The diameter of a penny = 3. The length and width of this piece of paper: Length = Width = Graduated Cylinders Measure the graduated cylinders located in the back of the room to the correct number of decimals. Include proper units. 4. Volume of Graduated Cylinder A = 5. Volume of Graduated Cylinder B = 6. Volume of Graduated Cylinder C = 7. Volume of Graduated Cylinder D = Buret Measure the burets located in the back of the room to the correct number of decimals. Include proper units. 8. Volume of Buret A = 9. Volume of Buret B = 10. Volume of Buret C = Chemistry 125 Laboratory Manual Page 10

11 Questions: 1. Does the data you collected during Part I. indicate that there is a significant difference between the masses of pennies minted in 1981 and earlier versus those minted in 1983 and later years? YES or NO (Circle One Answer) 2. Fill in the appropriate data and complete the question below. Pennies Dated 1981 and earlier Average ± average deviation = Pennies Dated 1983 and later Average ± average deviation = Comparing the averages and using the average deviations, support your answer from question #1. 3. Below you will find 3 different rulers each with different calibrations. Measure the bold line as accurately as possible, to the correct number of decimals, and with proper units. Length = Length = Length = Chemistry 125 Laboratory Manual Page 11

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13 Density Experiment Goals: Learn to use equipment (graduated cylinder, balance). Read equipment properly (estimate between lines). Learn to use significant figures. Calculate % error. Prelab Report on page 19. Turn in when you come to class! Background: In your study of chemistry, you will often have to make quantitative measurements and carry out calculations based on your data. If a correct or accepted result is known, you can compare your results to this correct result. This lab is an introduction to the techniques and mathematics you will need to use throughout the course. Significant Figures: The following rules can be used to determine the appropriate number of significant digits in a measurement or of a number as the result of a calculation. The significant figures in a measured quantity include all of the certain digits plus the last digit is considered to be the uncertain digit (not known precisely, guessed). The uncertainty in the last digit in a number is assumed to be ±1 unless otherwise noted. For example, suppose you measured the mass of a penny and reported it to be g. The first three digits are assumed to be exactly correct, but the last digit is assumed to have an uncertainty of at least ±1. In other words the mass could be as low as g or as high as g. So the value of the measurement is reported as (2.496±0.001) g or simply g. Often, further calculations must be done after collecting the data. Significant figures should also be used properly in these calculations. Your final answer should reflect the accuracy of the original measurement, rather than the number of digits displayed by the calculator. In general, the number of significant figures in the least accurate measurement limits the number of significant figures in the answer. Density: Mass The density of a substance is defined as the mass per unit volume: D = Volume The units of density are commonly g/ml or g/cm 3 for most solid and liquid substances (a milliliter is the same as a cubic centimeter). Density is a physical property, which has a characteristic value for pure substance. Some representative values are: Lead 11.3 g/ml Ethanol 0.79 g/ml Silver 10.5 g/ml Water (4 o C) 1.00 g/ml Aluminum 2.70 g/ml Mercury 13.6 g/ml Ice 0.92 g/ml Cork 0.26 g/ml Chemistry 125 Laboratory Manual Page 13

14 Measuring the density of an unknown substance can help you determine what the substance is. For example, if an unknown piece of shiny gray metal were found to have a density of 2.70 g/ml, you would strongly suspect that the metal is aluminum. Percent Error: When a correct or accepted value is known, you can determine if your measurements give an answer in close agreement with the accepted value. A common way to do this is to report the percent error. Refer to appendix A for discussion of percent error. Procedures: To determine density, mass and volume must be measured. Mass can be determined directly from a balance. Volume can be measured directly (measuring dimensions and using geometric formulas) or indirectly (water displacement). Taring the Balance: When weighing several small objects or a powder, it is impossible to place it directly on the balance pan. It is easiest to place the objects in a container, and have the balance subtract out the weight of the container. This process is called taring. Place the empty container on the balance and press the bar on the front. The display should now read zero (there will be some small fluctuations due to air currents). Remove the container from the balance (the balance will read a negative mass after the container is removed) and place the material to be weighed in the container. It is better to do this away from the balance to protect the balance from spills. The container with the material should then be placed back on the balance. The readout will give the weight of the material alone (the weight of the container has been subtracted out by the balance). Volume by Water Displacement: A common technique for determining volume is by using water displacement. This is often used when the volume of the solid is difficult to calculate from geometry or difficult to measure, such as an irregular shape. The technique involves measuring the amount of water in a container (graduated cylinder), adding the solid and remeasuring the water level. The water level will increase due to the presence of the solid. The difference between the levels (before and after adding the solid) is the volume of the solid. A. Regular Solid - Direct Measurement 1. Obtain a regularly shaped (cylinder or cube) piece of metal. 2. Determine the mass of the metal piece using the balance (record all the digits from the balance, do not round off). 3. Using a ruler, measure the dimensions (height, diameter or length, width, height). Estimate between the lines on the ruler (if each line is 0.1 cm, estimate to the nearest 0.01 cm). Record the measurements. Chemistry 125 Laboratory Manual Page 14

15 4. Show your calculations for the volume of the metal piece. Remember to use significant figures properly in recording your answer and give the appropriate units. 5. Show your calculations for the density of your metal piece. Remember to use significant figures properly in recording your answer and give the appropriate units. 6. When you are finished, return the metal piece and ruler to the appropriate place on the cart. 7. Using the CRC handbook available in the lab, look up the densities of several metals and alloys. Choose the one which seems to fit the appearance and density of your sample the best and calculate the % error. B. Regular Solid - Volume by Water Displacement 1. Obtain a cylinder made of metal. 2. Determine the mass of the metal cylinder (record all the digits from the balance, do not round off). 3. Pour some tap water into your 100 ml graduated cylinder. The amount should be enough so that it would completely cover the solid if it were submerged. 4. Record the level of the water. Estimate between the lines on the graduated cylinder - if the lines represent 1 ml, you should estimate to 0.1 ml. 5. Tip the cylinder and gently slide the metal piece down the side (if it is dropped in, the glass will crack). Record the level of the water with the metal included. 6. The volume of the metal is the difference between your two measurements. Show your calculations for the volume of the metal. Remember to use significant figures properly in recording your answer and give the appropriate units. 7. Show your calculations for the density of your metal piece. Remember to use significant figures properly in recording your answer and give the appropriate units. 8. When you are finished, dry off the metal piece and return it to the appropriate place on the cart. 9. Using the CRC handbook available in the lab, look up the densities of several metals and alloys. Choose the one which seems to fit the appearance and density of your sample the best and calculate the % error. Chemistry 125 Laboratory Manual Page 15

16 C. Irregular Solid 1. Obtain some marble chips (about chips). 2. Determine the total mass of your marble chips. 3. Place an empty plastic weigh boat on the balance pan and press the bar on the front of the balance to tare it. Remove the weigh boat and place the marble chips in it. Place the weigh boat with the marble chips back on the balance. The readout should give the mass of the marble chips alone. 4. Record the mass (record all the digits from the balance, do not round off). 5. The volume will be determined by the technique of water displacement (see above). Pour some tap water into your 100 ml graduated cylinder. The amount should be enough so that it would completely cover the solid if it were submerged. 6. Record the level of the water. Estimate between the lines on the graduated cylinder - if the lines represent 1 ml, you should estimate to 0.1 ml. 7. Gently drop the marble chips into the graduated cylinder. Make sure there are no air bubbles between the chips. Record the level of the water with the marble included. 8. The volume of the marble is the difference between your two measurements. Show your calculations for the volume of the marble chips. Remember to use significant figures properly in recording your answer and give the appropriate units. 9. Show your calculations for the density of marble. Remember to use significant figures properly in recording your answer and give the appropriate units. 10. When you are finished, return the marble chips to the appropriate container on the cart. (They are not placed back in the original container with the unused chips). 11. Using the CRC handbook, look up the density of marble. Calculate the % error. D. Liquid 1. Obtain a small amount (about 10 ml or less) of salt water in a beaker. 2. Clearly, a liquid cannot be poured onto the balance in order to be weighed. It must be in a container. The technique of taring the balance (see above) will be used to determine the weight of the liquid without the container. 3. Place an empty 100 ml graduated cylinder on the balance pan. Press the bar on the front of the balance. Remove the graduated cylinder and pour between 5 and 10 ml of the liquid in it (do this away from the balance so that any spill will not damage the balance). Place the graduated cylinder with the liquid back on the balance. The readout should give the mass of the liquid alone. Chemistry 125 Laboratory Manual Page 16

17 4. Record the mass (record all the digits from the balance, do not round off). Note that as the water evaporates, the mass will decrease somewhat. 5. The volume can be read directly from the graduated cylinder. Record the level of the liquid. Estimate between the lines on the graduated cylinder - if the lines represent 1 ml, you should estimate to 0.1 ml. 6. Show your calculations for the density of salt water. Remember to use significant figures properly in recording your answer and give the appropriate units. 7. In order to determine the density more precisely, the measurements will be repeated with a different piece of equipment. 8. Place a clean, empty 10 ml graduated cylinder on the balance and tare. Pour between 5 and 10 ml of the salt water into the graduated cylinder (again, do this away from the balance). Place the graduated cylinder with the liquid back on the balance. The readout should give the mass of the liquid alone. 9. Record the mass (record all the digits from the balance, do not round off). Note that as the water evaporates, the mass will decrease somewhat. 10. The volume can be read directly from the graduated cylinder. Record the level of the liquid. Estimate between the lines on the graduated cylinder - if the lines represent 0.1 ml, you should estimate to 0.01 ml. 11. Show your calculations for the density of salt water. Remember to use significant figures properly in recording your answer and give the appropriate units. 12. When you are finished, pour the salt water down the drain. Rinse out your graduated cylinder thoroughly so that the salt water does not dry inside and leave a residue. 13. Find out the accepted value of the density of the salt water (it may be written on the label of the container or may be given by your instructor). Calculate the % error. Chemistry 125 Laboratory Manual Page 17

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19 Density Experiment Prelab Exercise Name Sec How many significant figures are in each of the following measurements? ml g cm g L 1.20 x 10 3 m ml 1.05 x 10-4 g L x 10 6 cm Carry out the following calculations and round off the answer to the correct number of significant figures: g g = g g = 5.2 ml ml = g g = g / 2.0 ml = 10.0 g / 3 ml = g / 11.0 ml = g / 0.1 ml = ( ) x 25.5 = ( )/2.70 = Chemistry 125 Laboratory Manual Page 19

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21 Density Experiment Data and Calculations Name Record measurements as accurately as you can. Include units with all measurements and calculations. Report all calculated answers to the correct number of significant figures. Show work for calculations. Sec A. Regular Solid - Direct Measurement Mass Description of object Dimensions: (Be sure to include the units) Length Show Calculations Width Height Diameter Volume Density Accepted value Percent error B. Regular Solid - Volume by Water Displacement Mass Volume of water alone Volume of water and metal Volume of metal Density Reference Source for Accepted Value of Density: (Include the name, edition and page number) Identification of Metal Accepted Value for Density Percent error Chemistry 125 Laboratory Manual Page 21

22 C. Irregular Solid Show Calculations Mass Volume of water alone Volume of water and marble chips Volume of marble Density of Marble Accepted Value for Density Percent error D. Liquid: 100 ml graduated cylinder Mass of liquid Volume of salt water Density Percent Error 10 ml graduated cylinder Mass of liquid Volume of salt water Density Percent Error Accepted Value for Density Chemistry 125 Laboratory Manual Page 22

23 General Questions: 1. What is the limiting factor in determining the number of significant figures in your measurement of density - mass or volume? If you wished to determine the density to more significant figures, would you buy a better balance (to measure mass more precisely) or try to improve the measurement of volume (for example, by using calipers instead of a ruler, or a more finely calibrated graduated cylinder)? 2. In determining the density of the liquid, you weighed the liquid in a graduated cylinder and then determined the volume by reading the graduated cylinder. Suppose that you weighed the liquid in a beaker, and then transferred it to a graduated cylinder in order to determine the volume. Which approach would be more likely to give an accurate result? (Hint: What errors could be introduced by transferring liquid from one container to another?) 3. Suppose, after you determine the density of your solid, you discover that it is hollow. How would this affect your value for the measured volume (would your measured volume be higher or lower than the actual)? How would this affect your value for the density (would the calculated value be higher or lower than the actual value)? Chemistry 125 Laboratory Manual Page 23

24 4. Suppose that some air bubbles are trapped in between the marble chips when you measure the volume by water displacement. How would this affect your value for the volume (would your measured volume be higher or lower than the actual)? How would this affect your value for the density (would the calculated value be higher or lower than the actual value)? 5. Can water displacement be used to calculate the density of a cork? If so, how? If not, why not? 6. How does the density of 4 ml of water compare to the density of 2 ml of water? (Same, twice as much, half as much?) Explain your answer. Chemistry 125 Laboratory Manual Page 24

25 Separation of a Mixture Goals: Devise a scheme of separation based upon physical properties of substances. Write up a laboratory report outlining the procedures and equipment used. Cleanly separate a mixture and recover as much as possible of the components. Background: Physical properties of substances are those which can be measured without changing the identity of the substance. Examples of physical properties are density, color, melting point, magnetism and solubility. These properties can be significantly different from one substance to another and can be made use of in order to separate substances. For example, in doing laundry, clothes can be separated by color. Chemical properties, which involve the change of the substance, are not as useful since the substance will not be recovered in the same pure form. For example, in doing laundry, cotton materials are known to be flammable (they react with oxygen). However, setting the clothes on fire will destroy them. Materials and Procedures: You will be given a mixture of five components. The mixture contains copper(ii) sulfate, sand, iron filings, foam pieces, and boric acid. The following physical properties are known about the substances: copper(ii) sulfate: soluble in both hot and cold water; blue solid; density: 2.3 g/ml boric acid: soluble in hot but not cold water; white solid; density: 1.4 g/ml sand: not soluble in water; brown, tan and white solid; density is greater than 1.0 g/ml iron filings: not soluble in water; black solid; density: 7.9 g/ml foam pieces: not soluble in water; white solid; density is less than 1.0 g/ml You are to design an experimental procedure that will separate the five components into their natural states (all should be recovered as solids). You will be given a test tube containing one gram of the mixture. After separation, place the components into small baggies labeled with your name and the name of the component. Chemistry 125 Laboratory Manual Page 25

26 A flow chart is often helpful in clarifying the procedures to be used. For example: Dirty Clothes (White, delicate and permanent press) Examine each piece, separate by color Dirty White Clothes Place in washing machine Add detergent Add Bleach Set machine to hot water wash Run washing machine Dirty colored clothes (Delicate and permanent press) Read label of each piece Separate delicate from permanent press Clean white clothes Dirty colored permanent press clothes Dirty colored delicate clothes Clean colored permanent press clothes Place in washing machine Add detergent Set machine to warm water wash Run washing machine Place in washing machine Add special detergent Set machine to delicate Run washing machine Clean colored delicate clothes The flowchart describes the procedures used at each step, and gives a clear diagram of the separation. At each step, it is clear which component is being separated and which still remains in the mixture. Chemistry 125 Laboratory Manual Page 26

27 Separation of a Mixture Name Sec Flow Chart: Design a flow chart of your separation. At this point, check with your instructor before proceeding. Chemistry 125 Laboratory Manual Page 27

28 Procedure: List the steps in numerical order for your procedure. This should be a verbal explanation of your flow chart. Describe the procedures you are using. Correct terms should be used for equipment. The procedure should be written clearly enough that another person could understand and follow the instructions. Chemistry 125 Laboratory Manual Page 28

29 Equipment: List the equipment that you will need to accomplish the separation. Conclusions: 1. Discuss the sources of error in your experimental design. How might you do things differently to eliminate these errors? 2. List five specific examples from your experiment, which demonstrate how physical (not chemical) properties can be used to separate a mixture. Chemistry 125 Laboratory Manual Page 29

30 3. Using the information given, indicate how a mixture of the compounds might be separated: Use a flowchart for your answers. PbCl 2 and AgCl PbCl 2 : white solid, density 5.85 g/ml, soluble in hot water but not in cold water AgCl: white solid, density 5.56 g/ml, not soluble in hot or cold water Ice, methanol (CH 3 OH), and Styrofoam ice: density 0.9 g/ml, melting point 0 C methanol: density 0.79 g/ml, boiling point 65 C, melting point 97.8 C, soluble in water Styrofoam: white solid, density 0.47 g/ml, not soluble in water Fe 2 S 3 and Fe(NO 3 ) 3 Fe 2 S 3 : green solid, density 4.3 g/ml, decomposes upon heating, not soluble in water Fe(NO 3 ) 3 : violet solid, density 1.68 g/ml, decomposes upon heating, soluble in water Chemistry 125 Laboratory Manual Page 30

31 Determination of a Chemical Formula Goals: Carry out a chemical reaction and isolate the product. Determine the empirical formula of a compound. Prelab Report on page 33. Turn in when you come to class! Background: A formula of a chemical compound is a description of what chemical elements are present in the compound. When the elements are nonmetals, the compound is usually made up of molecules. The molecules are atoms covalently bonded together, such as H 2 O, water, or C 6 H 6, benzene. When metals and nonmetals combine to make a compound, there are usually positive and negative ions present. The formula is the simplest ratio of ions. There are no individual molecules in which a certain cation is bonded to a certain anion. For example, in sodium chloride each sodium ion is surrounded by six chloride ions and each chloride ion is surrounded by six sodium ions. The formula is NaCl because the smallest whole number ratio is sodium and chloride is one -to-one. In calcium chloride the formula is CaCl 2 because for each calcium ion there are two chloride ions present. Because atoms are so small, we cannot discover the formula of each chemical substance by counting atoms or ions. Instead, larger quantities of chemical substances are used. The masses of elements in chemical compounds are found. The analysis can be done in two ways: (1) The pure compounds are broken down and the amount of each element is found; (2) Elements are combined to give a compound and the amounts of the elements that react are found. Using known molar masses, gram amounts of the elements are converted to moles. The smallest whole number ratio of moles of each element is the chemical formula. For example: A compound contains 2.65 g aluminum and 2.35 g oxygen. Convert grams to moles: 2.65 g Al (1 mol Al/27.0 g Al) = mol Al 2.35 g O (1 mol O/16.0 g O) = mol O Find the mole ratio: mol Al/ = Al mol O/ = O Convert, if necessary, to a whole number ratio: 1.000Al x 2 = 2 Al O x 2 = O The formula for the compound is Al 2 O 3. Chemistry 125 Laboratory Manual Page 31

32 Procedure: You will react zinc with iodine to form zinc iodide. The mass of each reactant is measured, as well as the mass of the product and any leftover reactants. By calculating the mole Zn : mole I ratio, you will determine the chemical formula of zinc iodide. Methanol is poisonous and flammable. Avoid skin contact and avoid breathing the vapors. (Review the procedure for tarring a balance in the Density experiment). 1. Weigh out approximately 1.0 g of zinc metal. 2. Weigh out approximately 1.0 g of iodine solid. DO NOT LET ANY IODINE FALL ON THE BALANCE! It reacts with the metal and damages the balance. 3. Weigh a clean dry 125 ml Erlenmeyer flask and record the mass. 4. Add the zinc to the flask, reweigh, and record the mass. 5. Add the iodine to the flask, reweigh, and record the mass. 6. Add approximately 25 ml of methanol to the flask. Cover the flask with foil. Place the flask on a hot plate in the fume hood and adjust the temperature to heat the methanol to boiling. (Try about 3 on the dial of the hot plate for this temperature.) Swirl the flask occasionally. The reaction should be complete in approximately minutes. There will still be some particles of zinc metal left in the flask, but the iodine should have disappeared. 7. Weigh a clean 250 ml beaker and record the mass. 8. Carefully pour the liquid from the reaction flask into the beaker. Do not let any zinc leave the flask and fall into the beaker. To make certain that all the zinc iodide is transferred to the beaker, add 5 ml of methanol to the zinc left in the flask and swirl. Pour the liquid into the beaker with the zinc iodide liquid - again, do not let any particles of zinc leave the flask. Repeat the washing with another 5 ml of methanol. 9. Place the flask with the unreacted zinc back on the hot plate and allow it to dry. When it is dry, let the flask cool and then weigh the flask with the remaining zinc metal. 10. Place the beaker with the zinc iodide and methanol liquid on the hot plate. Adjust the temperature so that the methanol evaporates, leaving the zinc iodide in the beaker. Make sure all the methanol evaporates and the product is dry. Allow the beaker to cool and then weigh the beaker with the zinc iodide product. Chemistry 125 Laboratory Manual Page 32

33 Determination of a Chemical Formula Pre-lab Exercise Name Sec Lead forms two different compounds with chloride - one is an oily yellow liquid (compound A), the other is a white solid (compound B). A chemist adds a sample of compound A to a test tube with a mass of g (empty weight). On re-weighing the test tube with compound A he found the mass to be g. On heating, he found that some chlorine gas escaped and compound A was converted to compound B. He found that the mass of the test tube and compound B was g. Compound B was then heated in the presence of hydrogen gas, reducing the lead to the metallic elemental form and driving off the chloride as HCl gas. The weight of the remaining lead and the test tube was found to be g. What is the mass of compound A? What is the mass of compound B? What is the mass of the lead? The lead left in the test tube at the end was present in both compounds A and B, so they both contain the same amount of lead. The rest of the mass must be due to chlorine. What is the mass of chlorine in compound A? What is the mass of chlorine in compound B? The mole to mole ratio will give the simplest formula of the compounds. How many moles of lead were in the compounds? How many moles of chlorine were in compound A? How many moles of chlorine were in compound B? What is the formula of compound A? What is the formula of compound B? Show your calculations: Chemistry 125 Laboratory Manual Page 33

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35 Determination of a Chemical Formula Data and Calculations Name Before Reaction: Sec Show Calculations: Mass of empty Erlenmeyer flask Mass of flask and Zn Mass of flask, Zn and I After Reaction: Mass of empty beaker Mass of beaker and zinc iodide Mass of flask with leftover Zn Calculations: Before reaction: Mass of Zn Mass of I Moles of I After Reaction: Mass of zinc iodide Mass of leftover Zn Mass of Zn that reacted Moles of Zn that reacted Calculate the mole I reacted: mole Zn reacted ratio. According to your data, what is the simplest formula for the compound zinc iodide? Chemistry 125 Laboratory Manual Page 35

36 A way to check your technique in this experiment is to see if mass is conserved. Before Reaction: Mass of Zn before reaction Mass of I before reaction Total mass before reaction After Reaction: Mass of zinc iodide Mass of leftover Zn Total mass after reaction Questions: 1. In the procedures, why dissolve the iodine in methanol? Why not mix the two solids and allow them to react? 2. Suppose some leftover zinc metal was transferred to the product. How would this affect your calculations, and how would the formula of the product be affected? (Would the formula have a higher I:Zn ratio or lower?) 3. Suppose a non-reactive impurity was present in the zinc at the start of the reaction, and remained in the flask with the leftover zinc. How would this affect your calculations, and how would the formula of the product be affected? 4. What changes did you observe that indicated a chemical reaction? 5. What physical properties of zinc metal and zinc iodide did you use in order to separate the mixture? Chemistry 125 Laboratory Manual Page 36

37 Solubility Rules and Writing Formulas Goals: Accurately record observations of the result of mixing two salt solutions. Write net ionic equations for the reaction of salt solutions. Identify an unknown by reaction with various salt solutions. Background: When a salt dissolves in water to form an aqueous solution, it exists in solution as individual ions. For example, sodium phosphate (Na 3 PO 4 ) does not exist as an aqueous solution of Na 3 PO 4 molecules, but rather as separate Na + (sodium ions) and PO 4 3 (phosphate ions) surrounded by water. Na 3 PO 4 (s) H 2O 3Na + (aq) + PO 4 3 (aq) We can write the equation for the dissolving of silver nitrate as well: AgNO 3 (s) H 2O Ag + (aq) + NO 3 (aq) If two solutions, each containing a soluble salt, are mixed, then the ions from one solution are free to mix with the ions from the other. When silver nitrate and sodium phosphate are mixed, the species present before a reaction takes place are: Na + (aq) + PO 4 3 (aq) + Ag + (aq) + NO 3 (aq) We can mix a solution of aqueous sodium phosphate and aqueous silver nitrate. The products are aqueous sodium nitrate and solid silver phosphate. We get the complete balanced equation below: Na 3 PO 4 (aq) + 3AgNO 3 (aq) Ag 3 PO 4 (s) + 3NaNO 3 (aq) Na + NO 3 - Na + Na + Ag + Ag + NO - PO PO 3-4 Na + Na + NO - 3 Ag + Na + NO - 3 PO 3- Ag + 4 Rewriting the equation using the ionic species in the solution before and after they are mixed, we get the equation below. Silver phosphate is not a soluble salt and will precipitate (the solid that forms) from the solution. The aqueous solution that remains contains sodium ions and nitrate ions. This equation is sometimes called the complete ionic equation. 3Na + (aq) + PO 4 3 (aq) + 3Ag + (aq) + 3NO 3 (aq) 3Na + (aq) + 3NO 3 (aq) + Ag 3 PO 4 (s) NO 3 - Na + Na + - NO 3 Ag 3 PO 4 Ag 3 PO 4 If the resulting salts are soluble in the products side of the equation (such as sodium nitrate above) then no reaction has really taken place between these ions and they are the spectator ions. Thus the product of this chemical reaction is silver phosphate. The solution that remains at the end of the reaction contains the spectator ions Na + and NO 3. Because these spectators do not participate in the chemical reaction that has taken place, we do not need to write them in the overall reaction. Since the reactants are present as aqueous ions and the product is a solid, a reaction is said to occur. We can cancel the spectator ions out of the equation and rewrite the chemical reaction as the net ionic equation as shown below. 3Ag + (aq) + PO 4 3 (aq) Ag3 PO 4 (s) Chemistry 125 Laboratory Manual Page 37

38 The insoluble silver phosphate is seen as a fine yellowish solid, which will slowly settle to the bottom of the container. One method of establishing the solubility of salts is to mix solutions of known composition and observe the resulting mixture. If a precipitate forms, then at least one of the salts produced is insoluble. We could say, based upon our example, that silver phosphate is insoluble. But does one of the two ions have more influence on the reaction than the other? Are all silver salts soluble except the phosphate? Or are all phosphates soluble except the silver? In this experiment we will examine the results obtained when nine anions (in solution as sodium salts) are mixed with nine cations (in solution as nitrate salts). Procedure: Obtain one sheet of acetate with a 10x10 grid and obtain one set of the nine anions and one set of the nine cations. (There should be one set of each per table.) Each of the salts is present as 0.1 M aqueous solutions in bottles with a screw top and dropper. 1. Put an acetate sheet over a piece of white paper, and then put one drop of each of the anions (NaCl, NaI, etc.) in each of the squares of the proper column on the acetate sheet. 2. Place one drop of each of the cations (AgNO 3, KNO 3, etc.) in each square of its proper row on the acetate sheet. Drop the solutions onto the sheet carefully. DO NOT touch the dropper to the sheet or the solutions in the bottles will become contaminated. If you make a mistake, blot up the liquid with a KimWipe. 3. Most of the drops will remain clear, indicating that the salts formed are soluble and no reactions occurred; but several solutions change color and form a precipitate. Often it helps to see the precipitate if the white sheet is carefully slipped from under the acetate. Note on your report sheet any color changes and/or precipitate formation. There may be changes due to further reactions with vapors or light as time passes. On your report sheet, record only the reactions that occur. Leave the squares empty where no reaction occurs. 4. Write equations for each mixture which resulted in the formation of a precipitate on the report sheet. 5. Develop a set of statements or rules about the solubilities of ions based upon your observations (e.g., all nitrates are soluble). Write these solubility rules on a separate sheet of paper. Group your observations together so that you have no more than four solubility rules. 6. Obtain an unknown from your instructor. Record the unknown letter on your report. Your unknown contains both an anion and a cation. Following the procedure you used previously, place a drop of the unknown in each box in the column and the row labeled unknown. React the unknown with a drop of each anion and cation in the appropriate square. Compare the reactions of the unknown with those of the known anions and cations, and determine which column of known anions matches the column of the unknown, and which row of cations matches the row of the unknown. Note that the entire row or column must match, not just some of the squares. It may not be possible to completely identify the unknown, but you should be able to narrow down the possibilities (e.g., your unknown may contain the iodide ion as the anion and either the sodium or the ammonium ion as the cation. The cation cannot be narrowed down any further based upon the reactions you have observed in this lab). Chemistry 125 Laboratory Manual Page 38

39 Name Solubility Rules Sec anions cations Solubility Table PO 4 3 C2 H 3 O 2 NO 3 Cl SO 4 2 I OH CO3 2 Unknown anion K + Ag + Ba 2+ Na + Ni 2+ NH 4 + Mg 2+ Al 3+ Zn 2+ Unknown cation Chemistry 125 Laboratory Manual Page 39

40 Data: In the space below, write balanced net ionic equations for the formation of any precipitates you observed in the experiment. Unknown Letter Which cation was in the unknown? Which anion was in the unknown? Write the name and formula of the ionic compound in the unknown (e.g., if your unknown contained Ba 2+ and I -, you would write Barium Iodide, BaI 2 ). Briefly describe how you reached a decision about your unknown. You may use equations, but don t limit your explanation to them. Chemistry 125 Laboratory Manual Page 40

41 Questions: 1. Suppose you were asked to prepare solid silver chloride from a solution of silver nitrate. Describe the chemical reaction you would carry out (include any necessary reagents) and how you would separate the solid product from the other materials. 2. List your solubility rules on this page. Write rules that you can draw from this laboratory experiment. Do not simply restate the rules that you have learned in lecture. In other words, derive rules that are based on your experimental results. Group your observations together so that you have no more than four solubility rules. Chemistry 125 Laboratory Manual Page 41

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43 Acid - Base Titration Goals: Learn to read a buret. Learn how to accurately carry out a titration. Calculate the molarity of an unknown acid or base from titration data. Prelab Report on page 47. Turn in when you come to class! Background: In this experiment, you will be carrying out the following acid-base neutralization reaction: HCl(aq) + NaOH(aq) H 2 O(l) + NaCl(aq) In this reaction, one mole of acid will react exactly with one mole of base. This reaction can be used to determine the concentration of an unknown acid (or base) by reacting the unknown with a base (or acid) of known concentration. In this experiment you will titrate several different samples of HCl(aq) with NaOH(aq). Titration is a laboratory technique where a carefully measured amount of one solution is added from a buret to a fixed amount of another solution in an Erlenmeyer flask. In today's experiment, the buret is filled with NaOH(aq). The HCl(aq) samples are placed in Erlenmeyer flasks along with a color indicator. Sodium hydroxide solution is added to the Erlenmeyer flask until a color change occurs. This is called the endpoint. The following information is required to determine the concentration of the unknown acid: 1. The concentration of the base - this is determined through a careful titration with a precisely measured amount of acid (a process called standardization). You will be using a base, which has been previously standardized so that the molarity of the base is accurately known. 2. The volume of base used in the reaction - you will use a buret to measure the amount of base used in the reaction. Knowing the molarity of the base and the volume of base used, it is possible to calculate the number of moles of base used in the reaction. According to the titration reaction shown above, the number of moles of base added at the endpoint is equal to the number of moles present in the sample initially. 3. The volume of acid used in the reaction - you will use a buret to measure an accurate volume of the unknown acid. The molarity of the acid solution can then be calculated from the moles of acid (determined in the titration) and the volume of acid. 4. A means of determining the endpoint (the point at which just enough base has been added to neutralize the acid) - a substance called an indicator will be added to the reaction mixture. An indicator is a molecule which changes color as the concentration of acid (H + ) changes. The indicator we will use is called phenolphthalein. It is colorless in acid solutions, and pink in alkaline solutions. As NaOH is added to your HCl sample, hydrogen ions from the acid are removed from solution, and the indicator will change color (becoming pink). The endpoint is reached when the indicator just changes color. Near the endpoint the color change may occur very suddenly when a single drop of base is added to the flask. Chemistry 125 Laboratory Manual Page 43