GENERAL CHEMISTRY PHS 1015 Laboratory Manual INDEPENDENCE COMMUNITY COLLEGE Department of Chemistry 2014-2015 Table of Contents Introduction: Laboratory Policies ...............................................................................................2 Experiment 1: Measurement & Significant Figures .................................................................3 Experiment 2: Chemical & Physical Properties ......................................................................23 Experiment 3: Identification Unknown Ion (Flame Test) .......................................................31 Experiment 4: Endothermic & Exothermic Chemical Reactions .........................................40 Experiment 5: Conductivity ......................................................................................................44 Experiment 6: Lewis Dot Structure & Molecular Models .....................................................49 Experiment 7: Stoichiometry .....................................................................................................62 Experiment 8: Alum Synthesis .................................................................................................66 Experiment 9: Alum Analysis ...................................................................................................72 Experiment 10: Identification Unknown Solution (Solubility) ...............................................79 Experiment 11: Acid-Base Titration (Handout) .....................................................................90 Experiment 12: Household Acids & Bases ..............................................................................91 Experiment 13: Boyles Law & Guy-Lussac ................................................................................ Boyle’s Law .......................................................................................................................96 Guy-Lussac ......................................................................................................................100 Experiment 14: Molar Volume (Handout) ............................................................................105 Experiment 15: Heat of Fusion ...............................................................................................106 Experiment 16: Evaporation & Intermolecular Interactions ..............................................110 1|Page Introduction: Laboratory Policies: Laboratory part of science course are integral in the understanding of science. Science is about exploring, observing, and interpreting experiments and chemistry is no different. 40% of the grade for this course is based on the laboratory part of the class. There are 16 experiments to be completed. The schedule of the experiments is shown in the syllabus and rarely changes. Be prepared to conduct any experiment scheduled. There are no make-up days for missed laboratory experiments. If you are going to miss a laboratory experiment, you can plan to come into another class conducting the same experiment. You need to email the professor if you are going to miss an experiment. Proper laboratory attire is required, if you do not have proper laboratory attire you will be asked to leave until you have proper attire. If you do not return to class and do not come to another class time to make up the experiment, your will receive a zero (0) for the laboratory experiment (both results and question sections) Grading: The laboratory experiments will be broken into two parts. You will receive a maximum of 15 points for the results part of the lab and a maximum of 10 points for the questions. Some experiments have pre- and post- laboratory question and some have only post- laboratory question. The only parts due are the results and any questions (only parts you write on). The introduction and procedure portion of the laboratory book are for you to keep. These sections will come in handy when you do the laboratory practical. The laboratory experiments are due the next class date, late experiments will not be accepted without a valid excuse. An email in advance is required to turn anything in late. 2|Page Experiment 1: Measurement & Significant Figures Purpose: What are some of the techniques used to make scientific measurements in the laboratory? How does one determine the uncertainty in any measurement? What is meant by “significant figures”? Introduction: Conducting experiments or investigations are an important activity in all science classes. These experiments or investigations require measurements to be recorded and analyzed. Every measurement contains some amount of uncertainty due to the accuracy (or inaccuracy) in the device or technique used to make the measurement. Some of the instruments or techniques used to make measurements have large uncertainties and some have small uncertainties. Recognizing how to record the measurements and determining the uncertainty of each measurement is your goal for this experiment. Measuring Mass: The measurement of mass is conducted using a balance, because historically the mass of the sample was determined by balancing it against a standard or known mass. Figure 1 & 2 are examples of different scales or balances used to measure the mass of an object or sample. Figure 2 contains examples of the digital balances used in many chemistry courses across the country. Figure 1. Triple beam Balance. Figure 2. Example of spring-balance & electronic digital balances. Here are some guidelines for operating the electronic balance to successfully obtain a mass of any sample: 3|Page 1. Be certain the balance has been “zeroed” (meaning the digital readout is “0”) before you place any sample on the balance. 2. NEVER weigh chemicals directly on the balance pan. A suitable container or weighing paper should always be used. Often, you will weigh the sample container empty (to get an initial weight) and then weigh the contain containing the sample or chemical. Subtract the two measurements to determine the mass of the sample or chemical. 3. Be certain that air currents are not effecting your measurements. Someone walking past your balance or bumping the table that the balances are sitting on can effect a measurement. When making measurements to 0.001 g, one should always close the door or lid to the balance. 4. NEVER measure the mass of hot or warm objects, since the temperature difference will change the density of the air surrounding the balance and give an inaccurate measurement. 5. Record your measurement, to the proper number of significant figures. The balances used measure to the 0.001 g or the 0.01 g. 6. Once your measurements have been made, close the door to the balance (if necessary), and be sure the balance registers zero. Be sure to clean the balance if any chemical have been spilled. 7. Be gentle with the balances, these are sensitive, delicate instruments and like an person, responds best when treated properly. When recording a mass, the digit furthest to the right of the number recorded is the uncertainty of the measurement. For example, a mass of a penny was measured to be 1.57 g. The uncertainty of the measurement is 1.57 ± 0.01 g. The uncertainty recognizes the + 0.01 g or -0.01 g in the measurement. The uncertainty or last significant figure is always recognized as the last digit of the measurement. The number of significant figures is defined as the number of digits recorded in the measurement. If you recorded 1.57 g as the mass of the penny, then the number of significant figures is 3 significant figures. Measuring Volume: Volume is another common measurement made in the laboratory. The most common type of sample to measure the volume is a sample in the liquid state of matter. There are many different devices that can be used to measure the volume of a sample, depending on the accuracy required in your measurement. Typically volume is measured in mL but the SI unit for volume is L. It is helpful to know: 1 mL = 1 cc = 1 cm3 A. Beakers & Erlenmeyer Flasks: 4|Page Using a beakers and Erlenmeyer Flasks to measure volume is the least accurate of all the devices used to measure volume. The measurement can be off by as much as 10%. For a 250 mL beaker, the measurement could be ± 25 mL in error. Figure 3 is an example of breakers and Figure 4 is an example of Erlenmeyer flasks. Figure 3. 400 mL Breaker Figure 4. Various sizes of Erlenmeyer Flasks B. Graduated Cylinder: Using a graduated cylinder to measure volume somewhat more precise, the error is less than 1%. Most of the volume measurements conducted in Chemistry are conducted with graduated cylinders. Examples of graduated cylinders are shown in Figure 5. Figure 5. Various sizes of Graduated Cylinders. 5|Page Depending on the size of the graduated cylinder, the uncertainty can be from ± 0.1 mL to ± 0.01 mL. C. Pipets & Volumetric Flasks: Pipets and volumetric flasks are calibrated to contain specific volumes of liquid. When filled to the proper level, the volume contained is accurate to ± 0.01 mL. When using the pipet, proper technique must be utilized for accurate measurements. Examples of volumetric flasks are shown in Figure 6 and an example of a pipet is shown in Figure 7. The following steps make for the proper use of a pipet: 1. The pipet must be clean and dried. When the pipet is draining, there should be no drops left on the pipet walls. 2. Always use a pipet bulb or filler, NO mouth pipetting (pulling liquid into the pipet by sucking on the pipet like a straw). 3. Equilibrate the pipet by drawing some of the liquid into the pipet, remove the bulb or filler. Tilt the pipet so it is parallel to the floor and coat the inside wall of the pipet with the solution. Allow the solution to drain out the tip of the pipet. 4. Use the pipet bulb or filler to draw the liquid into the pipet. Continue to draw the liquid into the pipet until the meniscus is 1 – 2 cm above the calibration mark. Remove the pipet bulb or filler and place a finger or thumb over the opening of the stem. Let the liquid flow out of the pipet, into a waste beaker until the level of the meniscus is touching the top of the calibration mark. Place the tip of the pipet into the receiving flask, with the tip of the pipet touching the wall of the flask. The pipet needs to be straight up and down so the receiving flask will be tilted. Allow the liquid to drain from the pipet while the tip continues to touch the wall of the flask. Once the liquid has drained, separate the pipet from the receiving flask. There is a small amount of liquid in the tip of the pipet, DO NOT blow this out into the receiving flask. When properly used, the pipet is designed to deliver the exact volume of liquid (± 0.01 mL) into the receiving flask. Figure 6. Various sizes of Volumetric Flasks. Figure 7. A Volumetric Pipet. 6|Page D. Burets: A buret (Figure 8) is used to add precisely measured volumes of liquid. This is different from using a pipet to deliver the precisely measured volume of liquid in that unlike with a pipet the volume of liquid can vary depending on the amount required. When using a pipet the volume of liquid is dependent on the volume of the pipet, the buret has a valve that can be opened and closed (Figure 9) to allow the user to add the liquid until a visual change has been achieved. Figure 8. A Buret. Figure 9. Buret valve (stopcock) open & closed. As when using the pipet, working with the buret requires a particular technique to determine the volume of the sample to the ± 0.01 mL. The following procedure should be helpful: 1. Equilibrate the pipet using the same procedure to equilibrate the pipet. Be sure to open the stopcock to allow some of the solution to drain through the stopcock and tip. When the solution has been removed from the buret, add enough solution to fill the buret past the 0.00 mL mark. 2. Open the stopcock and tap the end with the stopcock to shake loose and air bubbles in the stopcock/tip of the buret. Close the stopcock and determine the initial volume. If the meniscus is still above the 0.00 mL mark, drain enough of the solution so the meniscus is touching the top of the 0.00 mL mark. If the meniscus is below the 0.00 mL mark, read the volume as the initial volume. Be sure to view the meniscus at eye level as shown in Figure 10. An example of an initial and final volume measurement is shown in Figure 11. 7|Page Figure 10. Proper method to read the meniscus of the buret. Figure 11. The initial reading and final reading of a measurement using a buret. 3. Notice in the measurement with the buret, the hundredth place is an estimate. For the initial reading, starting above the meniscus, the measurement is between the 9 and 10. So this means the measurement will be starting with 9 mL. Counting the graduations between 9 and 10, one views the meniscus to be between the 6th and 7th mark. This leads 8|Page to initial measurement to start 9.6 mL. The meniscus is closer to the 0.6 mark than the 0.7 mark, estimating the meniscus one can say the meniscus is about 0.03 and therefore the initial measurement of the volume is 9.63 mL. Read the final reading to the hundredth place and verify the final reading to be 24.16 mL. To determine the volume dispensed, subtract the initial reading from the final reading to get the volume used. This laboratory experiment correlates with the section on measurements & mathematics in your lecture book. Students are expected to read the introduction and appropriate lecture material to answer the Pre-Lab Questions. 9|Page Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Pre-Lab Questions: 1. What is the theoretical value for the density of water? __________________________________ 2. In one of the experiments students are asked to determine the density of 3 solid objects. Describe the method you can use to determine the volume of the solid. You might want to discuss this with your laboratory partners. ________________________________________________________________________ ________________________________________________________________________ ________________________________________________________________________ ________________________________________________________________________ ________________________________________________________________________ ________________________________________________________________________ ________________________________________________________________________ ________________________________________________________________________ ________________________________________________________________________ ______________________________________________________ 3. When using a buret to measure the volume of a liquid, the measurement should be made to the nearest ________ mL. 4. The quantity of mass of 5 pennies was measured to be 12.54876 g. Which number in the measurement has an uncertainty? __________________________________________________ 5. How many significant figures are contained in the measurement of the 5 pennies? ___________ 6. What is the average mass of each penny? Be sure to consider the uncertainty in the measurement and the number of significant figures. ______________________________________________ 7. What is the uncertainty of each of the following measurements and how many significant figures are contained in each of the measurements? a. 840.3 mL ± _______mL _______ Significant Figures b. 0.0486 kg ± ______kg _______ Significant Figures 10 | P a g e c. 4.184 L ± ______L d. 5.000 mm ± ______mm e. 760 mm Hg ± ______ mm Hg _______ Significant Figures _______ Significant Figures _______ Significant Figures 8. Round each of the measurements in problem 4 to contain the number of significant figures indicated in the parentheses. a. (3) ______ mL b. (2) ______ kg c. (1) ______ L d. (3) ______ mm e. (2) ______ mm Hg 8. Carry out the following calculations and write the result using the correct number of significant figures. All the numbers should be treated as measured quantities. a. 0.847 x 2.10 = ______ b. 79.06/0.21 = ______ c. 17.2 x 1.000 = ______ d. 1.002/84.01 = ______ e. 0.748 + 12.02 + 0.0001 = ______ f. 12.54 – (12.874 + 12.0) = ______ 11 | P a g e Procedure: A. Distance measurements in cm & mm. Using the ruler in your lab drawer, measure the length and width of shapes A, B, C, & D found on page 10. Record the measurements in centimeters (cm) and millimeters (mm) on the Data & Observations Sheet. Calculate the area and parameter of the shapes using the correct number of significant figures. B. Mass measurements. Measure the mass of the three objects you receive from your instructor. Press the tare button to zero the balance and place a container to hold the sample on the balance pan. Record the mass of the empty sample container on the Data & Observations Sheet. Using the crucible tongs, place the sample into the sample container. Record the mass of the sample and sample container. Calculate the mass of the sample by subtracting the mass of the empty sample container from the mass of the sample container with the sample. The difference is the mass of the sample. Repeat for the remaining samples. Record the mass using the balance measuring to ± 0.01 g and ± 0.001 g on the Data Sheet. C. Volume measurements. Measure a volume of water using a 100 mL beaker, 50 mL graduated cylinder, and the 10 mL graduated cylinder. Weigh the empty 100 mL beaker, 50 mL graduated cylinder, and the 10 mL graduated cylinder and record the weight of each on the Data & Observations Sheet. Using a Beral pipet, add 100 drops of water to each container and measure the volume. Remember each of the containers will have a different uncertainty. Weigh the containers with the water in them; determine the mass of the water in each of the containers. The density of an object or substance is the mass of the object or sample, and the volume that object or sample occupies. Based on the data collected, determine the density of water. Each of the containers used to determine the volume will have a different volume with different uncertainty or significant figures. Determine the percent error using the formula: Theoretical value – Actual value x 100% = Percent Error Theoretical value The theoretical value for the density of water is 1.000 g/mL. D. Determine the Density of Solids. Using the mass measurements from Part B, determine the density of the 3 objects and using Table 1, determine the unknown substances. You will need to determine the volume of the object as accurately as possible so to get the most number of significant figures for your 12 | P a g e calculations. Notice the number of significant figures used in Table 1 and plan your experiment accordingly. Table 1. Density of Possible Unknown Solids Element Chemical Symbol Density (g/mL) Element Chemical Symbol Density (g/mL) Zinc Zn 7.140 Copper Cu 8.960 Tin Sn 7.265 Iron Fe 7.874 Lead Pb 11.35 Aluminum Al 2.702 Nickel Ni 8.908 Chromium Cr 7.190 Silver Ag 10.49 Gold Au 19.32 E. Density Determination of Unknown Liquid. Record the unknown identification code of the liquid unknown. Mark 3-50 mL beakers #1, #2, #3 and measure the mass of each empty beaker. Using the buret, determine the volume of the unknown liquids. Determine the mass of the each of the volumes of unknown liquid samples. Using Table 2, determine the unknown composition of the liquid. Table 2. Density of Possible Unknown Liquids Unknown Solution Density (g/mL) Unknown Solution Density (g/mL) Saturated NaCl 1.200 90% Isopropyl Alcohol 0.812 Saturated Sucrose 1.343 70% Isopropyl Alcohol 0.845 Vinegar 1.101 50% Isopropyl Alcohol 0.898 13 | P a g e Print this page! B A C D 14 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Data & Observations: A. Distance measurements in cm & mm. Table 3. Measurement of shapes in cm. Shape Measured Length (cm) Measured Width (cm) A B C D Table 4. Parameter of shapes in cm. Shape Calculated Parameter (unrounded) (cm) Calculated Parameter (rounded) (cm) A B C D 15 | P a g e Table 5. Area of shapes in cm2. Shape Calculated Area (unrounded) (cm2) Calculated Area (rounded) (cm2) A B C D B. Mass of Object Measurements. Identification code of Object 1 __________ Identification code of Object 2 __________ Identification code of Object 3 __________ Mass of Objects using 0.01 g Electronic Balance Mass of Object 1 and container __________ g Mass of container empty _______________ g Mass of Object 1 _____________________ g Mass of Object 2 and container __________ g Mass of container empty _______________ g Mass of Object 2 _____________________ g 16 | P a g e Mass of Object 3 and container __________ g Mass of container empty _______________ g Mass of Object 3 _____________________ g Mass of Objects using 0.001 g Electronic Balance Mass of Object 1 and container __________ g Mass of container empty _______________ g Mass of Object 1 _____________________ g Mass of Object 2 and container __________ g Mass of container empty _______________ g Mass of Object 2 _____________________ g Mass of Object 3 and container __________ g Mass of container empty _______________ g Mass of Object 3 _____________________ g C. Volume Measurements. Mass & Volume Measurements, Density Determination using 100 mL Beaker Mass of 100 mL beaker with water __________ g Mass of 100 mL beaker empty ______________ g Mass of water in 100 mL beaker ____________ g 17 | P a g e Volume Measurement using 100 mL beaker __________ mL Density of water using 100 mL beaker ______________g/mL Percent Error of Density of water using 100 mL beaker __________ % Mass & Volume Measurements, Density Determination using 50 mL Graduated Cylinder Mass of 50 mL Graduated Cylinder with water __________ g Mass of 50 mL Graduated Cylinder empty ______________ g Mass of water in 50 mL Graduated Cylinder ____________ g Volume Measurement using 50 mL Graduated Cylinder __________ mL Density of water using 50 mL Graduated Cylinder ______________g/mL Percent Error of Density of water using 50 mL Graduated Cylinder __________ % Mass & Volume Measurements, Density Determination using 10 mL Graduated Cylinder Mass of 10 mL Graduated Cylinder with water __________ g Mass of 10 mL Graduated Cylinder empty ______________ g Mass of water in 10 mL Graduated Cylinder ____________ g Volume Measurement using 10 mL Graduated Cylinder __________ mL Density of water using 10 mL Graduated Cylinder ______________g/mL Percent Error of Density of water using 10 mL Graduated Cylinder __________ % 18 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ D. Determine the Density of Objects. Density of Object 1 Identification code of Object 1 ___________ Mass of Object 1 __________ g Volume of Object 1 __________ mL Density of Object 1 __________ g/mL Element Object 1 __________ Chemical Symbol Object 1 __________ Density of Object 2 Identification code of Object 2 ___________ Mass of Object 2 __________ g Volume of Object 2 __________ mL Density of Object 2 __________ g/mL Element Object 2 __________ Chemical Symbol Object 2 __________ Density of Object 3 Identification code of Object 3 ___________ Mass of Object 3 __________ g Volume of Object 3 __________ mL Density of Object 3 __________ g/mL Element Object 3 __________ Chemical Symbol Object 3 __________ 19 | P a g e E. Density of Unknown Liquid Density of Unknown Sample 1 Unknown Identification Code Sample _______ Mass of Sample 1 __________ g Volume of Sample 1 __________ mL Density of Sample 1 __________ g/mL Composition of Sample 1 __________ Density of Unknown Sample 2 Unknown Identification Code Sample _______ Mass of Sample 2 __________ g Volume of Sample 2 __________ mL Density of Sample 2 __________ g/mL Composition of Sample 2 __________ Density of Unknown Sample 3 Unknown Identification Code Sample _______ Mass of Sample 3 __________ g Volume of Sample 3 __________ mL Density of Sample 3 __________ g/mL Composition of Sample 3 __________ 20 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Post-Lab Questions: 1. Oil (0.784 g/mL) and vinegar (1.101 g/mL) are immiscible. If you vigorously shake a bottle of Italian salad dressing (primarily oil and vinegar with herbs and seasoning) and let the bottle sit on the counter for a few minutes, what will happen to the Italian dressing? Explain. 2. If you mistakenly did not tare the balance before you measured the mass of the graduated cylinder with water, would your calculation of the density of the water be greater or less than the actual density of water? You notice the balance has a reading of 0.03 g. Explain. 3. Were the volume measurements made with the 10 mL graduated cylinder more, less, or equally precise as the volume measurements made with the 50 mL graduated cylinder? Explain. 21 | P a g e 4. The density of water at 0 oC is 0.9999 g/mL, and the density of ice at 0 OC is 0.9168 g/mL. Based upon these densities, does ice float or sink in water? Explain. 5. Does water contract or expand when it freezes? Explain. 6. Suppose you have a 400 mL beaker with ice filled to the 250 mL mark (assume no air pockets). What is the volume of liquid water formed when the ice melts? Show your work. 7. The density of cod liver oil is 0.92 g mL-1, what is the volume occupied by a 1 lb bottle of cod liver oil? Show your work. 22 | P a g e Experiment 2: Chemical & Physical Properties Purpose: Can a student differentiate between chemical change/properties and physical change/properties? Introduction: The properties of matter can be divided into two classes, physical properties and chemical properties. Physical properties are those properties that can be observed or measure without any change to the chemical nature of the substance. The best way to think about a physical property is that a physical property describes the substance. The color, density, viscosity, solubility, melting point, and boiling point are all physical properties of a substance. Chemical properties describe a change in the chemical composition of the substance under specific conditions. Physical change is a change in the physical nature of the substance. Evaporation is the change of a substance from the liquid state of matter to the gas state of matter. The substance did not change in the chemical nature or composition, only in its state of matter. Chemical change is a change in the chemical composition of the substance. Fermentation is the process yeast use to convert sugar (glucose) in to ethyl alcohol, carbon dioxide, water, and energy. Evidence of a chemical change include: color change, gas production, formation of a precipitate (insoluble product), and release or absorption of heat. When heat is released from a chemical reaction, the reaction is termed exothermic (giving off heat) and when the chemical reaction absorbs heat, the chemical reaction is termed endothermic (absorbing heat). 23 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Pre-Lab Questions: 1. What are the three states of matter? a. Describe each state of matter on the microscopic level. b. Describe each state of matter on the marcoscopic level. 2. Identify each of the following as a chemical change/property or physical change/property. Explain your choice. a. an iron nail rusting b. a test tube shattering when dropped on the floor 24 | P a g e c. paper shredded in a paper shredder d. a silver fork losing its luster while sitting in the drawer e. heated copper metal leaving a black residue when heated f. sugar becoming a gel on a humid day 3. What is an observation? Give an example of an observation. 4. What is inference? Give an example of inference. 25 | P a g e Procedure: A. Investigating properties. Obtain the group of the five unknown coded, sealed vials, DO NOT open the vials. Based on the appearance (physical properties) of the material in each vial and using the CRC Handbook of Chemistry and Physics, determine which vial contains the carbon, iron, lead, sulfur, and aluminum. B. Chemical and physical properties of a substance. Obtain the vial containing the white, crystalline unknowns. Based on the appearance (physical properties) of the material in each vial and using the CRC Handbook of Chemistry and Physics, determine which vial contains the sodium chloride (table salt) and sucrose (table sugar). Determine if either or both the substances are soluble in water (dissolve in water). Place one of the white, crystalline unknowns into an evaporating dish and the other white crystalline unknown into another evaporating dish. Heat each evaporating dish with a Bunsen burner for 5 minutes. Record your observations and inferences. Hint: You might want to fill out the data table 2 with the physical properties from the CRC Handbook of Chemistry and Physics. C. Chemical and physical properties of a substance (Part 2). Obtain the vial containing the next four known substances. Record the physical appearance of each substance. Record the solubility in water (dissolve in water) of each of the substances. Place the paraffin, sulfur, malonic acid, and copper(II) sulfate pentahydrate into separate evaporating dishes, record the appearance before and after heating. 26 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Data & Observations: A. Investigating Properties. CRC Handbook of Chemistry and Physics Edition: ________ Table 1. Physical Properties & Identification of Five Unknowns. Unknown ID Code 1 2 3 4 5 Name of Unknown Physical State Molar Mass (g/mol) Density (g/mL) Melting Point (oC) Boiling Point (oC) Solubility in water at 0 oC, g/100 cc Element Name Chemical Symbol 27 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ B. Chemical and physical properties of a substance. Table 2. Physical Properties & Identification of Table Salt & Table Sugar. Unknown ID Code 1 2 Physical State Molar Mass (g/mol) Density (g/mL) Melting Point (oC) Boiling Point (oC) Solubility in water at 0 oC, g/100 cc Appearance before heating Appearance after heating Chemical Name Chemical Symbol 28 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ C. Chemical and physical properties of a substance (Part 2). Substance Physical Appearance Soluble in water? (Y/N) Appearance before heating Appearance after cooling Paraffin Sulfur Malonic Acid Copper(II) Sulfate pentahydrate 1. Description of the odor of heated malonic acid: 2. Observation of heated copper(II) sulfate pentahydrate before heating, after cooling, after adding a few drops of water on the sample: 3. What can be inferred about the copper(II) sulfate: 29 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Post-Lab Questions: 1. What is another observation or physical property that can be made in the identification of table salt & table sugar part of the lab? What would the inference be of this observation? 2. One warm summer evening, a group of students decide to go camping. For dinner they cut up beef, potatoes, and onions into cubes, wrap everything into packets using aluminum foil, and place the packets on the campfire. After 40 minutes, the contents of the aluminum foil packets were opened revealing a brown, well-cooked meal. For desert, the students ate crispy, golden brown, roasted marshmallows with graham crackers and chocolate. One of the students boiled water for hot cocoa. During the evening, the temperature dropped and the grass became damp with the evening dew. a. List the physical changes described in the story. b. List the chemical changes described in the story. 30 | P a g e Experiment 3: Identification Unknown Ion (Flame Test) Purpose: Using a flame test, can a student identify unknown ions? Introduction: Understanding the Bohr model of the atom is the basis for understanding the ability for one to identify unknown ions using a technique call a flame test. Each element emits unique colored bands combined to give a unique color to that element. In the Bohr model (Figure 1), the electrons exist in energy levels around the nucleus. When an electron absorbs energy, the electron moves to a higher energy level (further away from the nucleus). This situation can only occur for a limited time and the electron must return to a lower energy level. When the electron returns to a lower energy, the difference in the energies is released as a photon. Humans visualize the energy in a photon as visible light in the electromagnetic spectrum. The colored light seen corresponds to an energy difference between the electron in a higher energy level and the electron transitioning to a lower energy level (Figure 2). The colored light seen also corresponds to a specific wavelength. A wavelength (Figure 3) is defined as the distance from the peak to peak or trough to trough. Typical units for wavelength is the unit used to measure distance, nanometer, or another unit is the angstrom (Å). The wavelength is related to the number of wave’s passing over a point, in an amount of time. This is called the frequency, hertz, or cycles per second. The relationship between wavelength and frequency () is shown in equation 1. When the electron transitions from a higher energy level to a lower energy level, a photon is released. The photon corresponds to a specific amount of energy and related to a specific wavelength () by the equation 2. = c/ Equation 1 c = the speed of light (3.00 x 108 m/s) = frequency = wavelength E = h Equation 2 E = Energy h = Planck’s constant (6.626 x 10-34 J.s) 31 | P a g e Figure 1. Bohr Model of the atom. n = 1, n = 2, n = 3, are the different energy levels. Figure 2. Electron transition from higher energy level to lower energy level. 32 | P a g e Figure 3. Determining a wavelength measurement. 33 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Pre-Lab Questions: 1. Pretend your lab partner skipped class and you need to teach them how to light the Bunsen burner. Write out the procedure to light the Bunsen burner as you would explain the processes to your lab partner. 2. Using equations 1 & 2, rearrange equation 2 to solve for wavelength and substitute the rearranged equation 2 into equation 1. The new equation should show the relationship between the energy of the photon and the frequency of the photon. 3. What color flame indicates a “cool” flame? 4. What color flame indicates a “hot” flame? 34 | P a g e 5. Where is the flame the hottest? 6. What part of the Bunsen burner controls the gas flow? 7. What part of the Bunsen burner controls the air supply? 8. What is the method a student should use when the characteristic color of an ion is similar to another ion? 35 | P a g e Procedure: A. Flame test of known ions. Adjust the flame on the Bunsen burner to a very hot flame. Dip one end of a cotton swab into a solution containing a known ion. Place the tip of the cotton swab with the solution on it into the hot portion of the flame. You may need to repeat the flame test a couple of times to determine the color of the flame. Record the corresponding ion and the characteristic color of the flame emitted when the tip of the cotton swab was placed into the flame. The one side of the cotton swab should only be used to one solution. Do not dip the cotton swab into another solution! Use one cotton swab for each known solution. When you have completed the flame test for the known ions, place the cotton swabs into a beaker with some water to make sure you extinguish any hot embers on the cotton swab. B. Flame test of unknown ions. Each group is assigned two unknown solutions. Repeat the procedure to conduct the flame test of the known ion solutions. Record the unknown identification codes, the corresponding characteristic flame color, and the identity of the unknown ion. If you are uncertain of the identity of the unknown ion, conduct the flame test of the suspected known ion solution and the unknown ion solution at the same time. Sometimes the differences in the characteristic colors are very slight and the only way to correctly identify the unknown ion solution is to conduct the flame test side by side. 36 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Data & Observations: A. Flame test of known ion solutions. Table 1. Known ions and Characteristic color. Chemical Symbol Element Description of color Na K Li Sr Ba Cu Table 2. Unknown ions identity and Characteristic color. Unknown ID Code Description of Color Unknown Identity 37 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Post-Lab Questions: 1. Name the colors of light in the electromagnetic spectrum from highest energy. ______ ______ ______ ______ ______ ______ ______ 2. Red light has a longer wavelength than blue light. c. Which of these colors has the higher frequency? d. Which of these colors has the higher energy? 3. What are some of inaccuracies that may be involved when using the flame tests for identification purposes? Explain. 38 | P a g e 4. The alkali metals cesium & rubidium were can be identified based on their characteristic flame colors. Cesium was named after the sky and rubidium was named after the gem color. What characteristic colors of light do you think are emitted when these metals are heated in a flame? 5. When a glass rod is heated, the characteristic color of flame is a bright yellow flame. Based on the flame test experiment, what metal ion do you think is predominately found in glass? 6. What experiment would you conduct to verify the characteristic flame color was due to theh metal cation and not the anion? Explain. 39 | P a g e Experiment 4: Endothermic & Exothermic Chemical Reactions Purpose: Many chemical reactions give off energy. Chemical reactions that release energy are called exothermic reactions. Some chemical reactions absorb energy and are called endothermic reactions. You will study one exothermic and one endothermic reaction in this experiment. In Part I, you will study the reaction between citric acid solution and baking soda. An equation for the reaction is: H3C6H5O7(aq) + 3 NaHCO3(s) 3 CO2(g) + 3 H2O(aq) + Na3C6H5O7(aq) In Part II, you will study the reaction between magnesium metal and hydrochloric acid. An equation for this reaction is: Mg(s) + 2 HCl(aq) H2(g) + MgCl2(aq) Another objective of this experiment is for you to become familiar with using LabQuest. In this experiment, you will use the program to collect and display data as a graph or list, to examine your experimental data values on a graph, and to print graphs and data lists. In this experiment, you will study one exothermic and one endothermic reaction. Figure 1 Procedure: 1. Obtain and wear goggles. Part I Citric Acid plus Baking Soda 2. Connect the Temperature Probe to LabQuest. 40 | P a g e 3. Place a Styrofoam cup into a 250 mL beaker as shown in Figure 1. Measure out 30 mL of citric acid solution into the Styrofoam cup. Place the Temperature Probe into the citric acid solution. Note: The Temperature Probe must be in the citric acid solution for at least 30 seconds before doing Step 5. 4. Weigh out 10.0 g of solid baking soda on a piece of weighing paper. 5. You are now ready to begin collecting data. a. Record the temperature of the citric acid solution. b. Add the baking soda to the citric acid solution. Gently stir the solution with the Temperature Probe to ensure good mixing. c. Temperature readings (in °C) can also be monitored in a display box. d. Record the temperature of the solution when the temperature of the solution stabilizes. 6. Dispose of the reaction products as directed by your instructor. Rinse the Temperature Probe. Part II Hydrochloric Acid Plus Magnesium 7. Measure out 30 mL of HCl solution into the Styrofoam cup. Place the Temperature Probe into the HCl solution. Note: The Temperature Probe must be in the HCl solution for at least 30 seconds before doing Step 9. 8. Obtain a piece of magnesium metal from your instructor. 9. You are now ready to begin collecting data. a. Record the temperature of the citric acid solution. b. Add the magnesium strip to the HCl solution. Gently stir the solution with the Temperature Probe to ensure good mixing. c. Temperature readings (in °C) can also be monitored in a display box. d. Record the temperature of the solution when the temperature of the solution stabilizes. Caution: Do not breathe the vapors. 10. Dispose of the reaction products as directed by your instructor. Rinse the Temperature Probe. 41 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ DATA TABLE Part I Part II Final temperature, t2 °C °C Initial temperature, t1 °C °C Temperature change, t °C °C 42 | P a g e PROCESSING THE DATA 1. Calculate the temperature change, t, for each reaction by subtracting the initial temperature, t1, from the final temperature, t2 (t = t2 – t1). 2. Tell which reaction is exothermic. Explain. 3. Which reaction had a negative t value? Is the reaction endothermic or exothermic? Explain. 4. For any chemical reaction, describe three ways you could tell a chemical reaction was taking place. 5. Which reaction took place at a greater rate? Explain your answer. 43 | P a g e Experiment 5: Conductivity Purpose: Students investigate conductivity of aqueous solution using a number of dissolved solutes. What are the criteria for a solution’s ability to conduct electricity? What types of compounds, when dissolved in water, conduct electricity? Why are some compounds better at conducting electricity than others? Introduction: Electricity is the flow of negative charged particles. The ability for electricity to flow in an aqueous solution is measured by the conductivity of the solution. An aqueous solution that conducts electricity must have ions or be an ionic compound. When this type of compound dissolves in water, the ions separate into cations and anions. This is what allows the negative charged particles of electricity flow. Solutions that conduct electricity are called electrolytic solutions. Covalent compounds do not form ions and therefore cannot conduct electricity. These form non-electrolytic solutions and do not conduct electricity. The periodic table can be used to predict whether a compound is ionic or covalent and how much the compound will conduct electricity based on the location, of the elements, on the periodic table. This principle is called the “trends of the periodic table”. This experiment investigates this principle and the role of the periodic table in predicting the physical and chemical properties of the compounds. 44 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Pre-Lab Questions: 1. Define an atom. 2. Define an ion. 3. Define a polyatomic ion. 4. Do all compounds contain ions? Explain. (What other types of compounds are there?) 5. Define electrolytic solutions. 6. Define non-electrolytic solutions. 45 | P a g e Procedure: A. Taking Measurements with the Conductivity Probe. You should not need to calibrate the conductivity probe. Connect the conductivity probe to the LabQuest data collection device. To make a measurement, into 400 mL beaker (waste beaker), rinse the tip of the conductivity probe using the squirt bottle of distilled water and blot dry the tip of the probe using a Kimwipe (lab tissue) to remove any water droplets. Insert the tip of the conductivity probe into the sample to be tested. Important: Be sure the electrode surface is completely submerged in the liquid. While gently swirling the probe, wait for the reading to stabilize, this should take about 5 to 10 seconds. Rinse the end of the probe with the distilled water over the waste beaker before taking another measurement of storing the probe. Obtain a set of vial containing 0.01 M aqueous solutions for each group to test. Using the procedure above to take measurements using the conductivity probe, record the conductivity of each of the aqueous solutions. 46 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Data & Observations: A. Taking Measurements with the Conductivity Probe. Table 1. Conductivity of Aqueous Solutions. 0.01 M Solution Conductivity (mg/L) 0.01 M Solution Tap Water NaNO3 Distilled Water KNO3 LiCl Mg(NO3)2 NaCl Ca(NO3)2 KCl Sr(NO3)2 Sucrose Ba(NO3)2 Glucose MgCl2 Na2SO4 CaCl2 Na3PO4 SrCl2 K3PO4 BaCl2 Al(NO3)3 NiCl2 H2O2 CuCl2 CH3OH ZnCl2 C2H5OH MgSO4 C3H7OH CaSO4 Conductivity (mg/L) 47 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ 1. Which solutions were electrolytic? 2. Which solutions were non-electrolytic? 3. Using the periodic table, locate the elements of the compounds for the solutions that were electrolytic. Do you see a pattern for electrolytic solutions? Explain. 4. Using the periodic table, locate the elements of the compounds for the solutions that were non-electrolytic. Do you see a pattern for non-electrolytic solutions? Explain. 5. For the electrolytic solutions, describe if there is a pattern seen in the numerical readings? Explain the reason for these patterns. 48 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Post-Lab Questions: 1. How does the ionic and covalent bond differ? 2. Choose one of the electrolytic solutions and draw a view, at the molecular level, of the particles in the solution. 3. Choose one of the non-electrolytic solutions and draw a view, at the molecular level, of the particles in the solution. 49 | P a g e Experiment 6: Lewis Dot Structure & Molecular Models Purpose: Students investigate how to write Lewis Dot Structures and convert these structures into Electronic & Molecular geometric shapes. What is the significance of the electronic & molecular geometry shape on the physical & chemical properties of the compound? Introduction: Whether the chemical bonding be ionic or covalent in nature, the electron is the critical component to the bonding theory. The outermost electrons or valance electrons are the electrons on an atom that are chemically active and lead the element or compound having its physical & chemical properties. The Lewis Dot representation of the atom, molecule or ion depicts each of the valance electrons. Covalent bonds are bonds between atoms that share electrons. The ability to represent a molecule or polyatomic ion using the Lewis Dot Structure allow one to determine the special orientation of the electron cloud or electron density. The spatial orientation of the all the electron clouds around a central atom lead to the ability to determine the electronic and molecular geometries. The method used to determine the geometric shape of the molecules or polyatomic ions is called VSEPR (valance shell electron pair repulsion) model. The ability to represent the molecule or polyatomic ion using the Lewis Dot Structure will help determine molecules that have resonance and distinguish whether a molecule is polar or non-polar. 50 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Pre-Lab Questions: 1. What is meant by the “octet rule”? 2. What is meant by “electronegativity”? 3. Does a negative ion have more or less electrons than the neutral atom? Explain. 4. Does a positive ion have more or less electrons that the neutral atom? Explain. 5. What is the acronym “VSEPR” mean? 6. What is meant by “resonance”? 51 | P a g e Procedure: Complete the tables to show the Lewis Structures and Electronic & Molecular Geometry for the molecules and polyatomic ions, using the information from the textbook and the information given below. A. Writing the Lewis Dot representation for a monoatomic ion or atom. Determine the number of valance electrons for the monoatomic ion or atom. Write the chemical symbol for the monoatomic ion or atom. Imagine a box around the chemical symbol; place a dot for each valance electron around the chemical symbol. A box has four sides, so the first four electrons are unpaired, if there are more than four electrons, the electrons start getting paired. B. Writing the Lewis Dot representation for a compound or polyatomic ion. Determine the number of valance electrons for each of the atoms in the compound or polyatomic ion. The central atom tends to be the least electronegative atom but the central atom is NEVER hydrogen. The remaining atoms are placed around the central atom to form a skeleton referred to as a spider. Some general rules for forming the bonding and non-bonding pairs of electrons. 1. Valance electrons are the only electrons allowed to form bonds between atoms. 2. Atoms should initially be paired so one electron from each atom is shared to form a single bond between the two atoms. 3. Check for compliance with the octet rule. 4. Oxygen atoms will not bond together except in rare cases (O2, O3, peroxides) 5. Electrons should be paired, no individual electrons. 6. If the octet rule is not fulfilled, try forming multiple bonds. 7. If there are free electrons or the octet rule is still being violated, try forming molecules with resonance structures. C. Determining the Electronic & Molecular Geometry of a Compound or Ion Determine the electronic and molecular geometry of each of the molecules or polyatomic ions. 1. Determine the correct Lewis Dot Structure. 2. Determine the number of regions containing bonding electrons and non-bonding electrons (lone pair electrons). 3. Based on the number of total regions of electrons, determine the electronic geometry. 4. Based on the number of regions of bonding electrons and non-bonding electrons, determine the molecular geometry. 5. Using the molecular models build the molecules or polyatomic ions. 52 | P a g e 6. Table 1 contains the information for electronic and molecular geometry, and bond angles. Table 1. Table of Molecular Geometric Information. Regions of electronic density Electronic Geometry 2 Linear 3 4 5 6 Trigonal Planer Tetrahedral Trigonal Bipyramidal Octahedral Bond Angle Number of Bonding regions Number of Non-bonding (lone pair) Regions Molecular Geometry 180 2 0 Linear 1 1 Linear 3 0 Trigonal Planer 2 1 Bent 4 0 3 1 2 2 5 0 4 1 3 2 2 3 6 0 5 1 4 2 3 3 2 4 120 109.5 120, 90 90 Tetrahedral Trigonal Pyramidal Bent Trigonal Bipyramidal See-Saw “T”- Shaped Linear Octahedral Square Pydramidal Square Planer “T” – Shaped Linear 53 | P a g e D. Determining if a Molecule is Polar or Non-Polar. Once the geometry of the molecule has been determined, the polarity of the molecule can be determined using the following method. 1. Use the steps to determine the electronic and molecular geometries. 2. Determine if the molecule is symmetric or non-symmetric. a. Symmetric molecule: Non-polar molecule b. Non-Symmetric molecule: Polar molecule 54 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Data & Observations: A. Writing the Lewis Dot Structure for monoatomic ion or atoms. Table 2. Lewis Dot Structure for atom or monoatomic ion. Element Chemical Symbol Number of Valance Electrons Lewis Dot Structure Potassium Arsenic Magnesium Sulfur Bromine Boron Sodium ion 55 | P a g e Iodide ion Oxygen ion Rubidium ion Calcium ion B. Molecular Models. Group 1 Group 2 Group 3 Group 4 Group 5 Group 6 SiCl4 CHBr3 CH2Cl2 CF4 SiH4 NH4+ PCl3 ICl2+ ClO3- OF2 AsBr3 SF2 NO3- NO2- SO2 COH2 O3 COF2 SO42- PO43- ClO4- CH3- BrO4- NH3 O2 N2 CO2 CS2 CN- CSO HOOH BF3 BBr3 CH3OH CO32- BCl3 56 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ _______________ (Assigned molecule/ion #1) _______________ (Assigned molecule/ion #2) 1. Number of valance electrons __________ 1. Number of valance electrons __________ 2. Lewis Dot Structure: 2. Lewis Dot Structure: 3. Electronic Geometry: 3. Electronic Geometry: 4. Molecular Geometry: 4. Molecular Geometry: 5. Bond Angles: 5. Bond Angles: 6. Molecular Model ______________ (Instructors Initials) 6. Molecular Model ______________ (Instructors Initials) 7. Polar or Non-Polar molecules: (Explain) 7. Polar or Non-Polar molecules: (Explain) 57 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ _______________ (Assigned molecule/ion #3) _______________ (Assigned molecule/ion #4) 1. Number of valance electrons __________ 1. Number of valance electrons __________ 2. Lewis Dot Structure: 2. Lewis Dot Structure: 3. Electronic Geometry: 3. Electronic Geometry: 4. Molecular Geometry: 4. Molecular Geometry: 5. Bond Angles: 5. Bond Angles: 6. Molecular Model ______________ (Instructors Initials) 6. Molecular Model ______________ (Instructors Initials) 7. Polar or Non-Polar molecules: (Explain) 7. Polar or Non-Polar molecules: (Explain) 58 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ _______________ (Assigned molecule/ion #5) _______________ (Assigned molecule/ion #6) 1. Number of valance electrons __________ 1. Number of valance electrons __________ 2. Lewis Dot Structure: 2. Lewis Dot Structure: 3. Electronic Geometry: 3. Electronic Geometry: 4. Molecular Geometry: 4. Molecular Geometry: 5. Bond Angles: 5. Bond Angles: 6. Molecular Model ______________ (Instructors Initials) 6. Molecular Model ______________ (Instructors Initials) 7. Polar or Non-Polar molecules: (Explain) 7. Polar or Non-Polar molecules: (Explain) 59 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Answer the following questions using the molecules/ions assigned to your group. 1. Which molecules are polar? 2. Which molecules are non-polar? 3. Which molecules/ions violate the “octet rule”? Explain. 4. Is it necessary to determine the polarity of an ion? Explain 60 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Post-Lab Questions: 1. Explain the factors to determine the molecular geometry of a molecule or polyatomic ion? 2. Explain how the molecular geometry is different from the electronic geometry. 3. What is the key difference between a polar molecular and a non-polar molecule? 61 | P a g e Experiment 7: Stoichiometry Purpose: In this activity you and your lab partners will analyze several sets of data and solve the problems based upon the information provided. The reactants will be graham crackers (Gc), chocolate (Ch), and marshmallows (M) and the final product will be a S’more (Gc2Ch2M) which can be abbreviated as Sm. The formula for the reaction is as follows: 2Gc + 2Ch + 1M Gc2Ch2M (can be abbreviated as Sm) 1 mole Gc = 1 square of a graham cracker = 1.94 g 1 mole Ch = 1 rectangle of chocolate = 2.51 g 1 mole M = 1 marshmallow = 0.53 g Procedure: 1. Use the balanced equation provided above to answer the data analysis questions. 2. Determine if your answers are correct by comparing your responses with your classmates. Redo any problems that are not. 3. After washing your hands, obtain the necessary supplies to make your S’mores. 4. Use safety precautions when lighting the Bunsen Burners. 5. Always wear goggles when using Bunsen Burners. 6. ENJOY! 7. Make additional S’mores depending upon supplies. 8. Pay close attention to clean up so that you can partake in Cooking Chemistry Labs in the future :) 62 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Data Analysis: 1. How many grams are found in 2 moles of a graham cracker? 2. How many grams are found in 6 moles of chocolate? 3. How many grams are found in 10 moles of marshmallows? 4. How many moles are in 5.82 g of Gc? 5. How many moles are in 5.02 g of Ch? 6. How many moles are in 5.30 g of M? 7. What is the molar mass of a S'more? 8. How many atoms will be found in two moles of chocolate? 9. How many molecules will be found in one mole of a S'more? 63 | P a g e 10. What is the mole ratio between chocolate and graham crackers? Marshmallows and graham crackers? Marshmallows and S'mores? 11. If you have three marshmallows, how many S'mores can you make? 12. If you have 5 squares of graham crackers, how many complete S'mores can you make? 13. Given 5 moles of Gc and excess of the other reactants, how many S'mores can you make? How does this question relate to question #12? 14. If 10 moles Ch reacts with M and Gc, how many moles of S'mores will be produced? 15. If 12 moles of Sm were produced, how many moles of Ch were required to make this product? 16. If you wanted to make enough S'mores for a class of 37 students, how many moles of Ch, Gc, and M would you need? 17. Given 21.6 g of M, how many moles of S’mores can you make? 18. Given 36.7 g of Gc, how many grams of M will you need? 64 | P a g e 19. Given 8 moles of Gc, how many grams of M will be needed? 20. To produce 97.8 g of Sm, how many moles of Gc will you need? 21. Based on the amount of Gc, Ch, and M your group received, which reactant would be the limiting reagent (reactant)? 22. How many moles of S’mores did you produce? 23. How many grams of S’mores did you produce? 24. What were the reactants in excess? 25. How many moles of each excess reactant were left? 26. How many grams of each excess reactant were left? 65 | P a g e Experiment 8: Alum Synthesis Purpose: The term alum is a general family name for a crystalline substance composed of cations with 1+ and 3+ charges. In this experiment, you will synthesize a type of alum called potassium aluminum sulfate dodecahydrate, KAl(SO4)2•12H2O. You will synthesize this compound by placing the appropriate ions in one container in aqueous solution and then evaporate the water to form the alum crystals. 2 Al(s) + 2 KOH(aq) + 6 H2O(l) 2 KAl(OH)4(aq) + 3 H2(g) This particular compound has been chosen because it is relatively simple to prepare a pure sample. The process of synthesizing this compound is interesting in that it involves both chemical and physical reactions. Chemically, aluminum is oxidized from aluminum foil to prepare the Al3+ ions. Physically, as the solution that contains the mixture of ions evaporates, crystals will form which contain six waters of hydration bonded to the aluminum ion and six waters bonded to the potassium ion. Aluminum is considered a reactive metal, but because its surface is usually protected by a thin film of aluminum oxide, it reacts slowly with acids. It does, however, dissolve quickly in basic solutions. Excess hydroxide ion converts the aluminum to the tetrahydroxoaluminate (Al(OH)3) precipitates. Continued addition of acid causes the hydroxide ions to be completely neutralized, and the aluminum exists in solution as the hydrated ion [Al(H2O)6]3+. Aluminum hydroxide is considered to be an amphoteric hydroxide because it dissolves in both acids and bases. In this experiment, you will Synthesize a sample of potassium aluminum sulfate dodecahydrate (alum). Observe and record the process of synthesizing a compound. Calculate the percent yield of your synthesis. Procedure: 1. Obtain and wear goggles. 2. Obtain a piece of aluminum foil and measure its mass. For best results, you should have about 1.00 g of aluminum. Tear the foil into small pieces and place the pieces in a 250 mL beaker. 3. Set up a Büchner funnel and filter flask so that you are ready to filter the reaction mixture that will be produced in Step 4. 4. Conduct the first part of the synthesis. CAUTION: Potassium hydroxide solution is caustic. Avoid spilling it on your skin or clothing. a. Use a graduated cylinder to measure out 25 mL of 3 M KOH solution. 66 | P a g e b. Slowly add the KOH solution to the beaker of aluminum pieces. Notice that the reaction is exothermic. Allow the reaction to proceed until all of the foil is dissolved. c. Carefully pour the reaction mixture through your Büchner funnel and filter flask setup, and rinse the filter paper with a small amount of distilled water. Note: The reaction mixture contains three ions: K+, [Al(OH)4–], and excess OH–. d. Rinse the beaker with distilled water, and pour the filtered liquid back into the beaker. 5. Allow the solution to cool to near room temperature. If you are pressed for time, you may cover the beaker with plastic wrap or Parafilm, and store the liquid overnight. 6. Clean the Büchner funnel and filter flask, and prepare it for more filtering that you may need to do in Step 7 or Step 10. 7. Complete the synthesis. a. Use a graduated cylinder to measure out 35 mL of 3 M H2SO4 solution. CAUTION: The reaction mixture must be cooled to room temperature before proceeding. Handle the H2SO4 solution with care. It can cause painful burns if it comes in contact with the skin. b. After the reaction mixture has cooled, slowly add the sulfuric acid solution to the beaker of liquid. Stir the mixture constantly. The reaction is strongly exothermic, so be careful as you stir the mixture. Note that aluminum hydroxide will precipitate initially. It will dissolve as more sulfuric acid is added. c. If there is some solid remaining in the beaker after the 35 mL of sulfuric acid has been added, pour the mixture through the Büchner funnel and filter flask to separate the undissolved solid from the mixture. 8. Gently boil your mixture until you have about 50 mL of liquid in the beaker. 9. Cool the beaker of solution. Choose one of the two methods listed below. a. Allow the solution to cool overnight. In most cases, this gradual cooling forms a good crop of alum crystals. b. Prepare an ice bath for the 250 mL beaker. Place your beaker of solution, uncovered, in the ice bath. Do not move the ice bath or the beaker. After about fifteen minutes, crystals of alum will appear in the beaker. If there are no crystals after fifteen minutes, scrape the bottom of the beaker with a glass stirring rod to create a rough spot for crystal growth. You may also heat the solution to evaporate more water and cool the solution again. 10. Collect your alum crystals by pouring them onto the Büchner funnel and filter-flask setup. Use vacuum filtration to wash the crystals on the filter paper with 50 mL of an aqueous ethanol solution (50%). The crystals will not dissolve in this solution. 67 | P a g e 11. Remove the filter and crystals from the Büchner funnel and allow the crystals to dry at room temperature. Measure and record the mass of your sample of alum. Store the crystals for further analysis. 68 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ PRE-LAB QUESTIONS 1. What are the two cautions you need to take when conducting the synthesis of alum experiment? 2. Fill in the blanks for the physical properties of alum: a. Molecular Weight: _____________g/mol b. Melting Point: _________________ oC 3. A student conducting the synthesis of alum experiment started with 0.9156 g of aluminum foil. a. What is the theoretical yield of alum? (Show your work) b. The student isolated 8.3181 g of alum. What was the percent yield the student obtained? (Show your work) 69 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ DATA TABLE Mass of Aluminum Foil used: ___________________g Moles of Aluminum Foil used: __________________mol Mole Ratio of Aluminum Foil to Alum: ___________ratio Moles of Alum theoretically obtained: ____________mol Theoretical Mass of Alum: _____________________g Actual Mass of Alum obtained: _________________g DATA ANALYSIS 1. Determine the theoretical yield of the alum. Use the aluminum foil as the limiting reagent and presume that the foil was pure aluminum. 2. Calculate the percent yield of your alum crystals. 3. Discuss the factors that affected the percent yield. 70 | P a g e 4. Write the balanced equations for the following: (a) aluminum and potassium hydroxide, yielding [Al(OH)4] – and hydrogen gas; (b) hydrogen ions and [Al(OH)4] –, yielding aluminum hydroxide; (c) aluminum hydroxide and hydrogen ions, yielding [Al(H2O)6]3+; and (d) the formation of alum from potassium ions, sulfate ions, [Al(H2O)6]3+, and water. 71 | P a g e Experiment 9: Alum Analysis Purpose: After a compound has been synthesized, tests should be carried out to verify that the compound formed is indeed the compound desired. There are a number of tests that can be performed to verify that the compound is the one desired. In the previous experiment, you prepared alum crystals, KAl(SO4)2•12H2O. In this experiment, you will conduct a series of tests to determine if your crystals are really alum. The first test is to find the melting temperature of the compound and compare this value with the accepted (published) value for alum (92.5°C). The second test determines the water of hydration present in the alum crystals. The third test is a chemical test to determine the percent sulfate in your sample of alum. In this experiment, you will Determine the melting temperature of a sample of alum. Determine the water of hydration of a sample of alum. Procedure: Part I: Determine the Melting Temperature 1. Obtain and wear goggles. 2. Connect the Temperature Probe to LabQuest. 3. Take a piece of weighing paper folded in half and using a micro-spatula remove a small amount of alum from the sample you prepared in the past experiment. Place the alum into the fold of the weighing paper, fold the weighing paper over, and using the edge of the micro-spatula pulverize the alum sample. Use the micro-spatula to pile the alum in the weighing paper. Push the open end of a capillary tube into the pile of the alum powder. Pack alum into the capillary tube to a depth of about 0.5 cm by tapping the tube lightly on the table top. 4. Use a rubber band to fasten the capillary tube to the Temperature Probe. The tip of the tube should be even with the tip of the probe. Use a utility clamp to connect the Temperature Probe to a ring stand (see Figure 1). 72 | P a g e Figure 1 5. Prepare a water bath to be heated by a hot plate. 6. Monitor the temperature readings on the Main screen. Immerse the capillary tube and Temperature Probe in the water bath. Warm the alum sample at a gradual rate so that you can precisely determine the melting temperature. The white powder will become clear when it is melting. Observe the temperature readings and record the precise melting temperature when the substance is completely clear. 7. Conduct a second test with a new sample of alum in a new capillary tube. Part II Determine the Water of Hydration 8. Heat a crucible with cover over a burner flame until it is red hot. Allow the crucible to cool, and then measure the total mass of crucible and cover. Handle the crucible with tongs or forceps to avoid getting fingerprints on it. 9. Place about 2 g of your alum crystals in the crucible, and then measure the mass of the crucible, cover, and alum. Record this measurement in the data table. 10. Set up a ring, ring stand and triangle over a lab burner. Use tongs or forceps to set the crucible at an angle on the triangle and place the cover loosely on the crucible. Use a lab burner to very gently heat the crucible of alum until you can see no vapor escaping from the crucible. It is important that the vapor does not carry any alum with it. After the vapor is gone, heat the crucible more strongly for five minutes, and then cool the crucible. 11. Measure and record the mass of crucible, cover, and alum after drying. This would be the mass after the first heating. 12. Reheat the crucible and alum sample for five additional minutes. Cool and measure the mass of the crucible again. This would be the mass after the second heating. If the two masses are 73 | P a g e the same (or very nearly so), the test is done. If not, repeat the heating and weighing until a constant mass is obtained. 74 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ PRE-LAB QUESTIONS 1. Write the definition/explanation of each of the following terms: a. Hydrated Salt b. Anhydrous Salt c. Hygroscopic substance d. Desiccant substance 75 | P a g e 2. A student was asked to identify an unknown hydrate by following the following the procedure in this experiment to determine the percent water of alum. A 2.752 g sample of the unknown sample is heated and weighted after cooling to constant weight of 1.941 g. The unknown is believed to be one of the following compounds: LiNO3.3H2O, Ca(NO3)2.4H2O, or Sr(NO3)2.4H2O. a. Calculate the percent water in the unknown sample. b. Calculate the percent water in the three known samples. c. What is the identity of the unknown sample? Explain your reasoning. 76 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ DATA TABLE Part I Melting Temperature Test Results Trial 1 Trial 2 Trial 1 Trial 2 Melting Temperature (°C) Part II Water of Hydration Test Results Mass of crucible and cover (g) Mass of crucible, cover, and alum before heating (g) Mass of Alum before heating (g) Mass of crucible, cover, and alum after 1st heating (g) Mass of anhydrous Alum after 1st heating (g) Mass of crucible, cover, and alum after 2nd heating (g) Mass of anhydrous Alum after 2nd heating (g) Mass of crucible, cover, and alum after final heating (g) Mass of anhydrous Alum after final heating (g) Average mass of anhydrous Alum (g) Moles of anhydrous Alum (mol) Average mass of water lost (g) Moles of water (mol) Ratio of moles Alum : moles water Chemical Formula of Alum 77 | P a g e DATA ANALYSIS 1. Is your sample alum? Use the results of the three tests to support your answer. Discuss the accuracy of your tests and possible sources of experimental error. 2. Suggest other tests that could be conducted to verify the composition of your alum. 3. If the melting temperature test was the only test that you conducted, how confident would you be in the identification of your sample? Explain. 78 | P a g e Experiment 10: Identification Unknown Solution (Solubility) Purpose: Students investigate the procedure to identify unknown solutions bases on the observation of chemical changes. What are the identities of the each of the unknowns? Introduction: Chemical change occurs when molecules of two substances are allowed to come in contact with each other with enough energy for change to occur with the molecules. One can observe if a chemical reaction occurred if there is an observable change. Observable changes are: 1. 2. 3. 4. Precipitate forms (solid formed when two solutions are mixed) Color change Bubbles form (gases are released) Heat generated Observing one of these indicates a chemical reaction has occurred between two molecules. These signs of a chemical reaction are especially evident when dealing with ionic compounds. Ionic compounds are made up of ions, positively or negatively charged atoms or molecules. Positively charged atoms or molecules are called cations, and negatively charged atoms or molecules are called anions. The charged molecules are termed polyatomic ions. The combination of cations and anions form the neutral ionic compound. Many ionic compounds dissolve in water and when the ionic compound dissolves in the water, the ions separate into the cations and anions. These are free to interact with other ions; cations react with other anions and anions reacting with other cations. Many of the interactions of the cation and anion do not generate one of the signs of a chemical reaction. But if there is a sign of a chemical reaction, knowing the initial ionic compounds can be used to determine unknown solutions. Even the observable signs of a chemical reaction can help determine unknown solutions: 1. Precipitation reactions produce a solid when two solutions are mixed together. The solutions are ionic solutions and the combination of the specific cation and anion result in the precipitate formation. The rules for precipitation can be used to determine whether two solutions will produce a precipitate. 2. Color changes tend to involve Transition metal cations. 3. Bubbles forming tend to be single replacement/displacement reactions (redox). Or CO2 generated reaction involving carbonate or bicarbonate. 4. Heat generating reactions tend to be acids reacting bases, either the acid or base needs to be strong. 79 | P a g e The rules for solubility (soluble means dissolved, insoluble means does not dissolve) outline when two solutions are mixed together, whether the resulting mixture will have a precipitate formed. These are the rules for solubility: 1. All soluble when cation is alkali metal or ammonium ion. 2. All soluble when anion is nitrate, acetate, perchlorate. 3. All soluble when anion is halide. a. Except when cation is silver ion, lead(II) ion, mercury(II) ion. 4. All soluble when anion is sulfate. a. Except when the cation is calcium, barium, 5. All insoluble when anion is hydroxide. a. Except when the cation is alkali metal or alkaline earth metal. 6. All insoluble when anion is sulfide, carbonate, phosphate. a. Except when cation is alkali metal. 80 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Pre-Lab Questions: 1. When a gaseous product is formed, will one see 1-large bubble or many small bubbles form? (Hint: Think of a carbonated beverage.) Explain. 2. Define a precipitate. 3. Match the terms with the abbreviations, place the number of the abbreviation in the space provided of the matching term: _____ a. solution 1. unk _____ b. precipitate 2. aq _____ c. reaction 3. rxn _____ d. aqueous 4. ppt _____ e. unknown 5. soln 81 | P a g e 4. Complete the chemical equation for each of the following. Be sure to include the state of each compound or if it is in solution. If a chemical reaction is observed, write the net ionic equation. BaCl2(aq) + KIO4(aq) Sr(NO3)2(aq) + Na2SO4(aq) Pb(NO3)2(aq) + KOH(aq) CuSO4(aq) + Na2CO3(aq) HCl(aq) + NaOH(aq) 82 | P a g e Procedure: Students will conduct the procedure for this experiment twice, once with a known set of solutions and once with the unknown set of solutions. The object of this experiment is to determine the unknown solutions. A. Studying the reactions of the known set of solutions. Students will obtain the six known solutions and a 96-well plate. Using Table 1 in the data & observation section, construct a similar matrix using the 96-well plate. Choose one of the solutions to be added to row 1, this will be solution 1. Write the name of the compound for solution 1 in the space provided on Table 1. Add one drop of the solution to each of the first five wells, in the first row, of the 96-well plate. Choose one of the solutions to be added to row 2, this will be solution 2. Write the name of the compound for solution 2 in the space provided on Table 1. Add one drop of the solution to each of the first four wells, in the second row, of the 96-well plate. Choose another solution for row 3; write the name of that solution in the space provided on Table 1. Add one drop of the solution to each of the first three wells of the 96-well plate. Choose a fourth solution to be added to row 4. Write the name of the solution on row 4 of Table 1. Add a drop of solution 4 to the first 2 wells of the 96-well plate in row 4. Choose a fifth solution to be added to row 5. Write the name of the solution on row 5 of Table 1. Add a drop if solution 5 to the first well of the 96-well plate in row 5. The name of the sixth solution will be written into the first column of Table 1. Add one drop of solution 6 to the first five wells of column 1. Take solution 5 and write the name of solution 5 in the space provided in column 2. Add one drop of solution 5 to the first four wells of column 2. Take solution 4 and write the name of solution 4 in the space provided in column 3. Add one drop of solution 4 to the first three wells of column 3. Take solution 3 and write the name of solution 3 in the space provided in column 4. Add one drop of solution 3 to the first two wells of column 4. Take solution 2 and write the name of solution 2 in the space provided in column 5. Add one drop of solution 2 to the first well of column 5. Observe each well for a sign of a reaction. If no observable sign of a reaction, write N.R. on Table 1 where the two solutions intersect on the table. If there is a color change, write the color of the solution. If a precipitate forms, write “ppt” and the color of the precipitate. If bubbles form, write “bubbles”, and if heat was generated, write “heat”. B. Studying the reactions of the unknown set of solutions. Students will obtain a set of six unknown solutions and another 96-well plate. Using Table 2 in the data & observation section, repeat the above procedure. In the space provided, write the unknown number. Based on the results of the reactions of the unknowns and comparing those data with the data in Table 1, the student will determine each of the unknowns and report those results. 83 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Data & Observations: A. Studying the reactions of the known set of solutions. Table 1. Matrix of reactions of known solutions. Solution 6 Solution 5 Solution 4 ___________ ___________ ___________ Solution 3 Solution 2 ___________ ___________ Solution 1 __________ Solution 2 __________ Solution 3 __________ Solution 4 ___________ Solution 5 ___________ 84 | P a g e B. Studying the reactions of the unknown set of solutions. Table 2. Matrix of reactions of unknown solutions. Unknown 6 Unknown 5 Unknown 4 ___________ ___________ ___________ Unknown 3 Unknown 2 ___________ ___________ Unknown 1 __________ Unknown 2 __________ Unknown 3 __________ Unknown 4 ___________ Unknown 5 ___________ 85 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ C. Identify the Unknowns. Unknown 1: ____________________ Experimental reasoning for your identification: Unknown 2: ____________________ Experimental reasoning for your identification: Unknown 3: ____________________ Experimental reasoning for your identification: Unknown 4: ____________________ Experimental reasoning for your identification: Unknown 5: ____________________ Experimental reasoning for your identification: Unknown 6: ____________________ Experimental reasoning for your identification: 86 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ 1. Write the molecular equation for each of the reactions that produced a precipitate. 2. Write the net ionic equation for each of the reactions that produced a precipitate. 87 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ Post-Lab Questions: 1. A student was given the following five known solutions; NiCl2, Pb(NO3)2, NaOH, Na2C2O4, and CuSO4. After conducting an experiment mixing the solutions in pairs and recording their observations in Table 3. Table 3. Precipitate formation by mixing the pairs of known solutions. CuSO4 Na2C2O4 NaOH Pb(NO3)2 NiCl2 N.R. N.R. grenn ppt N.R. Pb(NO3)2 white ppt white ppt white ppt NaOH blue ppt N.R. Na2C2O4 white ppt The student was then given the same five solutions as unknowns labeled Solutions A-E. After conducted the same mixing experiment as with the known solutions, the student recorded their observations in Table 4. 88 | P a g e Table 4. Precipitate formation by mixing the pairs of unknown solutions. Solution E Solution D Solution C Solution B Solution A blue ppt white ppt N.R. green ppt Solution B N.R. N.R. N.R. Solution C white ppt white ppt Solution D white ppt Identify the unknowns: Solution A: ____________ Solution B: ____________ Solution C: ____________ Solution D: ____________ Solution E: ____________ 89 | P a g e Experiment 11: Acid-Base Titration 90 | P a g e Experiment 12: Household Acids & Bases Purpose: Many common household solutions contain acids and bases. Acid-base indicators, such as litmus and red cabbage juice, turn different colors in acidic and basic solutions. They can, therefore, be used to show if a solution is acidic or basic. An acid turns blue litmus paper red, and a base turns red litmus paper blue. The acidity of a solution can be expressed using the pH scale. Acidic solutions have pH values less than 7, basic solutions have pH values greater than 7, and neutral solutions have a pH value equal to 7. In this experiment, you will use litmus and a pH Sensor to determine the pH values of household substances. After adding red cabbage juice to the same substances, you will determine the different red cabbage juice indicator colors over the entire pH range. In this experiment, you will Use litmus paper and a pH Sensor to determine the pH values of household substances. Add cabbage juice to the same substances and determine different red cabbage juice indicator colors over the entire pH range. Figure 1 91 | P a g e Procedure: 1. Obtain and wear goggles. CAUTION: Do not eat or drink in the laboratory. Part I Litmus Tests 2. Label 7 test tubes with the numbers 1–7 and place them in a test tube rack. 3. Measure 3 mL of vinegar into test tube #1. Refer to the data table and fill each of the test tubes 2–7 to about the same level with its respective solution. CAUTION: Ammonia solution is toxic. Its liquid and vapor are extremely irritating, especially to eyes. Drain cleaner solution is corrosive. Handle these solutions with care. Do not allow the solutions to contact your skin or clothing. Wear goggles at all times. Notify your teacher immediately in the event of an accident. 4. Use a stirring rod to transfer one drop of vinegar to a small piece of blue litmus paper on a paper towel. Transfer one drop to a piece of red litmus paper on a paper towel. Record the results. Clean and dry the stirring rod each time. 5. Test solutions 2–7 using the same procedure. Be sure to clean and dry the stirring rod each time. Part II Red Cabbage Juice Indicator 6. After you have finished the Part I litmus tests, add 3 mL of red cabbage juice indicator to each of the 7 test tubes. Record your observations. Dispose of the test-tube contents as directed by your teacher. Part III pH Tests 7. Prepare the pH Sensor for data collection. a. Connect the pH Sensor to LabQuest and choose New from the File menu. If you have an older sensor that does not auto-ID, manually set up the sensor. b. Remove the pH Sensor from the sensor storage solution bottle by unscrewing the lid. Carefully remove the bottle, leaving the cap on the sensor body. c. Rinse the tip of the sensor with distilled water and place the sensor tip into a beaker containing sensor soaking solution. Use a utility clamp to fasten the pH Sensor to a ring stand, as shown in Figure 1. 8. Raise the pH Sensor from the sensor soaking solution and set the solution aside. Use a wash bottle filled with distilled water to thoroughly rinse the pH Sensor. Catch the rinse water in a 250 mL beaker. 9. Obtain one of the 7 solutions in the small container supplied by your teacher. Raise the solution to the pH Sensor and swirl the solution about the sensor. When the pH reading stabilizes, record the pH value. 92 | P a g e 10. Prepare the pH Sensor for reuse. a. Rinse it with distilled water from a wash bottle. b. Place the sensor into the sensor soaking solution and swirl the solution about the sensor briefly. c. Rinse with distilled water again. 11. Determine the pH of the other solutions using the Step 9 procedure. You must clean the pH Sensor between tests using the Step 10 procedure. 12. When you are finished, rinse the sensor with distilled water and return it to the sensor soaking solution. 93 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ PROCESSING THE DATA 1. Which of the household solutions tested are acids? How can you tell? 2. Which of the solutions are bases? How can you tell? 3. What color(s) is red cabbage juice indicator in acids? In bases? 4. Can red cabbage juice indicator be used to determine the strength of acids and bases? Explain. 5. List advantages and disadvantages of litmus and red cabbage juice indicators. 94 | P a g e DATA TABLE Test Tube Solution 1 vinegar 2 ammonia 3 lemon juice 4 soft drink 5 drain cleaner 6 detergent 7 baking soda 8 antacid Blue Litmus Red Litmus Red Cabbage Juice pH 95 | P a g e Experiment 13: Boyles Law & Guy-Lussac Boyles Law Purpose: The primary objective of this experiment is to determine the relationship between the pressure and volume of a confined gas. The gas we use will be air, and it will be confined in a syringe connected to a Gas Pressure Sensor (see Figure 1). When the volume of the syringe is changed by moving the piston, a change occurs in the pressure exerted by the confined gas. This pressure change will be monitored using a Gas Pressure Sensor. It is assumed that temperature will be constant throughout the experiment. Pressure and volume data pairs will be collected during this experiment and then analyzed. From the data and graph, you should be able to determine what kind of mathematical relationship exists between the pressure and volume of the confined gas. Historically, this relationship was first established by Robert Boyle in 1662 and has since been known as Boyle’s law. Figure 1 Procedure: 1. Prepare the Gas Pressure Sensor and an air sample for data collection. a. Connect the Gas Pressure Sensor to LabQuest. If you have an older sensor that does not auto-ID, manually set up the sensor. b. With the 20 mL syringe disconnected from the Gas Pressure Sensor, move the piston of the syringe until the front edge of the inside black ring (indicated by the arrow in Figure 1) is positioned at the 10.0 mL mark. c. Attach the 20 mL syringe to the valve of the Gas Pressure Sensor. 2. You are now ready to collect pressure and volume data. It is easiest if one person takes care of the gas syringe and another enters volumes. a. Move the piston so the front edge of the inside black ring (see Figure 2) is positioned at the 5.0 mL line on the syringe. Hold the piston firmly in this position until the pressure value displayed on the screen stabilizes. 96 | P a g e Figure 2 b. Continue this procedure using syringe volumes of 10.0, 12.5, 15.0, 17.5, and 20.0 mL. 97 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ DATA AND CALCULATIONS Volume (mL) Pressure (kPa) Constant, k (P / V or P • V) 5.00 10.0 12.5 15.0 17.5 20.0 PROCESSING THE DATA 1. Based on your data, what would you expect the pressure to be if the volume of the syringe was increased to 40.0 mL. Explain or show work to support your answer. 2. Based on your data, what would you expect the pressure to be if the volume of the syringe was decreased to 2.5 mL. 3. What experimental factors are assumed to be constant in this experiment? 98 | P a g e 4. One way to determine if a relationship is inverse or direct is to find a proportionality constant, k, from the data. If this relationship is directly proportional, k = P/V. If it is inverse proportionsal, k = P•V. Choose one of these formulas based on your knowledge of Boyle’s Law, and calculate k for the seven Pressure-Volume data pairs in your data table. Based on your answers to in the Data and Calculations Table, does the relationship hold true? Explain your answer. 5. Graph the pressure vs volume relationship or the pressure vs inverse volume (1/volume) relationship using Excel. Using the graph, determine if the relationship between pressure and volume is a direct or inverse proportionality. Since the graph showing proportionality should be a linear graph, you must have the trendline, r2 value, and the equation for the line on the graph. Your graph should also have a title, and axis’s labeled including units. 99 | P a g e Experiment 13: Boyles Law & Guy-Lussac Guy-Lussac Law Purpose: Gases are made up of molecules that are in constant motion and exert pressure when they collide with the walls of their container. The velocity and the number of collisions of these molecules are affected when the temperature of the gas increases or decreases. In this experiment, you will study the relationship between the temperature of a gas sample and the pressure it exerts. Using the apparatus shown in Figure 1, you will place an Erlenmeyer flask containing an air sample in water baths of varying temperature. Pressure will be monitored with a Pressure Sensor and temperature will be monitored using a Temperature Probe. The volume of the gas sample and the number of molecules it contains will be kept constant. Pressure and temperature data pairs will be collected during the experiment and then analyzed. From the data and graph, you will determine what kind of mathematical relationship exists between the pressure and absolute temperature of a confined gas. You may also do the extension exercise and use your data to find a value for absolute zero on the Celsius temperature scale. OBJECTIVES In this experiment, you will Study the relationship between the temperature of a gas sample and the pressure it exerts. Determine from the data and graph, the mathematical relationship between pressure and absolute temperature of a confined gas. Find a value for absolute zero on the Celsius temperature scale. Figure 1 100 | P a g e PROCEDURE 1. Obtain and wear goggles. 2. Prepare a boiling-water bath. Put about 800 mL of hot tap water into a l L beaker and place it on a hot plate. Turn the hot plate to a high setting. 3. Prepare an ice-water bath. Put about 700 mL of cold tap water into a second 1 L beaker and add ice. 4. Put about 800 mL of room-temperature water into a third 1 L beaker. 5. Put about 800 mL of hot tap water into a fourth 1 L beaker. 6. Prepare the Temperature Probe and Gas Pressure Sensor for data collection. a. Connect the Gas Pressure Sensor to Channel 1 of LabQuest and the Temperature Probe to Channel 2. Choose New from the File menu. If you have older sensors that do not auto-ID, manually set up the sensors. Figure 2 b. Obtain a rubber-stopper assembly with a piece of heavy-wall plastic tubing connected to one of its two valves. Attach the connector at the free end of the plastic tubing to the open stem of the Gas Pressure Sensor with a clockwise turn. Leave its two-way valve on the rubber stopper open (lined up with the valve stem as shown in Figure 2) until Step 6d. c. Insert the rubber-stopper assembly into a 125 mL Erlenmeyer flask. Important: Twist the stopper into the neck of the flask to ensure a tight fit. Figure 3 d. Close the 2-way valve above the rubber stopper—do this by turning the valve handle so it is perpendicular with the valve stem itself (as shown in Figure 3). The air sample to be studied is now confined in the flask. 9. Start recording data. Pressure readings (in kPa) and temperature readings (in °C) and are displayed on the screen. 101 | P a g e 10. Collect pressure vs. temperature data for your gas sample. a. Place the flask into the ice-water bath. Make sure the entire flask is covered (see Figure 3). b. Place the Temperature Probe into the ice-water bath. c. When the temperature and pressure readings have both stabilized, record the temperature and pressure readings. 11. Repeat Step 10 using the room-temperature bath. 12. Repeat Step 10 using the hot-water bath. 13. Use a ring stand and utility clamp to suspend the Temperature Probe in the boiling-water bath. CAUTION: Do not burn yourself or the probe wires with the hot plate. To keep from burning your hand, hold the tubing of the flask using a glove or a cloth. After the Temperature Probe has been in the boiling water for a few seconds, place the flask into the boiling-water bath and repeat Step 10. 102 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ PROCESSING THE DATA 1. In order to perform this experiment, what two experimental factors were kept constant? 2. Based on the data and graph that you obtained for this experiment, express in words the relationship between gas pressure and temperature. 3. Write an equation to express the relationship between pressure and temperature (K). Use the symbols P, T, and k. 4. One way to determine if a relationship is inverse or direct is to find a proportionality constant, k, from the data. If this relationship is direct, k = P/T. If it is inverse, k = P•T. Based on your answer to Question 3, choose one of these formulas and calculate k for the four ordered pairs in your data table (divide or multiply the P and T values). Show the answer in the fourth column of the Data and Calculations table. How “constant” were your values? 5. Graph the pressure vs volume relationship or the pressure vs inverse volume (1/volume) relationship using Excel. Using the graph, determine if the relationship between pressure and volume is a direct or inverse proportionality. Since the graph showing proportionality should be a linear graph, you must have the trendline, r2 value, and the 103 | P a g e equation for the line on the graph. Your graph should also have a title, and axis’s labeled including units. 6. According to this experiment, what should happen to the pressure of a gas if the Kelvin temperature is doubled? Check this assumption by finding the pressure at –73°C (200 K) and at 127°C (400 K) on your graph of pressure versus temperature. How do these two pressure values compare? DATA AND CALCULATIONS Pressure (kPa) Temperature (°C) Temperature (K) Constant, k (P / T or P•T) 104 | P a g e Experiment 14: Molar Volume 105 | P a g e Experiment 15: Heat of Fusion Purpose: Melting and freezing behavior are among the characteristic properties that give a pure substance its unique identity. As energy is added, pure solid water (ice) at 0°C changes to liquid water at 0°C. In this experiment, you will determine the energy (in joules) required to melt one gram of ice. You will then determine the molar heat of fusion for ice (in kJ/mol). Excess ice will be added to warm water, at a known temperature, in a Styrofoam cup. The warm water will be cooled down to a temperature near 0°C by the ice. The energy required to melt the ice is removed from the warm water as it cools. To calculate the heat that flows from the water, you can use the relationship q = Cp•m•t where q stands for heat flow, Cp is specific heat capacity, m is mass in grams, and t is the change in temperature. For water, Cp is 4.18 J/g°C. In this experiment, you will Determine the energy (in Joules) required to melt one gram of ice. Determine the molar heat of fusion for ice (in kJ/mol). Procedure: Part I Freezing 1. Connect the Temperature Probe to LabQuest and choose New from the File menu. If you have an older sensor that does not auto-ID, manually set up the sensor. 2. Use a utility clamp to clamp the Temperature Probe on a ring stand as shown in Figure 1. 3. Place a Styrofoam cup into a 400 mL beaker as shown in Figure 1. 4. Use a 100 mL graduated cylinder to obtain 100.0 mL of water at about 60°C from your teacher. 5. Obtain 7 or 8 large ice cubes. 6. Lower the Temperature Probe into the warm water (to about 1 cm from the bottom). 7. Record the temperature reading, in °C, is displayed to the right of the graph. You will need to wait until the temperature reaches a maximum (it will take a few seconds for the cold probe to reach the temperature of the warm water). This maximum will determine the initial temperature, t1, of the water. As soon as this maximum temperature is reached, fill the 106 | P a g e Styrofoam cup with ice cubes. Shake excess water from the ice cubes before adding them (or dry with a paper towel). Record the maximum temperature, t1, in your data table. Figure 1 8. Use a stirring rod to stir the mixture as the temperature approaches 0°C. Important: As the ice melts, add more large ice cubes to keep the mixture full of ice! 9. When the temperature reaches about 4°C, quickly remove the unmelted ice (using tongs). Continue stirring until the temperature reaches a minimum (and begins to rise). This minimum temperature is the final temperature, t2, of the water. Record t2 in your data table. 10. Use the 100 mL graduated cylinder to measure the volume of water remaining in the Styrofoam cup to the nearest 0.1 mL. Record this as V2. 107 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ PROCESSING THE DATA 1. Use the equation t = t1 – t2 to determine t, the change in water temperature. 2. Subtract to determine the volume of ice that was melted (V2 –V1). 3. Find the mass of ice melted using the volume of melt (use 1.00 g/mL as the density of water). 4. Use the equation given in the introduction of this experiment to calculate the energy (in joules) released by the 100 g of liquid water as it cooled through t. 5. Now use the results obtained above to determine the heat of fusion, the energy required to melt one gram of ice (in J/g H2O). 6. Use your answer to Step 5 and the molar mass of water to calculate the molar heat of fusion for ice (in kJ/mol H2O). 7. Find the percent error for the molar heat of fusion value in Step 6. The accepted value for molar heat of fusion is 6.01 kJ/mol. 108 | P a g e DATA AND CALCULATIONS Initial water temperature, t1 °C Final water temperature, t2 °C Change in water temperature, t °C Final water volume, V2 mL Initial water volume, V1 mL Volume of melt mL 109 | P a g e Experiment 16: Evaporation & Intermolecular Interactions Purpose: In this experiment, Temperature Probes are placed in various liquids. Evaporation occurs when the probe is removed from the liquid’s container. This evaporation is an endothermic process that results in a temperature decrease. The magnitude of a temperature decrease is, like viscosity and boiling temperature, related to the strength of intermolecular forces of attraction. In this experiment, you will study temperature changes caused by the evaporation of several liquids and relate the temperature changes to the strength of intermolecular forces of attraction. You will use the results to predict, and then measure, the temperature change for several other liquids. You will encounter two types of organic compounds in this experiment—alkanes and alcohols. The two alkanes are pentane, C5H12, and hexane, C6H14. In addition to carbon and hydrogen atoms, alcohols also contain the -OH functional group. Methanol, CH3OH, and ethanol, C2H5OH, are two of the alcohols that we will use in this experiment. You will examine the molecular structure of alkanes and alcohols for the presence and relative strength of two intermolecular forces—hydrogen bonding and dispersion forces. In this experiment, you will Study temperature changes caused by the evaporation of several liquids. Relate the temperature changes to the strength of intermolecular forces of attraction. Figure 1 110 | P a g e Pre-lab exercise: Prior to doing the experiment, complete the Pre-Lab table. The name and formula are given for each compound. Draw a structural formula for a molecule of each compound. Then determine the molecular weight of each of the molecules. Dispersion forces exist between any two molecules, and generally increase as the molecular weight of the molecule increases. Next, examine each molecule for the presence of hydrogen bonding. Before hydrogen bonding can occur, a hydrogen atom must be bonded directly to an N, O, or F atom within the molecule. Tell whether or not each molecule has hydrogen-bonding capability. Procedure: 1. Obtain and wear goggles! CAUTION: The compounds used in this experiment are flammable and poisonous. Avoid inhaling their vapors. Avoid contacting them with your skin or clothing. Be sure there are no open flames in the lab during this experiment. Notify your teacher immediately if an accident occurs. 2. Connect the Temperature Probes to LabQuest. If you have older sensors that do not auto-ID, manually set up the sensors. 3. Wrap Probe 1 and Probe 2 with square pieces of filter paper secured by small rubber bands as shown in Figure 1. Roll the filter paper around the probe tip in the shape of a cylinder. Hint: First slip the rubber band on the probe, wrap the paper around the probe, and then finally slip the rubber band over the paper. The paper should be even with the probe end. 4. Stand Probe 1 in the ethanol container and Probe 2 in the 1-propanol container. Make sure the containers do not tip over. 5. After the probes have been in the liquids for at least 30 seconds, start data collection. Monitor the temperature for 15 seconds to establish the initial temperature of each liquid. Then simultaneously remove the probes from the liquids and tape them so the probe tips extend 5 cm over the edge of the table top as shown in Figure 1. 6. Monitor the temperature of both probes to determine when each probe maintains a constant minimum temperature. Record your minimum temperature value for each probe (t2). 7. For each liquid, subtract the minimum temperature from the maximum temperature to determine t, the temperature change during evaporation. 8. Repeating Steps 3–7 using 1-butanol with Probe 1 and n-pentane with Probe 2. 9. Repeating Steps 3–7, using methanol with Probe 1 and n-hexane with Probe 2. 111 | P a g e Name: ________________________________ Date: ______________________________ Group Members: _______________________ Class & Section: ____________________ PROCESSING THE DATA 1. Two of the liquids, n-pentane and 1-butanol, had nearly the same molecular weights, but significantly different t values. Explain the difference in t values of these substances, based on their intermolecular forces. 2. Which of the alcohols studied has the strongest intermolecular forces of attraction? The weakest intermolecular forces? Explain using the results of this experiment. 3. Which of the alkanes studied has the stronger intermolecular forces of attraction? The weaker intermolecular forces? Explain using the results of this experiment. 4. Plot a graph of t values of the four alcohols versus their respective molecular weights. Plot molecular weight on the horizontal axis and t on the vertical axis. Be sure to label your axis, have title, and since the graph is a linear graph you should have an equation for the line and r2 value. 112 | P a g e PRE-LAB Substance Formula Structural Formulas Molecular Weight Hydrogen Bonding (yes or no) ethanol C2H5OH 1-propanol C3H7OH 1-butanol C4H9OH n-pentane C5H12 methanol CH3OH n-hexane C6H14 DATA TABLE Substance t1 (oC) t2 (oC) t (t1 – t2) (oC) ethanol 1-propanol 1-butanol n-pentane methanol n-hexane 113 | P a g e