CHEM 213- Synthesis of Aspirin Raw Data Sample # m_salicylic_acid V_acetic_anhydride m_product 1 1.99 5.2 4.535 2 2 5 2.332 3 2.004 5 3.952 4 2.243 5 2.855 5 2.03 5 1.671 6 2.038 5 3.696 7 2.082 5.03 2.089 8 2.092 5 6.488 9 2.121 5 3.332 10 1.759 5 1.143 11 2.16 5 1.95 12 2.013 5 2.253 13 2.007 5.02 3.05 14 2.23 5 5.635 15 2.073 5 2.32 16 2.041 5 7.262 17 2.043 5 2.198 18 2.08 5 2.533 19 2.034 5 3.048 20 2.078 5 1.404 21 2.003 5.1 4.522 22 2 5 2.098 23 2.008 5 2.97 24 2.2 5 0.014 *Pick one sample number to use as your own data SYNTHESIS OF ASPIRIN CHEM 213 Fall 2020 Overview ■ Experimental Background ■ Experimental Procedure – Purpose of the lab – Show video ■ Lab Reports – Raw Data – Results – Calculations – Discussion Questions ■ Questions Today’s Reaction Calculations ■ The following calculations must be handwritten/typed as part of your lab report – Moles of salicylic acid – Moles of acetic anhydride – Limiting reagent – Theoretical yield – % Yield Results ■ Follow the data analysis instructions in the lab manual that is on Blackboard – Raw and Processed Class Data ■ Remember all the guidelines from the Excel Tutorial ■ If you have any questions, please email your instructor Discussion ■ Give a thorough interpretation of the results obtained ■ Which reactant, Acetic Anhydride or Salicylic acid, is the limiting reagent? Why is it called so? ■ What is the most likely impurity in the final aspirin product? How is its presence detected? ■ Discuss the results of the Ferric Chloride tests ■ What are the possible sources of error in this experiment? Questions? CHEMISTRY 213, 214, 272 Laboratory Manual 2020 Suzanne W. Slayden and Faculty in the Department of Chemistry & Biochemistry George Mason University -i- 2020 CHEM 213/214/272 Laboratory Manual Summer 2020 George Mason University -ii- Table of Contents Preface ............................................................................................................................................................................................................... iv General Safety Guidelines and Operations in the General Chemistry Labs ............................................................................v KEEPING A LABORATORY NOTEBOOK ............................................................................................................................................... vii CHEM 213 / CHEM 272 EXPERIMENTS ................................................................................................................................................ 1 Experiment 1 Pennies and good Sense ...................................................................................................................... 2 Experiment 2 Measurements ......................................................................................................................................... 7 Experiment 3 Qualitative Analysis of Cations ....................................................................................................... 19 Experiment 4 Density of Solutions ............................................................................................................................ 24 Experiment 5 Empirical Formula of Zinc Iodide ................................................................................................. 36 Experiment 6 Synthesis of Aspirin ............................................................................................................................ 43 Experiment 7 Volumetric Analysis: Titration of Vinegar ................................................................................. 52 Experiment 8 The Ideal Gas Law ................................................................................................................................ 67 Experiment 9 Heat of Reaction: Hess’s Law .......................................................................................................... 80 Experiment 10 Solvent Polarities from a Solvatochromic Dye...................................................................... 94 Experiment 11 VSEPR and Molecular Modeling ............................................................................................... 103 Experiment 12 Estimation of Avogadro’s Number .......................................................................................... 111 Experiment 13 Graphing ............................................................................................................................................ 114 Experiment 14 Thin-Layer Chromatography..................................................................................................... 118 Experiment 15 Chemical Interactions .................................................................................................................. 124 Experiment 16 Absorption Spectroscopy............................................................................................................ 132 Experiment 17 Periodic Properties: Cation Analysis...................................................................................... 148 Experiment 18 Periodic Properties: Anion Analysis....................................................................................... 156 CHEM 214, CHEM 272 EXPERIMENTS ............................................................................................................................................. 160 Experiment 19 Synthesis of a Coordination Compound ............................................................................... 161 Experiment 20 Determination of Water of Hydration in KTOF3 ............................................................... 170 Experiment 21 Redox Titration of Oxalate in KTOF3 ..................................................................................... 176 Experiment 22 Determination of Molecular Mass by Freezing Point Depression.............................. 184 Experiment 23 Chemical Kinetics ........................................................................................................................... 191 Experiment 24 Measurement of an Equilibrium Constant ........................................................................... 200 Experiment 25 Acidity Constant from a Titration Curve .............................................................................. 211 Experiment 26 Dissociation Constants of Acids and Bases .......................................................................... 217 Experiment 27 Solubility Product Constant (Ksp) and Common-Ion Effect .......................................... 226 -iii- Experiment 28 Thermodynamics of the Dissolution of Borax .................................................................... 235 Experiment 29 .Electrochemical Measurements .............................................................................................. 242 Experiment 30 Identification of an Organic Acid ............................................................................................. 250 APPENDIX ...................................................................................................................................................................................................... 255 A 1. Instructions for the pH probe ................................................................................................................... 257 A 2. Temperature probe ........................................................................................................................................ 261 A 3. Spectro Vis instructions................................................................................................................................ 263 PREFACE The experiments is in this lab manual are designed for students that are currently enrolled in a science / engineering program at George Mason University. Overall, there are two categories of experiments: those that are meant to complement a lecture discussion on the same topic and those for which there are no relevant lecture text commentaries but which exemplify experimental principles and techniques. Visible-region spectroscopy, which is emphasized throughout the lab courses, is the precursor to the study of ultraviolet, infrared, and NMR spectroscopies covered in the organic chemistry laboratory courses. Organic compounds also figure prominently here because rarely is there time in lecture to survey organic chemistry and because biology majors (our main constituency) are exposed early to organic and biochemical structures in their biology studies. As might be expected for a new edition, many people have contributed toward its improvement. GMU chemistry professors Gerald Weatherspoon, James Schornick and John Schreifels suggested procedures which were gratefully received. Most consequential have been those students (graduate teaching assistants Arion Ross, Yukiko Yarnall, Allison Leninger, Megan Devine, and Christopher Kenedy) who offered constructive comments from their practical experience with the previous edition. -iv- GENERAL SAFETY GUIDELINES AND OPERATIONS IN THE GENERAL CHEMISTRY LABS 1. Safety goggles must be worn in the laboratory at all times, whether you are actually performing an experiment or not. (The wearing of contact lenses is strongly discouraged.) 2. In case of fire, accidents or injury (no matter how minor) notify the lab instructor at once. 3. Note the location of fire extinguishers, safety showers, eyewash and blankets so that you can use them quickly in an emergency. 4. If either a corrosive acid, base, or any other chemical comes in contact with your skin or clothing, wash the affected area immediately with copious amounts of water. Notify the instructor. 5. Always read a label carefully before removing a chemical from its container. 6. Exercise great care in noting the odor of fumes, and avoid breathing in fumes of any kind. 7. Never point a test tube or other vessel containing a reacting mixture (especially when you are heating it) toward your neighbor or yourself. 8. NO SMOKING, EATING, DRINKING, OR CHEWING OF GUM in the laboratory. Do not taste anything in the lab. (This applies to cigarettes, food, drinks and gum as well as chemicals.) 9. Under no circumstances should you perform any unauthorized experiments. 10. Keep aisles and benches clear of books, clothing, etc. Store unused belongings in the front area of the lab or in the cabinets provided. Do not use lab benches as seats. 11. Work areas are to be cleaned and put back in order prior to exiting the lab (includes electronic balance area, fume hoods, etc.) 12. Full and proper attire is required for all students entering the teaching labs---including proper shoes. 13. Students are required to wear FULL SHOES----no abbreviated shoes means no open-toed or open-heeled shoes. Ballet flats, flats,pumps and heels, wedges (including wedge boots), and moccasins with socks are not acceptable for lab entry nor any other shoes that do not completely cover the top of the students' feet since they do not offer enough protection to the body. Flip flops, sandals, top-sider/boat shoes, espadrilles, and loafersare not acceptable forms of protective shoes for entering the chemistry lab. 14. Cotton fabric shoes are absorbent and do not provide the necessary protection from chemical spills. -v- 15. Low-cut socks are unacceptable for lab entry. Full socks (crew length) must be worn at all times. 16. Tights/leggings are not acceptable as "pants". Pants worn to lab should allow enough space so that chemicals coming in contact with the body will not be held tightly to the body resulting in additional harm. 17. Skinny jeans/pants are not allowed in the chemistry labs. 18. Yoga pants and running tights are not allowed in the chemistry labs. 19. Shorts, capri pants and any other form of abbreviated lower body wear are not allowed. Long pants worn must completely reach your shoes and not leave any area of the legs or feet exposed. 20. Skirt lengths must adhere to the same requirements as pants; skirts should extend all the way to the shoes and not leave any areas of the legs or feet exposed. Skirts that do not meet this criterion are forbidden in the chemistry lab. 21. Always keep long hair confined. 22. Tank tops and abbreviated upper body wear are not allowed in the chemistry labs. As an additional layer of protective wear, the upper body should be covered by short-sleeve clothing. 23. Lab coats are provided by the chemical preps room OR may be purchased directly from the GMU Bookstore. Lab coats purchased individually must be knee length; waist length laboratory coats are not allowed in the laboratory. 24. Children, acquaintances and pets are not allowed in the lab. 25. Cell phones, laptops and other portable electronic devices, may not be used in the laboratory. Cell phones should be stored in backpacks prior to entering lab and should not be removed until the student exits the lab. Violation will result in a loss of 40 points for the lab exercise being performed that day. 26. Students that refuse to wear the proper protective equipment (PPE) will be dismissed from the lab for that day and a grade of zero recorded for the laboratory exercise. Additional information regarding PPE violations can be found on Blackboard. 27. Students will receive one general warning about PPE violations, which is given during the first class meeting. The second violation will result in a loss of 40 points for the lab exercise being performed that day. Repeat violators will be referred to the General Chemistry Laboratory Coordinator, with the recommendation that the student be permanently removed from the lab course. 28. Refer to the lab syllabus (posted on Blackboard) regarding laboratory operations. 29. You will sign an acknowledgement sheet on the first day of the lab, stating that we went over these rules and regulations and you read them on Blackboard too. -vi- KEEPING A LABORATORY NOTEBOOK General Instructions 1. Students enrolling in General Chemistry lab courses (CHEM213/ CHEM214 / CHEM251) are required to purchase the carbonless spiral bound notebook from the George Mason University Bookstore. 2. Put your name and your instructor’s name on both the outside and inside of the notebook. This will prove useful in case the lab notebook is misplaced. 3. Leave the first page blank, except for writing the word "Index" at the top of the page. The Index will be a chronological listing of the experiment numbers, the experiment titles and their page numbers. 4. Make all notebook entries in PERMANENT INK using a blue or black ballpoint pen. Use the right-hand pages (white pages with light blue grid lines) for main entries and the left-hand pages (yellow backside of the carbon pages) for scratch work only. 5. Do not try to erase or white-out; draw a single line through the incorrect work or entry. 6. Maintain chronological entries; do not leave blank pages. Each main page must be dated at the top of the page. 7. All data and observations are to be recorded directly in your lab notebook. Never record data or observations anywhere except in your notebook. 8. You will turn in the carbon copy of the raw data at the end of each formal lab exercise, retaining the notebook for yourself. Additional details will be provided on the first day of lab and posted on Blackboard. Outline of Notebook Format Generally, you will find that each experiment can be organized in your notebook under the headings: Title; References; Purpose; Apparatus, Equipment and/or Chemicals; Procedure and Observations; Data and Results (includes sample calculations); Discussion; and Conclusion. Some of these terms are explained as follows: References: This states the source, including page numbers, of the information used to perform the experiment, usually the laboratory manual. Your textbook and other sources also should be read and referred to. -vii- Purpose: This should be a brief simple statement of the intent of the experiment. It should be written so that later, when you reread it, you have no difficulty remembering the experiment. Usually, the manual contains such a statement and you may choose to adapt that wording. Apparatus, Equipment, Chemicals: All important materials used in the experiment should be listed. If any new equipment is used, it should be described with a drawing which includes essential features such as approximate dimensions and relative locations. If apparatus described previously is used, the original description should be cited, i.e., "p. 37". Procedure and Observations: This section is much like a diary and should describe what you did and what you saw. The procedural steps should be written, as a part of the pre-lab, prior to entering the lab. The observations should be entered in the notebook at the time the operation is performed, not before. Do not simply write the same instructions which are in the manual. Number the steps so that it is easy for you to keep track of where you are in the experiment. Make sure that the steps are clear and easy for you to understand (no praragraphs). Do not write in first person. Never record on a pad of paper or on scrap paper and then transcribe into the notebook. Always record directly in the notebook. The notebook is not the spot for polished writing. Instead, be as detailed as possible so that someone else can duplicate what you have done by reading your account. This means including contradictions and failures. If there is a conflict, enter a description of it rather than omitting it. The negative results may be important later for figuring out what happened. Data should be entered directly into the lab notebook as it is gathered before performing calculations. Leave adequate room for observations and, where necessary, calculations. Do not write in a jumbled/congested and disorganized manner. Refer to the grading rubric (see Blackboard). Data and Results: This section must begin on a new page, titled "Data and Results". You may wish to model your Data Sheet on the sample report sheets found in your lab manual. The quantitative data, as well as qualitative data, is to be organized, neat, and recorded with the appropriate significant figures and units. Data must be written legibly. Discussion: In this section, you will address the discussion questions from the lab manual or posted on Blackboard. In addition, your individual experimental outcomes, data analysis and error analysis will be addressed and discussed as it relates to outcomes for the entire class. Proper sentence structure and grammar are to be used at all times. Conclusion: -viii- In this section you should record any thoughts or conclusions about your experiment. You may offer explanations for low yields for example, or for improvements in procedure or you may comment on the experiment itself. ***A sample lab report from a previous semester is posted on Blackboard. The report is to be used as a guide----not for plagiarism purposes. "A contemporary instance of the importance of laboratory notebooks is provided by the case of Gordon Gould, who as a young physicist filed an application for a basic laser patent in 1959. Gould failed to get the patent, which was awarded instead in 1960 to Charles Townes and Arthur Schawlow. Gould went to court, claiming he was the true inventor. His challenge was based in part on his research notebook which showed, among other items, a sketch, a statement of the main idea, and a derivation of the acronym LASER--Light Amplification by Stimulated Emission of Radiation. "In October 1977, after a series of litigated oppositions, Gould was granted a patent for optically pumped laser amplifiers. The world market has been estimated at between 100 million and 200 million dollars. "As this example suggests, well-kept notebooks are valuable documents. They provide complete, accurate records of ongoing work. In the event of litigation or contests for patent rights, they are submitted as evidence. They serve the important role of corroboration should the researcher or inventor have to prove origin or substantiate statements and conclusions. They are valuable documents to validate a company's claims to funds spent for research, particularly in support of tax deductions. "The uses of laboratory notebooks are not limited to legal issues. They are vehicles for organizing and focusing the thinking of the writer, as well as being receptacles for detailed procedural information that might not be available in highly compressed journal articles. Finally, they may serve not only the researcher or inventor but also the public. If properly maintained, they are a record of success and failure, a safeguard against error and carelessness in such important areas as the testing of drugs and chemicals."1 1 "Keeping a Laboratory Notebook", A. Eisenberg; J. Chem. Educ., 59 (12), 1045 (1982). -ix- CHEM 213 / CHEM 272 EXPERIMENTS . -1- Experiment 1 PENNIES AND GOOD SENSE Adapted from Reference: "An Experiment in Thinking Scientifically", D. J. Sardella, J. Chem. Ed. 69, 933 (1992). Purpose: The purpose of this exercise is to determine experimentally the variation in mass of pennies both within a single mint year and over a range of years. Equipment: Electronic balance, pennies Procedure and Observations: Obtain a vial of pennies from your instructor and record the mint year. Using a single balance, determine the precise mass (in grams) of each of your pennies (up to a maximum of 10) by placing them on the balance and weighing them using a weigh boat. Make sure to zero the balance with the weigh boat before measuring the mass of each penny. Record all your data in your laboratory notebook as you weigh the pennies. Be sure to record the mass to the nearest 0.001 g, which is the limit of precision of the balance. Make sure to observe the penny’s condition – corroded, dirty, etc. Data: After you have recorded all the masses, enter your own data into the appropriate Web-based data entry form, using one of the computers in the lab. Setup a table in your lab notebook that has columns for masses, deviations, and deviations squared. Calculate the average mass and standard deviation for the mint year of pennies assigned by your lab instructor. Verify the calculated standard deviation value with your lab instructor prior to entering data into the database. Results: Outside of class, retrieve the class data in spreadsheet form from: http://chem.gmu.edu/results/. Use Excel to calculate the average mass and the standard deviation for each mint year, using the AVERAGE() and STDEV() spreadsheet functions. Make sure the precision of the calculations matches the precision of the data. Finally, calculate the standard deviation for ALL the pennies, considered as a single large sample. Use Excel to prepare a scatter plot (select CHART, then options) showing the average mass vs. year on the primary y-axis and standard deviation vs. mint year on the secondary y-axis. A sample is provided on the next page. Row and column labels should always be printed for each Excel sheet. Examples of a data spreadsheet and charts are shown following this page. Discussion: In your discussion section address the following questions: -2- • Describe the trend observed in the variation of average mass with mint year. • Describe the trend observed in the variation of the standard deviation with mint year. • How does the standard deviation among the pennies in a given year compare with the standard deviation among all the pennies (all years)? What can you conclude from this? • Give a possible explanation for these trends and discuss how your observations support this explanation. Include a discussion of the contribution of the density of each (look it up for each metal) and the composition of each penny to the total mass of the pennies. -3- Chemistry 213-2A1 Measurements I: Sample Sheet of Data and Results Mint year Avg. mass1 mass2 mass3 mass4 mass5 mass6 mass7 mass8 mass9 mass10 Mass number (grams) (grams) (grams) (grams) (grams) (grams) (grams) (grams) (grams) (grams) (grams) 1971 9 3.041 3.048 3.064 3.090 3.112 2.966 3.067 0.045 1972 6 3.088 3.029 3.14 3.046 3.075 3.090 3.078 0.039 1974 10 3.132 3.072 3.105 3.090 3.039 3.110 3.096 3.113 3.128 3.084 3.097 0.028 1975 10 3.089 3.098 3.043 3.067 3.032 3.053 3.098 3.092 3.123 3.078 3.077 0.028 1976 7 3.111 3.155 3.084 3.098 3.127 3.053 3.080 3.101 0.034 1977 8 3.080 3.072 3.011 3.085 3.094 3.128 3.090 3.049 3.076 0.034 1978 10 3.143 3.068 3.056 3.131 3.107 3.074 3.117 3.054 3.093 3.102 3.095 0.031 1979 10 3.095 3.130 3.135 3.073 3.056 3.053 3.117 3.091 3.063 3.126 3.094 0.032 1980 10 3.094 3.096 3.078 3.094 3.058 3.131 3.076 3.065 3.124 3.094 3.091 0.023 1981 10 3.072 3.083 3.100 2.999 3.083 3.074 3.108 3.102 3.115 3.064 3.080 0.033 1982 10 3.020 3.092 3.040 3.086 3.088 3.062 3.028 3.105 3.084 3.055 3.066 0.029 1983 10 2.509 2.517 2.586 2.535 2.543 2.519 2.517 2.496 2.514 2.514 2.525 0.025 1983 8 2.545 2.550 2.492 2.486 2.527 2.539 2.580 2.518 2.530 0.031 1984 10 2.529 2.529 2.553 2.530 2.464 2.532 2.522 2.582 2.529 0.029 -4- 3.083 3.089 3.108 Std. Dev. 2.519 2.525 1985 10 2.534 2.579 2.507 2.575 2.525 2.501 2.502 2.543 2.492 2.511 2.527 0.031 1986 10 2.485 2.561 2.490 2.520 2.569 2.494 2.614 2.482 2.557 2.560 2.533 0.045 1987 10 2.511 2.545 2.468 2.508 2.503 2.496 2.521 2.465 2.485 2.500 2.500 0.024 1988 10 2.544 2.504 2.479 2.470 2.516 2.473 2.505 2.492 2.518 2.486 2.499 0.023 1990 10 2.485 2.502 2.539 2.505 2.458 2.515 2.487 2.519 2.499 2.498 2.501 0.022 Shaded areas indicate where measurements were not performed. -5- -6- Experiment 2 MEASUREMENTS In all sciences, careful measurements of materials and conditions are essential. There must be well-understood, consistent methods of describing the weight, volume, temperature, and other characteristics of a sample or method that is investigated. The metric system is the accepted system for scientific measurements worldwide and will be employed throughout your chemistry lecture and laboratory. Tables listing the relationships between the English and metric systems of measurement can be found in your textbook and other reference works. You should be familiar with the metric prefixes and with the units used in measuring mass, volume, temperature, and distance. It is never possible to make an exact measurement because no matter how sophisticated the measuring instrument or how carefully the measurement is made, errors cannot be avoided. We distinguish between two categories of experimental errors: systematic and random. It is possible to eliminate or correct systematic errors; we cannot eliminate random errors. Random errors are the errors, which are inherent in every physical measurement and are beyond the observer's control. Systematic errors often are avoidable and may arise from faulty equipment or poor technique. (However, a systematic error is not the same as a mistake or a blunder.) In order to balance random errors and more closely to approximate the accepted or true value, at least two, and preferably three or more, measurements should be made. The best value is usually the average of the individual results (although often the question arises whether to include a result that seems out of line with the others). The arithmetic mean, or average, of this set of results is calculated by the equation: Mean = Average = x ave = [ x1 + x2 + x3 +... xi ] n where X1,2,3,...,i indicates the measured values in the sample and n is the number of measurements in the sample. The error of a measurement is the numerical difference between the observed value and the "true" or accepted value. The smaller the error, the more accurate the observed value. Thus, accuracy is the nearness of a result to the true value. If the true or accepted value is known, the average and/or individual values can be compared to it and the accuracy judged. A useful comparison is the percent error: % Error = accepted value − observed value accepted value -7- •100 Because the “true” value is seldom known (and hence the accuracy), we often rely on the precision of the data as an indication of its accuracy. Precision is the closeness or agreement of a set of results among themselves. It is thus an indication of the reproducibility or reliability of the measurements. Precise results may be accurate or inaccurate; that is, measurements of high precision, made with a faulty measuring instrument, can be wrong. We often speak of the uncertainty of the measurement using a particular piece of equipment. Since no measurement, or derived quantity, can be exactly determined, the number of significant figures used in recording its magnitude expresses the uncertainty of a measurement. All certain digits are recorded, while the final digit is significant but estimated, or uncertain. (Consult your textbook for a discussion on the correct use of significant figures.) For example, some electronic balances weigh an object with an uncertainty of + 0.001 g (plus or minus 1 milligram), while other, less precise balances weigh only to + 0.01 g (plus or minus 1 centigram). If the digital display on a milligram electronic balance reads 10.382 g, then the sample’s maximum mass is 10.383 g and its minimum mass is 10.381 g (10.382 + 0.001 g). Therefore, the measurement and its uncertainty indicate the range a data value may have because of the need to estimate the last, uncertain digit. If the final digit of a display is a zero, such as 8.410 g, the zero must be recorded to indicate that the uncertainty is in the milligram range. Dropping the zero to give 8.41 g would indicate an uncertainty of only 0.01 g. For equipment which does not display a digital value, estimate the uncertainty by how well you can read the graduation markings on the equipment. In many cases, you can assume that the last digit is uncertain by + 1 unit. For example, suppose the scale below represents graduations of 1 cm and you have determined the length of an object to be at the mark L. You are certain that the length is greater than 2 cm, but you are uncertain about its exact location between 2 and 3 cm; that is, you are uncertain about the final digit. Mentally divide the distance between 2 and 3 cm into tenths (0.1), and then if you estimate the location to be about 7/10's along the line, the length is 2.7 + 0.1 cm. 1 2 3 4 5 ^ L The scale below is similar to the one above, but suppose this one is further subdivided into millimeters. 1 2 3 4 ^ L -8- 5 Now as you look at the mark L, you are certain that it is at least 3.5 cm, but again, you are uncertain about the final digit. Mentally divide the distance between 3.5 and 3.6 cm into tenths, (0.01 cm = 0.1 mm) and estimate the position of L. Your answer could be 3.52 cm but others might decide 3.51 or 3.53 cm. The measurement you would record is 3.52 + 0.01 cm. Notice that the closer the graduations, the more difficult it is to use an uncertainty of + 1 unit. An easier division to estimate might have been 0.2 or 0.5 cm. With some practice, you can decide what is the most appropriate division. Relative uncertainty is the ratio of the uncertainty to the magnitude of the measurement. It is frequently expressed as a percent. % relative uncertainty = uncertainty •100 magnitude of measurement Notice that with a constant uncertainty, the relative uncertainty becomes smaller as the magnitude of the measurement becomes larger. The precision of a group of measurements is often expressed in terms of deviation (d). The deviation of a result is the numerical difference between an experimental value and the average value, (xi — xavg). If the deviations for a set of results are small, as compared to the average value, then the measurement is precise. Although we cannot always express the accuracy of a measurement (because we may not know the true value), we can always express the precision of several measurements. In the statistical treatment of experimental data, a common measure of precision is the standard deviation, x, which characterizes the average uncertainty in the individual measurements. The standard deviation for a small sample is the “square root of the average of the squared deviations” and can be calculated as follows: (𝑑12 + 𝑑22 + 𝑑32 + ⋯ . +𝑑𝑖2 ) 𝜎𝑥 = √ (𝑛 − 1) where the series is the sum of the squared deviations, di2 =(xi — xavg)2, and n is the number of items in the sample. (For theoretical reasons, the value [n-1] is used rather than n.) The standard deviation, x, is expressed in the same units as the data. (See the sample calculation.) A small standard deviation indicates high precision and a large standard deviation indicates low precision. Generally, high precision indicates high accuracy, though this is not necessarily so. For example, when a constant procedural error is repeated in every measurement, high precision may be observed even though the result is not an accurate determination of the true value. If we assume that most measurements are corrected to eliminate systematic errors and that random errors are small and as likely to be positive as to be negative, then a graph of the distribution of many measurements approaches a bell-shaped curve. (See Figure 1.) Such a graph can be generated by plotting the value of a measurement versus the number of times (frequency) that particular value is obtained. -9- The average value is found at the peak maximum and is considered the most probable value (more likely to occur frequently). Those values which depart widely from the average occur less frequently than those for which the deviation is small. If most of the measured values cluster closely around the mean, the curve will be narrow and steep and the measurements will have a high precision. If the measured values are widely dispersed from the mean, the curve will be broad and the measurements imprecise. Inspecting Figure 1, you will notice that although both curves have the same mean (and thus the same accuracy), one set of results is much more precise than the other. This is a very effective depiction of the difference between accuracy and precision. Number of times the value occurred Figure 1 Value Often one value in a group of results differs significantly from the others. In such cases, it is likely that some inadvertent error was made in a measurement. There is a simple rule that you may use in deciding whether a value should be discarded. If the deviation of the suspected value from the mean (determined without the suspected value) is equal to three times (or more) the average deviation from the mean of the other values, the suspected value should be discarded. ***Some of the common laboratory equipment you will use this year appear in Figure 2. Information on expected accuracy and uncertainty is provided in Table 1. -10- Figure 2 -11- TABLE 1 Instrument Graduation Uncertainty centigram 0.01 g 0.01 g milligram 0.001 g 0.001 g Top-loading electronic balance Graduated cylinder 10-mL 0.1 mL 100-mL 1 mL Buret, 50-mL 0.1 mL Volumetric flask not graduated -- 25-mL 0.05 mL fill to mark 50-mL 100-mL Pipet, volumetric, 10-mL not graduated -fill to mark 110° thermometer 1° C 0.2° C Mercury barometer 0.1 mm 0.2 mm -12- Calculating a Standard Deviation Sum = Measured Value Deviation Deviation Squared xi di=(xi - xavg) di2=(xi - xavg)2 5.21 +0.01 0.0001 5.18 - 0.02 0.0004 5.23 +0.03 0.0009 5.20 0.00 0 5.17 - 0.03 0.0009 25.99 𝐴𝑣𝑒𝑟𝑎𝑔𝑒 = 𝑥̅ = 𝑥𝑎𝑣𝑔 = 0.0023 ∑(𝑥1 + ⋯ 𝑥𝑖 ) 25.99 = = 5.20 𝑛 5 (0.0023) (𝑑12 + 𝑑22 + 𝑑32 + ⋯ . +𝑑𝑖2 ) 𝑆𝑡𝑑. 𝐷𝑒𝑣. = 𝜎𝑥 = √ =√ = 0.02 (𝑛 − 1) (5 − 1) Problems (show work) 1. Explain why it is correct to speak of the uncertainty of a single result but not the precision of a single result. Is it correct to speak of the accuracy of a single result? 2. What are the maximum and minimum values for a mass which has been reported as 7.23 ± 0.05 g? 3. Show that for an accuracy of 1% in weighing a 50 g sample, a balance which is uncertain to only 0.5 g is required. 4. Calculate the percent error in a molecular weight determination if the experiment yielded values of 121.5, 122.3, 121.9 and the actual molecular weight is 122.1. 5. The following volumes were measured in a 100-mL graduated cylinder: 74.6, 75.2, 73.7, 74.2, 75.0, 73.9. Calculate the average and the standard deviation for the measurements. 6. If your calculator has a built-in statistics program, verify the calculations in the Example and in Problem #6. . 7. Suppose the following percentages were determined for the amount of water in a crystal hydrate: 36.20%, 36.52%, 35.50%, and 35.95%. Determine if one of the values is suspect and may be discarded. Answers to Selected Problems: #4. 0.2% #5. 74.4, 0.6 -13- THE EXPERIMENT It is important that you learn how to use common equipment correctly in order that your measurements in the laboratory are as accurate and precise as possible. The validity of your experiment depends on the reliability of your measurements. You will use the equipment from this laboratory exercise to perform your experiments throughout the year. At the beginning of the laboratory period, your instructor will demonstrate the proper use and function of this equipment. During the experiment, you will be expected to acquire certain skills and techniques as you use the equipment. Today you will practice making mass and volume measurements. Mass measurements will be made on an electronic balance and their precision assessed. You should make all of your mass measurements on the same balance. Volume will be measured with a variety of glassware -- graduated cylinder, pipet, and buret. The accuracy of these devices will be determined by comparing a measured volume with the calculated volume. The conversion factor, which relates mass and volume, density, is expressed as mass per unit volume. At room temperature, one gram of water occupies approximately one milliliter of volume (that is, the density of water = 1 g/mL). PROCEDURE You will need a calculator for this experiment. It may be helpful to review the built-in standard deviation program on your calculator. Although such a program is not necessary for the experiment, it will make your calculations of standard deviation easier. Also, make sure you understand rounding of numbers and the determination of significant figures as explained in your textbook. Before you begin, place approximately 250 mL of distilled water in a large beaker and allow it to come to room temperature. All of your volume measurements should be done with this water. Place a thermometer in the water and record its temperature. Record the temperature of the air in the room. Measuring Mass 1. Weigh a dry 25-mL graduated cylinder on an electronic balance. Record the uncertainty of the measured mass. Repeat twice, zeroing the balance before each trial. Determine the average mass. Calculate the standard deviation. 2. Weigh a dry 150-mL beaker on an electronic balance. Record the uncertainty, and then determine the average mass and the standard deviation. Assess the precision of the balance. Based on your triplicate mass measurements, how many times do you think you should measure mass in future experiments? Assessing the Accuracy of Glassware Your results from the preceding section probably showed that mass measurements using the same balance are very precise (and presumably accurate because the balances have been calibrated with known masses). We next want to answer the question "which glassware most -14- accurately measures the volume of a given amount of water?" To determine this, you will use the mass of water as a basis for comparing and judging the accuracy of the volumes of water you measured. The conversion factor, which relates mass and volume, is the density. The density of water at various temperatures has been accurately determined in other laboratories. Alternatively, you may use the formula: density = 1.0028 − (0.000228)T where T is in oC and the density is in g/mL. 1. Into the same graduated cylinder you used in the Measuring Mass section above, add about 10 mL of distilled water at room temperature. Do not attempt to adjust the volume in the graduated cylinder to exactly 10.0 mL; this may introduce error and will take too long. Record the volume of water in the graduated cylinder (this is the measured volume; it should not be exactly 10 mL.). Weigh the graduated cylinder containing the water and determine the mass of water in the graduated cylinder. 2. Set up a clean buret and rinse with distilled water. Fill the buret with distilled water and drain enough water through the tip to flush out any air bubbles. Bring the water level to below the 0.0-mL mark. Record the initial volume reading. Do not attempt to adjust the initial volume to the 0.00 mL mark. This is unnecessary and is too time consuming. Then drain about 10 mL from the buret into the 150-mL dry beaker you used in the previous section. Record the final volume reading from the buret to obtain the measured volume of water delivered from the buret. Weigh the beaker and water to determine the mass of water delivered from the buret. From the known density of water at today’s room temperature and the measured mass of water delivered from each piece of glassware, determine the calculated volume of water delivered from each. Finally, calculate the error and the percent error in the measured volume, taking the calculated volume as the accepted/true value. Volume (mL) = mass (g) density (g/mL) Compare the accuracy of the three types of glassware. What are the advantages and disadvantages of the three types of glassware? -15- Temperature (°C) Density of water (g/mL) 20 0.99823 21 0.99800 22 0.99777 23 0.99767 24 0.99733 25 0.99708 -16- ___________________________________________________________________________________________________ MEASUREMENTS ___________________________________________________________________________________________________ MEASURING MASS Mass of 10-mL graduated cylinder Mass of 150-mL beaker __________ (g) __________ (g) Avg. Mass __________ (g) Std. Deviation _______ __________ (g) __________ (g) Avg. Mass __________ (g) Std. Deviation _______ __________ (g) Uncertainty + _______ __________ (g) Uncertainty + _______ ***Sample data table provided above. The table in your lab notebook should follow the same format. -17- Assessment of Precision: GLASSWARE ACCURACY Temperature of room air: __________ Temperature of water: __________ Density of water (from table or calculation) ___________ REQUIRED Sample Calculations for the following: Graduated Cylinder: Measured Vol. of H2O Mass of Cylinder + H2O Mass of H2O Calculated Vol. of H2O Error Percent Error Initial Vol. in Buret Final Vol. in Buret Measured Vol. of H2O Mass of Beaker + H2O Mass of H2O Percent Error Calculated Vol. of H2O Error Percent Error Buret: Assessment of Accuracy: -18- Experiment 3 QUALITATIVE ANALYSIS OF CATIONS Introduction: This experiment will analyze known and unknown solutions for the presence of specific cations. The general approach to finding out what ions are in a solution is to test for the presence of each possible component by adding a reagent that will cause that component, if present, to react in a certain way. This method involves a series of tests, one for each component, carried out on separate samples of solution. Difficulty sometimes arises, particularly in complex mixtures, because one of the species may interfere with the analytical test for another. Although interferences are common, many ions in mixtures can usually be identified by simple tests. In this experiment, you will analyze unknown solutions that may contain one or more of the following ions in solution: Fe3+, Ni2+, Mg2+, Cu2+ First, you will perform various tests to detect the presence of individual ions. Once you have observed these specific reactions, you will perform the same tests on a series of unknown solutions. Safety Precautions: In some of the tests, you will be required to use fairly concentrated acids and bases. When in contact with skin, most of these chemicals can cause severe burns if not removed promptly. Wear goggles when working with any of the reagents required in this experiment. Experimental Set-Up: For each experimental procedure, arrange four test tubes, so that the test tubes are organized like the data tables. The first row of test tubes corresponds to row 1 on the data tables. The first column of the test tubes corresponds to column A on the data tables. Procedure for Cation Tests: -19- Controls - First, you will test known solutions and observe the results (rows 1-4). For each cation there will be a unique set of positive reactions that indicate the presence of that cation. The test solutions for Cation are sodium iodide in acetone (NaI), sodium ferrocyanide (Na4Fe(CN)6), dimethylglyoxime in ethanol/water solution, and sodium hydroxide (NaOH). 1. Place 20 drops of the first cation solution (Fe(NO3)3 ) in the first 4 test tubes. 2. Repeat for test tubes 5-16 using the appropriate cation solution. 3. Place 4 drops of the first test solution (NaI) in the first test tube of each cation solution. 4. Repeat for test tubes 5-16 using the appropriate test solution. 5. Observe and record any reactions in the data table. Be sure to record the color and amount of any precipitates. Unknowns 6. Identify the cations in the unknown solutions in the vial. 7. Repeat steps 3-5 for the unknown vial to identify unknown cation solution. 8. Clean Up – carefully pour all of the solutions from the test tube into the waste container, then clean and dry the test tubes. If needed, rinse the test tube with soap or acetone to remove residue. -20- Name_______________________ Section___________ Lab Partner_____________ Data Tables: Cation Identification (descriptions will vary) 1 0.25 M Fe(NO3)3 2 0.1 M Ni(NO3)2 3 0.1 M Mg(NO3)2 4 0.1 M Cu(NO3)2 A B D E 7.5 % NaI 0.25M Na4Fe(CN)6 0.25 M Dimethylglyoxime 1 M NaOH 1. Balance and classify the following reactions: combination (C), decomposition (D), single replacement (SR), double replacement (DR) or combustion (CB). Classification a. ____C3H8 (g) + ____O2 (g) → ____CO2 (g) + ____H2O (l) ____________ b. ____K (s) + ____H2O (l) → ____KOH (aq) + ___H2 (g) ____________ c. ____MgSO4 (aq) + ____Na2CO3 (aq) → ____MgCO3 (s) + ____Na2SO4 (aq) ____________ d. ____Cl2 (g) + ____NaI (s) → ____NaCl (s) + ____I2 (s) -21- e. ____NaNO3 (s) → ____NaNO2 (s) + ____O2 (g) ________ f. ________ ____H2 (g) + ____Cl2 (g) → ____HCl (g) g. ____Ca(C2H3O2)2 (aq) + ____H3PO4 (aq) → ____Ca3(PO4)2 (s) + ____HC2H3O2 (aq) __________ h. ____H2O2 (l) → ____H2O (l) + ____O2 (g) _________ i. ____Na (s) + ____S (s) → ____Na2S (s) _________ j. ____BaCl2 (aq) + ____K2CrO4 (aq) → ____BaCrO4 (s) + ____KCl (aq) ________ 2. Balance each of the following equations. Then write the total (complete) ionic equation for each reaction. Finally write the net ionic equation for each reaction. Remember: insoluble substances are not present as separate ions in solution. a. ____Pb(NO3)2 (aq) + ____NaOH (aq) → ____Pb(OH)2 (s) + ____NaNO3 (aq) Complete Ionic equation: Net ionic equation: b. ____Na3PO4 (aq) + ____Ca(NO3)2 (aq) → ____Ca3(PO4)2 (s) + ____NaNO3 (aq) Complete Ionic equation: Net ionic equation: c. ____NaCl (aq) + ____AgNO3 (aq) → ____AgCl (s) + ____NaNO3 (aq) Complete Ionic equation: Net ionic equation: d. ____Na2SO4 (aq) + ____AgNO3 (aq) → ____Ag2SO4 (s) + ____NaNO3 (aq) Complete Ionic equation: Net ionic equation: -22- e. ____Na2CO3 (aq) + ____BaCl2 (aq) → ____BaCO3 (s) + ____NaCl (aq) Complete Ionic equation: Net ionic equation: f. ____Na2CO3 (aq) + ____HCl (aq) → ____NaCl (aq) + ____H2O (l) + ____CO2 (g) Complete Ionic equation: Net ionic equation: -23- Experiment 4 DENSITY OF SOLUTIONS Density is a physical property of a substance; it relates the mass of the substance to the volume it occupies and it is thus characteristic of the substance. Densities can be measured for gases, liquids, and solids and are expressed as mass per unit volume. Gas density is usually expressed as g/liter, while for the more dense liquids, the density is expressed as g/mL. Densities of many pure substances have been determined and you can look them up in reference books. For example, the density of water at 4°C is 1.000 g/mL (density changes somewhat with temperature, so it is necessary to specify the temperature to which the density refers). To understand how the densities of substances can vary widely, imagine two trucks, each with the same volume of cargo space. One truck is filled with lead shot, the other with packing peanuts. Although the volumes occupied are identical, the truck filled with lead will be much heavier than the truck filled with packing peanuts. Thus, the density of lead is much greater than the density of packing peanuts since lead has the larger mass-to-volume ratio. Solutions are homogeneous mixtures of two or more substances in which the boundaries of the component substances can no longer be identified. A common solution is that of a solid-in-liquid, such as sugar in water. Before preparing the solution, you can see the solid and you can see the liquid, that is, you can identify their boundaries. After mixing and stirring them together, you cannot see the sugar anymore, but it is still present, each sugar molecule separated from the others and surrounded by water molecules. If you were to drink this solution, you would know the sugar had not vanished. The substance which is present in minor amount in the solution is the solute and the substance which is present in greater amount is the solvent. The solute is said to dissolve in, or be soluble in the solvent. A solution itself can be either solid (a metal alloy such as brass), liquid (the sugar-water example above), or gaseous (the air we breathe); solutions can be composed of any combination of solid, liquid, and/or gaseous substances, although some combinations are rare. In the chemistry laboratory, we most often encounter solid-in-liquid or liquid-in-liquid solutions. Densities of mixtures, as well as of pure substances, can be measured. Knowledge of the density of a solution not only provides useful information about its composition, but can also aid in determining the concentration of an unknown solution. Today, it is a well known fact and accepted principle that the density of a solution varies proportionately with the mass of the solute present. -24- Archimedes' cry of "eureka" supposedly occurred when he realized that pure gold and adulterated gold (a solid solution) have different densities. Those who would pay for pure gold wanted to know that it was, in fact, pure. Considering a solid-in-liquid solution, such as the sugar water example above, you can see how the density should increase as the ratio of sugar to water increases. (It is often the case, as it is here, that the solid solute has a greater density than the liquid solvent.) Suppose you have a liter of pure water, which has a density of about 1 g/mL (what is the mass of water present?). If you dissolve a certain amount of sugar in the water, and then weigh a liter of the resulting solution, there will be a greater mass of solution present in that one liter volume than there was of the pure water, hence a greater density for the solution. As you add more and more sugar, each time you weigh a liter of the solution it will have an increasingly greater mass and greater density. A convenient aspect of the solution density concept is that the density of a solution and its concentration are linearly related. Concentration is expressed in many different units (e.g. molarity, weight %, volume %), but it refers to the ratio of the amount of solute to the amount of solution (rarely, it refers to the ratio of solute to solvent, an important distinction). Whenever the solute/solvent ratio changes, the concentration of the solution changes and there will be a proportional change in the density of the solution. Look at the graph to see the linear relationship shown as line A. 1.04 1.03 Density (g/mL) A 1.02 # 1.01 1.00 0.99 0 1 2 3 4 5 6 Concentration (M) In this graph, density versus concentration is plotted. If each mark on either axis represents one unit of density or concentration, then by inspection of this example, you can see that each two-unit change in concentration results in a one-unit change in density. How could you exploit this relationship if you do not know the concentration of a solution you have been given? You could prepare a series of solutions using the same solute and solvent, with known, but varying concentrations; these are called standard solutions. After determining the density (by measuring the mass and the volume) of each standard solution, you would plot a graph of density versus concentration. If you then measure the density of the solution of unknown concentration -25- and mark that density on the line of the graph, the corresponding concentration you read from the graph will be that of the unknown solution. The graph illustrates this. Suppose the density you measure is denoted by the # symbol on the density axis. If you then draw a horizontal line (-----) across the graph until you come to the data line A, and then a draw a vertical line down from that point, the concentration of the solution has the value where the vertical line intersects the concentration axis. THE EXPERIMENT In this experiment, you will put together the principles and concepts, which were discussed above. Although determining the unknown concentration of a solution is the ultimate goal, efficiently preparing the series of standard solutions and accurately measuring their densities will consume most of the laboratory period. You will prepare the standard solutions by the method of dilution. This procedure requires taking a known amount of a concentrated stock solution – a solution of known concentration which is used to make diluted solutions - to which is added pure solvent. You will begin by delivering, from a buret into a volumetric flask, a known volume (Vconc) of the stock solution with known concentration (Cconc); then you will add pure solvent until you reach a new known volume (Vdil). Calculating the concentrations of the diluted solutions is simple. No matter what the choice of concentration units, all dilutions can be represented by the equation: (Volume)dil * (Concentration)dil = (Volume)conc * (Concentration)conc or Vdil * Cdil = Vconc * Cconc where the subscript “dil” refers to the dilute solution and “conc” refers to the more concentrated solution. Rearranging the above equation to solve, and substituting known numerical values, gives the concentration of the diluted solution (Cdil). The concentration units will be molarity (M), moles of solute/liter solution. In your measurements and in your calculations, remember to pay attention to uncertainties and significant figures. As you prepare the graph of your results (whether on graph paper or electronically), keep in mind some features of good graphing technique: • • • • label each axis clearly, denoting units; choose unit divisions that can be easily subdivided; divide each axis so that the data fills the graph paper; assign the x- and y-axes so that you avoid drawing your data line with either a steep or a shallow slope; and draw the best straight line of the data points using a transparent straight edge. Problems (show work) -26- 1. Calculate the mass of NaCl needed to prepare 100 mL of a 3.00-M sodium chloride solution. 2. Calculate the volume of solution necessary to make 25 mL of a 2.00-M solution from the 3.00-M solution. 3. Calculate the molarity of 25 mL of solution, which was diluted from 10 mL of a 2.0-M solution. Answers to Selected Problems 1). 17.5 g 2). 16.7 mL 3). 0.80-M PROCEDURE The solid salt used to make your solutions will depend on what the stockroom has available. They will issue it as a 3.00-M standard solution. Your instructor will inform you of the identity of the salt. Since solid salts are convenient to use, the word "salt" will be used throughout the instructions. Preliminary Calculations: As a preliminary step, calculate the volumes of 3.00-M stock salt solution required to prepare 10.00 mL each of: (a) 2-M, (b) 1-M and (c) 0.5-M salt solutions. These are "rough" calculated volumes to guide you when you make your actual dilutions, below. The result of the first calculation is shown in the table below. Vconc = (Vdil  Cdil ) / Cconc Vdil 10 mL Cdil 2M Cconc 3M Vconc  6.7 mL 1. Before beginning your density measurements, obtain a clean, dry 10.00-mL volumetric flask with stopper. Weigh the volumetric flask and stopper on the electronic balance to three decimal places. For all subsequent mass determinations in this experiment, use the same balance. 2. Dilutions a. Rinse your buret with a small quantity (~5 mL) of the 3.00-M stock solution and then pour it out; add about 15 mL of the 3.00-M stock solution to the buret. Record the initial volume of solution in the buret. -27- b. Deliver directly into a 10-mL volumetric flask a volume of 3.00-M salt solution which closely corresponds to (but is not exactly the same as) one of the preliminary calculations above. Record the final volume of solution in the buret. c. Add distilled or deionized water until the volumetric flask is half to three-quarters full, then swirl and mix the contents of the flask. Continue adding water, dropwise if necessary, until the bottom of the meniscus just touches the calibration mark on the volumetric flask. Mix well. Measure the mass of the solution. After you have measured the mass of the solution (below), calculate the actual concentration of the solution in the flask. 3. Densities a. Weigh the volumetric flask containing the salt solution prepared in Step 2 above. From the mass of the solution and its volume, calculate the density. b. Empty and rinse the volumetric flask well with distilled water and shake out excess water. Using the same 10-mL volumetric flask, prepare each of the remaining diluted standard solutions using the same method as in Steps 2b and 2c above. Determine the density of each solution as in Step 3a. c. Determine the density of the 3.00-M stock solution. The buret is not required for this step; simply fill the flask to the mark. d. To compensate for any errors you might have made in preparing the solutions it is helpful to measure the density of water. This value, which will be the smallest density of all (i.e., it has no salt and so is 0.0-M), will be at the left edge of your graph and will help you decide how to draw the best straight line through your data points. 4. In your lab notebook, plot a graph of density versus concentration. Review the principles of good graphs before you begin. Include a linear regression line, or draw the best straight line that fits the data points. This experiment is not performed in pairs and each student must estimate his/her own unknown concentration prior to leaving the lab. Your lab instructor will write your true/accepted unknown concentration value in your lab notebook and you may then enter the data into the computer’s database. 5. Density and concentration of a solution of salt of unknown concentration: a. Obtain a sample of salt solution of unknown concentration from your lab instructor (unknown solutions are prepared by Chem Preps Room). b. Remember to write down the unknown letter. c. Determine the density of the salt solution by transferring the solution into the clean, preweighed 10 mL volumetric flask to the mark. d. Reweigh the volumetric flask and its contents. e. Determine the solution concentration based on the data on your graph. -28- Consider possible sources of error that could account for any difference in your value and the true density value. -29- ______________________________________________________________________________________________________ DENSITY OF SOLUTIONS ______________________________________________________________________________________________________ 1. Mass of dry 10-mL volumetric flask with stopper ___________ (g) 2. Calculations of Actual Concentrations of Diluted Solutions Vdil  Cdil = Vconc  Cconc Volume of stock solution Volume of stock solution delivered in Buret Initial (mL) Final (mL) Concentration of dilute solution Vconc (mL) Cdil (M) 3. Determination of Densities Solution Mass Container Conc. & Solution (Cdil) (g) Mass Container (g) -30- Mass of Volume Density Solution of solution of solution (g) (mL) (g/mL) Stock Solution ~3M ~2M ~1M ~0.5M H2O Unknown Enter your actual concentrations; above values provided as a guide to organize your table. -31- DATA ANALYSIS SUPPLEMENT Data: After you have entered all the masses and volumes, enter your own data into the appropriate Web-based data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures. Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/. Then use Excel to calculate the densities of the various diluted solutions, stock solution and water. Your lab report should include the following in addition to the pages from your lab notebook: • An Excel sheet “Raw Data” showing the raw data for the entire class (Landscape orientation; appropriate column labels and units; sigfigs; student identifying information). • A 2nd Excel sheet “Class Results” showing the calculated results for each student in the class – each in a separate row. The columns should include: o the concentrations of the five solutions (including the distilled water). o the measured densities of the five solutions o the slope and intercept for the density vs. conc. plot. Use the Excel functions =SLOPE(y-range, x-range) and =INTERCEPT(y-range, x-range) o the measured density of the unknown o the calculated concentration of the unknown. • A 3rd Excel sheet “My Results” showing your own results, only. Be sure to include your unknown number at the top of this sheet! There will be six rows, one for each of the known solutions and one for the unknown. There should be two columns: o density (g/mL) o concentration (M) • An Excel chart of density vs. concentration for your own data only, based on the “My Results” sheet described above. Density (dependent variable) should be on the y-axis and concentration on the x-axis. The chart should have five points – one for each of the known solutions. Include a linear regression fit to your points and show the equation of the line on the graph; include the R2 value. Format your charts so that they look like examples on the following pages. Make sure to print graphs and tables in LANDSCAPE orientation. -32- Discussion: In your discussion section address the following questions: • Making the simplistic assumption that the dissolved NaCl(s) does not affect the volume of the solvent water, determine the constants m and b in the equation Molarity = mdensity + b that relates the NaCl molarity to the NaCl(aq) density. Take the density of water to be 1.00 g/mL and the molar mass of NaCl to be 58.5 g/mol. • How does your experimental value for the slope m and intercept b compare with the values derived from the simple model above? • What are the units for slope and intercept for the plot? -33- SAMPLE CALCULATIONS Conc A = C_A (M) =RawData!G39*RawData!B39/RawData!H39 Density A = D_A (g/mL) =RawData!J40/RawData!K40 Slope =SLOPE(G3:K3,B3:F3) Intercept =INTERCEPT(G3:K3,B3:F3) [UNKNOWN] =L3*M3+N3 %ERROR =ABS((O3-P3)/P3)*100 -34- -35- Experiment 5 EMPIRICAL FORMULA OF ZINC IODIDE One of the simplest type of chemical reactions is the combination of two pure elements to form a binary compound. The term "binary" refers to the number of different elements in the compound. A binary compound is composed of just two elements, although there may be more than one atom of an element in the formula unit of compound. For instance, copper metal reacts with sulfur to form the binary compound copper(I) sulfide, according to the chemical equation: 2 Cu + S Cu2S. Even though the reaction may appear to be simple and straightforward, the morphology of the starting materials and reaction conditions determine the actual stoichiometric resultant product. The above reaction generates the desired product if Cu (powder) reacts with molten sulfur or if copper is reacted with sulfur vapour. However, if Cu (pellets) are used as a starting material, then a much higher reaction temperature is required in order to form copper sulfide. In a likewise manner, it is also possible to produce copper(II) sulfide, CuS, and other stoichiometric copper sulfides based on the reaction route of choice. The quantities (masses) of the substances consumed in such a reaction obey two important laws. The Law of Conservation of Mass states that the total mass is conserved in the reaction. In the example above, this means that the total masses of copper and sulfur that reacted must equal the mass of copper sulfide produced. The Law of Definite Proportions states that the ratio of the masses of the elements in a given compound is constant. Thus, the ratio of the masses of copper and sulfur that react will be a simple multiple of the ratio of the atomic weights of those elements. For instance, in the example above, mass of Cu reacted atomic weight of Cu = X• mass of S reacted atomic weight of S The multiple, X, is 2 in this case because copper combines with sulfur in the ratio of two atoms of copper for every one atom of sulfur. Another way of stating this is that copper combines with sulfur in the ratio of two moles of copper for every one mole of sulfur. This information is conveyed in the empirical formula for copper(I) sulfide, Cu2S, where the formula Cu2S is the conventional way of representing Cu2S1. -36- THE EXPERIMENT The purpose of this experiment is to determine the empirical formula of the binary compound, zinc iodide, which is the product of the reaction of the elements zinc and iodine. This is done by preparing a sample of zinc iodide and measuring the reacting masses of the constituent elements. Zn + I2 → ZnxIy The x and y subscripts are to be experimentally determined. In actuality, the absolute values for x and y themselves cannot be determined directly, only their ratio y:x. Then, x and y are chosen as the smallest integers which will give the proper ratio. For example, suppose one finds experimentally that the ratio of y to x is 1.5. The smallest integers which will give this ratio are y = 3 and x = 2. The ratio must be expressed as a ratio of integers since atoms are indivisible and would result in Zn2I3 as the empirical formula. In the reaction of Zn and I2, one of the two reactants will be present in excess. This means that there will be a “left-over”, unreacted quantity of this particular starting material when the reaction is complete. The other reactant (not present in excess) is called the limiting reagent. This reactant will be entirely consumed in the reaction; none will remain when the reaction is complete. The mass of the limiting reagent in the final product is the same as its starting mass, since all of it reacts. The mass of the reactant present in excess must be found indirectly by measuring the mass that is not consumed and then subtracting the value from its mass before the reaction began. Example Consider a reaction in which 1.132 g of copper and 0.620 g of bromine (Br2) were mixed and allowed to react to form the binary compound copper bromide, CuxBry. Suppose that when the reaction was complete, it was found that all of the bromine had reacted, i.e. bromine was the limiting reagent, and that 0.640 g of the copper remained unreacted (excess). The empirical formula for the copper bromide produced can now be determined. First calculate the moles of each element which reacted to form product. moles Br2 = mass Br2 0.620 g = = 0.00388 mol Br2 molecular wt. Br2 2 • 79.9 g/mol Bromine exists naturally as the diatomic molecule Br−Br. We need to conceptually “break” the molecule into atoms so moles of Br atoms = 2 • (moles of Br2 molecules) = 0.00776 mol Br moles Cu = mass Cu consumed (1.132 - 0.640) g = = 0.00774 mol Cu atomic wt. Cu 63.55 g/mol The mole ratio of atoms in the product can now be determined: -37- moles Br atoms 0.00776 = =1 moles Cu atoms 0.00774 Since the mole ratio is 1, the ratio of the integers in the empirical formula y:x is also 1 and the smallest integers giving this ratio are x = 1 and y = 1. Thus the empirical formula must be CuBr (short for Cu1Br1). When using your experimental data in the calculation, do not be alarmed if the final mole ratio obtained is not exactly an integer. It should be close to an integer however, and your number should be rounded to the nearest integer in order to generate a valid empirical formula. THE EXPERIMENT The quantities given below for the reactants are approximate in the sense that they are guidelines. It is not necessary (or even advantageous) to use precisely those quantities in order to obtain a successful outcome for the experiment. However, it is absolutely necessary that the precise quantities that are actually used be recorded in your laboratory notebook for use later in your calculations. For example, when directed to weigh out 2 g of zinc for the reaction, this does not mean that you must use exactly 2.000 g. A mass within a few percent of this is just as good or better, i.e. 1.852 grams. However, you must record your exact mass, for instance 1.852, and use this number in your calculations. One advantage to weighing out a quantity of zinc that is within a few percent of 2 g (not exactly 2.000) is that this can be done much more quickly. Another advantage is that your are less likely to obtain a biased mass. Safety Precautions To avoid potential allergic reactions to iodine vapor, the experiment should be performed in the hood up to the point at which the iodine limiting reagent is consumed. Be careful to avoid inhalation of iodine vapor and avoid contact between the iodine and your skin. Procedure 1. Record the mass of a clean and dry porcelain evaporating dish (PED). directly into your lab notebook; (a) in the data table. 2. Weigh approximately 1 gram of zinc granules directly into the PED. Record the mass of the PED. + the zinc granules; (b) in the data table. 3. Determine the mass of zinc granules by difference; (c) in the data table. 4. Obtain a vial containing iodine crystals and weigh the vial c (LEAVE THE CAP ON WHILE WEIGHING); (d) in the data table. 5. Using a 10-mL graduated cylinder measure out 5 mL of distilled water. 6. Bring the PED, graduated cylinder and vial containing the iodine crystals to the hood. Transfer the iodine crystals directly into the previously weighed PED that contains the zinc granules. (Do not put iodine directly on the balance pan.). -38- Add distilled water to this mixture as well as 5-10 drops of 6.0 M acetic acid and carefully stir and grind (stirring rod) the zinc and iodine mixture. It usually takes 10 minutes for the reaction to go to completion (evidenced by a clear and colorless solution and granules). 7. Re-weigh the vial + cap assembly; record the mass directly into your lab notebook; (e) in the data table. 8. Determine the mass of iodine crystals transferred to the PED by difference; (f) in the data table. 9. Observe and record any evidence you see that a chemical change is occurring. Stir the solution periodically with a stirring rod. It usually takes 10 minutes for the reaction to go to completion (evidenced by a clear and colorless solution and granules). The rest of the experiment may be performed outside the hood. 10. Decant the solution containing the zinc iodide into a beaker; DO NOT DISCARD THE GRANULES from the PED. Wash the excess zinc in the dish three times with 1 mL portions of distilled water acidified with a drop or two of the 6 M acetic acid solution. 11. Dry the excess unreacted zinc by placing the PED on a hotplate (if a drying oven is not available) for at least 20 minutes. Keep the hotplace setting between 2 and 4. You may be asked to use a hot plate for up to 40 minutes to dry the unreacted zinc. 12. Once the zinc is completely dry, allow the dish to cool completely to room temperature (1015 minutes). 13. Weigh the P.E.D. containing the dried unreacted zinc; (g) in the data table. 14. Determine, by difference, the mass of unreacted zinc; (h) in the data table. 15. Determine the mass of zinc that reacted with the iodine; (i) in the data table. Calculate the mole ratio of I:Zn and then determine the empirical formula. Enter your empirical formula into your data table. Pay attention to significant figures, using the appropriate number of them to reflect the precision of your mass measurements. Do not round-off until the final calculation of the empirical formula. Disposal and Cleanup --- Dispose of the excess solid zinc in a waste jar (not the sink). -39- _______________________________________________________________________________________________ EMPIRICAL FORMULA OF ZINC IODIDE _______________________________________________________________________________________________ (a) Mass of porcelain evaporating dish (P.E.D.) _______________ (g) (b) Mass of P.E.D. + Zn granules at beginning _______________ (g) (c) Mass of Zn at beginning _______________ (g) (c) = (b) – (a) --------------------------------------------------------------------------------------(d) Mass of screwcap vial + I2 _______________ (g) (e) Mass of screwcap vial _______________ (g) (f) Mass of I2 at beginning _______________ (g) (f) = (d) – (e) --------------------------------------------------------------------------------------(g) Mass of P.E.D. + unreacted Zn _______________ (g) (h) Mass of unreacted Zn _______________ (g) (h) = (g) – (a) -40- (i) Mass of reacted Zn _______________ (g) (i) = (c) – (h) (j) Moles of reacted I atoms (in product) _______________ (k) Moles of reacted Zn atoms (in product) _______________ (l) Mole ratio (I to Zn) _______________ (m) Empirical formula _______________ Show calculations for h, i, j, k and l on the corresponding pages of your data table. DATA ANALYSIS SUPPLEMENT Data: After you have recorded all the masses, enter your own data into the appropriate Web-based data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures. Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/. Then use Excel to calculate the moles of Iodine and Zinc in the product, the Mole ratio and the Empirical Formula. Your lab report should include the following in addition to the pages from your lab notebook: • An Excel sheet “Raw Data” showing the raw data for the entire class (refer to guidelines stated on page 42). Your personal data should be in boldface. • A 2nd Excel sheet “Class Results” showing the calculated results for each student in the class – each in a separate row. Your personal data should be in boldface. The columns should include: o Mass of unreacted zinc o Mass of reacted zinc o Moles of reacted I atoms o Moles of reacted Zn atoms o Your Mole ratio (I to Zn) -41- o • Your Empirical formula An Excel chart showing a histogram of the distribution of mole ratio results. Use equally spaced “bins” chosen to display the “spread” in the class results. Leave space between the vertical bars in the histogram. Discussion: In your discussion section address the following questions: • What was your own percent error for the mole ratio? • What percentage of the students in your section obtained a mole ratio that was within (a) 10% and (b) 25% of the true value? • What percentage of the students in your section obtained a mole ratio that was higher than the true value? • What experimental error(s) might have caused the calculated mole ratio to come out too high? • What experimental error(s) might have caused the calculated mole ratio to come out too low? -42- Experiment 6 SYNTHESIS OF ASPIRIN Aspirin is one of the best known and most widely used medicines in modern society. The Bayer Chemical Company devised the pharmaceutical formulation, acetylsalicylic acid, more than 100 years ago. Its medicinally active part, salicylic acid, is a naturally-occurring substance that can be extracted from willow tree bark and meadowsweet flowers. It had long been known and used as an herbal fever reducer, but salicylic acid is too irritating to the mucous membranes of the mouth, esophagus and stomach to be administered comfortably as such. Chemically modifying the salicylic acid by attaching an acetyl group minimizes the adverse side effects while retaining the desirable analgesic, antipyretic, and anti-inflammatory properties. The same strategy of chemical modification was used later to modify morphine, a powerful but addictive painkiller. However, when morphine was converted into diacetylmorphine, the result was heroin. The modification intensified the addictive property instead of reducing it. As you inspect the molecular structures of salicylic acid and aspirin, below, notice the small change in structure as a hydrogen atom in salicylic acid (in boldface) is replaced by the acetyl group, CH3 C O Note: The vertices of the “hexagons” in salicylic acid and aspirin are occupied by carbon atoms. H H O C O H H O CH3 C O C H H O C O H CH3 C O H CH3 O C CH3 H O H H H Salicylic acid C7H6O3 O Acetic anhydride O O Acetyl salicylic acid Acetic acid C9H8O4 C2H4O2 C4H6O3 A substance related to aspirin is oil of wintergreen, or methyl salicylate. Both of these compounds are synthesized from salicylic acid. Notice the similarity of names, which implies a similarity of molecular structure. Oil of wintergreen is largely used as a flavoring agent, but ingestion of relatively small amounts may cause severe poisoning and death. Oil of wintergreen is used in many liniments, however, as it is absorbed through the skin. Once absorbed, it may be cleaved back to salicylic acid; thus it is a source of pain relief, albeit localized pain relief. -43- H H O C O H H H O H C O CH3 H O H H C O H O H H H O H H H H Salicylic acid Methanol C7H6O3 CH4O Methyl salicylate C8H8O3 Water H2O THE EXPERIMENT In this experiment you will synthesize aspirin in much the same way as it is manufactured industrially. You will also calculate the percent yield of the aspirin you synthesized. In the synthesis of aspirin from salicylic acid and acetic anhydride, the most likely impurity in the final aspirin product is salicylic acid itself, which is present either from incomplete reaction or from hydrolysis (the reverse reaction) of the aspirin product during the isolation steps. The salicylic acid impurity is removed during the various stages of the purification and in the final crystallization of the product. The presence of salicylic acid is detected by its colorful reaction with ferric chloride. _________________________________________________________________________________________________________ Percent Yield The theoretical yield for any reaction is the maximum number of grams (or moles) of product that may be obtained from the reaction based on the reaction stoichiometry. Before a yield can be calculated, it is necessary to first balance the chemical equation for the reaction. Next, the moles of each reactant are calculated from the grams of each reactant (or volume if the reactant is liquid) which were used in the reaction. If the ratio of the moles of reactants in the balanced equation is different from the ratio of the moles of reactants actually used for the reaction, it is necessary to determine the limiting reagent in the actual synthesis. The limiting reagent is the reactant which theoretically would be completely consumed in the reaction based on the molar ratio calculated from the actual quantities of reactants. Any other reactants are therefore present in excess, and will not be consumed completely. There can be no more moles of product formed than that amount available from the moles of the limiting reagent. The limiting reagent determines the maximum or theoretical yield of product(s). If all of the reactants in the synthesis are actually used in the synthesis in the same molar ratio as in the balanced equation, any of the reactants may be used to calculate a theoretical yield. -44- 𝑡ℎ𝑒𝑜. 𝑦𝑖𝑒𝑙𝑑 (𝑔) = (𝑚𝑜𝑙 𝑙𝑖𝑚𝑖𝑡𝑖𝑛𝑔 𝑟𝑒𝑎𝑔𝑒𝑛𝑡 𝑐𝑜𝑛𝑠𝑢𝑚𝑒𝑑)𝑥 ( 𝑚𝑜𝑙 𝑝𝑟𝑜𝑑𝑢𝑐𝑡 ) 𝑥 (𝑚𝑜𝑙. 𝑤𝑡. 𝑝𝑟𝑜𝑑𝑢𝑐𝑡) 𝑚𝑜𝑙 𝑙𝑖𝑚𝑖𝑡𝑖𝑛𝑔 𝑟𝑒𝑎𝑔𝑒𝑛𝑡 The actual yield is the number of grams (or moles) of product actually obtained in the synthesis reaction. The percent yield describes the efficiency of the reaction: %𝑦𝑖𝑒𝑙𝑑 = ( 𝑎𝑐𝑡𝑢𝑎𝑙 𝑦𝑖𝑒𝑙𝑑 ) 𝑥 (100) 𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑦𝑖𝑒𝑙𝑑 Sample Calculation A sample calculation using an example of the aspirin preparation is shown below. The molecular weights of the compounds can be found in Table 1. From the balanced chemical equation you can determine the molar ratios among the reactants and products. Here, one mole of salicylic acid reacts with one mole of acetic anhydride to give one mole of aspirin product and one mole of acetic acid. The molar ratio of the reactants is 1:1 as is the ratio of salicylic acid (reactant) to aspirin (product). Next you must determine which of the two reactants (if either) is the limiting reagent in a laboratory synthesis. Suppose you started with 4.21 grams of solid salicylic acid and 15.0 mL of liquid acetic anhydride (density = 1.08 g/mL). Moles of salicylic acid used = 4.21 g • Moles acetic anhydride used = 15.0 mL • 1 mol = 0.0305 mol 138.12 g 1.08 g 1 mol • = 0.159 mol 1 mL 102.09 g Therefore, since salicylic acid and acetic anhydride react in a 1:1 ratio, but the ratio actually used is 0.0305/0.159 or 1 : 5.2, the acetic anhydride used here is clearly in excess and thus salicylic acid is the limiting reagent. The theoretical yield must be calculated using moles of salicylic acid. Theoretical yield (grams) = 0.0305 mol salicylic acid • 1 mol aspirin 180.17 g aspirin • = 5.50 g aspirin 1 mol sal. acid 1 mol aspirin Suppose the actual yield = 3.92 g aspirin, then: Percent yield = (3.92 g ÷ 5.50 g)  100 = 71.3% -45- Vacuum Filtration A common laboratory procedure is the separation of a solid from a liquid by filtration. Suction filtration requires a water vacuum aspirator and special equipment, but once set up, it proceeds rapidly. How to filter with a Buchner Funnel: 1. Assemble the equipment (ring stand, clamp, filter flask, rubber tubing, Buchner funnel, and filter paper (see Figure). Clamp the filter flask to the ring stand. Attach the flask side arm to the vacuum aspirator with the rubber tubing. 2. Place the filter paper in the funnel, and moisten it with water or the solvent to be used. Do not allow it to dry before proceeding to the next step. 3. Turn the vacuum aspirator on. 4. Pour into the funnel, at a moderate rate, the solution to be filtered. 5. To interrupt the vacuum, remove the rubber tubing from the flask. Do not turn off the water first. 6. Rinse the solid in the funnel by pouring into the funnel a small amount of cold solvent. Stir the contents with a stirring rod, being careful not to tear the filter paper. Allow the solid to dry by drawing air through the funnel. 7. Every time you do another filtration, use a fresh piece of filter paper. Figure 1 ALWAYS SECURE SETUP USING A RING STAND AND UTILITY CLAMP. -46- PROCEDURE Hazard: Acetic anhydride is a corrosive liquid and is quite flammable. Avoid breathing vapors and carry out all transfers in a HOOD. Concentrated sulfuric acid is also highly corrosive. Handle it with great care. Before you begin the synthesis, you need to make a water hot bath. Half-fill a 400 mL beaker with tap water and heat it on an electric hot plate. Bring the temperature of the water to 45-50°C (hot plate setting of about 3-4). This will serve as the hot bath. It is important that the temperature of the water bath is constant before heating the reactants. 1. Precisely weigh, using a weigh boat, approximately 2 g of salicylic acid. Carefully transfer the salicylic acid into a clean and dry 125-mL Erlenmeyer flask. 2. Take your flask containing the salicylic acid crystals to the HOOD and carefully add 5 mL of acetic anhydride. Measure the volume in a small (0-mL graduated cylinder. 3. In the hood, carefully add 4 drops of concentrated sulfuric acid (H2SO4) using a dropper. (Sulfuric acid acts as a catalyst in the reaction and is not the limiting reagent.) Gently swirl the mixture. 4. Heat the flask in the water bath (45-50°C) for 10 minutes, stirring until all the salicylic acid dissolves. (Do not stir with a rubber spatula.) 5. Remove the flask from the bath and allow it to cool to room temperature. Solid particles of acetylsalicylic acid (aspirin) should begin to precipitate out of the reaction mixture as it cools if it has not done so already. If precipitation does not occur, scratch the inside walls below the liquid level of the flask with a glass rod. This will provide nucleation sites for crystals to form. 6. Cool the mixture in an ice bath until no more aspirin precipitates from the liquid solution. While you are waiting, assemble the vacuum filtration apparatus. 7. After precipitation is complete, add 20 mL of ice cold water to the reaction flask, stir to mix, and again cool the mixture in an ice bath. 8. Collect the solid product in a Buchner funnel by vacuum filtration while the solution is still cold. The filtrate (the liquid solution that collects in the filter flask) can be used to rinse more product from the Erlenmeyer flask in which you performed the synthesis until all solid has been collected. 9. Rinse the solid twice with small portions of cold water. Continue drawing air through the crystals in the Buchner funnel by suction until the crystals are relatively dry. -47- 10. Transfer dried crystals to a pre-weighed watch glass for drying and allow then to dry for at least twenty minutes. 11. Record the mass. _____________________________________________________________________________ PART A - Ferric chloride test In order to determine if unreacted salicylic acid is present with your aspirin product, perform a ferric chloride color test. 1. Into a small test tube containing about 5 mL of water, add a small amount of aspirin (about the size of a pen tip). Stir to dissolve. 2. Into a second test tube put 5 mL of water and about the same amount of salicylic acid as the amount of aspirin you put in the first test tube. Stir to dissolve the solid. 3. Into a third tube, put 5 mL of the filtrate you saved from the synthesis procedure. 4. To each test tube, add 10 drops of 1% FeCl3 and stir. 5. Compare the color in the three test tubes. Explain your results. _________________________________________________________________________________________________________ Weigh the aspirin product and then calculate the percent yield. (If your aspirin is still wet, your percent yield may be more than 100%.) Submit your product to your instructor in a vial, labeled with your name, the product name, the yield in grams, and percent yield. -48- TABLE 1 The molecular weights, formulas, and densities for the substances which have been mentioned in this experiment are given below. [CRC Handbook (92nd Ed.)] Compound Formula M.W. (g/mole) Density (g/mL) salicylic acid C7H6O3 138.12 acetic anhydride C4H6O3 102.09 acetylsalicylic acid C9H8O4 180.17 acetic acid C2H4O2 60.05 1.0492 methanol CH4O 32.04 0.7914 methyl salicylate C8H8O3 1.0820 152.16 Problems (show work) 1. What is the purpose of (a) recrystallization; (b) the sulfuric acid in aspirin synthesis? 2. If you use 12.22 g of salicylic acid and an excess of acetic anhydride in the synthesis of aspirin, what would be the theoretical yield of acetylsalicylic acid in moles? In grams? Determine the percent yield if 10.50 g of aspirin is isolated? Answers to Selected Problems 2) 8.847 x 10-2 mol; 15.94 g; 65.87% -49- _____________________________________________________________________________________________________ SYNTHESIS OF ASPIRIN _____________________________________________________________________________________________________ Chemical Equation for synthesis of Aspirin: Show all calculations. Pay attention to significant figures. Reagent volume mass mole salicylic acid ----------- ____________ ___________ acetic anhydride ___________ ___________ ___________ limiting reagent ______________________________________ Theoretical Yield of Aspirin ____________________________ Mass of aspirin isolated _______________ Percent yield of aspirin _____________ Part B Observations and Conclusion: (These go directly in your lab notebook.) DATA ANALYSIS SUPPLEMENT Data: After you have recorded all the masses and volumes, enter your own data into the appropriate Web-based data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures. Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/. Then use Excel to calculate the moles of Salicylic acid and Acetic anhydride, and the Theoretical yield and Percent yield of aspirin. Your lab report should include printouts of data and result tables from your spreadsheet in addition to the pages from your lab notebook. Each table should be formatted as discussed in earlier experiments. • An Excel sheet “Raw Data” showing the raw data for the entire class (properly formatted column headings, units, sigfigs, etc.). Row and column labels should always be printed for each Excel sheet. -50- • • A 2nd Excel sheet “Class Results” showing the calculated results for each student in the class – each in a separate row. Each table should be formatted as in previous classes with lines separating the table headings from the contents. A line below the contents should also be drawn. Make sure to show an example calculation for each column. This can be done by copying the equation to a blank cell below the elements in the table. An apostrophe in front of it makes it text instead of an equation. The columns should include: o Moles of salicylic acid o Moles of acetic anhydride o Theoretical yield of aspirin o Mass of aspirin isolated o Percent yield of aspirin An Excel chart showing a histogram of the distribution of percent yield results. Use equally spaced “bins” chosen to display the “spread” in the class results. Discussion: In your discussion section address the following questions: • Which reactant, Acetic anhydride or Salicylic acid is the limiting reagent? Why is it called so? • What is the most likely impurity in the final aspirin product? How is its presence detected? • What are the possible sources of error in this experiment? Refer in your answer to the class histogram as well as to your own results. Printing Tables: • Large tables should be printed in the landscape mode and sometimes it might be necessary to go to “page setup” and select “Fit to 1 page” to make the page fit on one page. • Include column and row headings on the printout so that the instructor can more easily understand equations used. This can be done by going to File → Page Setup → Sheet tab → Select “Row and Column Headings”. -51- Experiment 7 VOLUMETRIC ANALYSIS: TITRATION OF VINEGAR Volumetric analysis is the quantitative determination of substances by titration. Titration is a laboratory procedure in which one solution is added, by means of a buret, to a second solution until all of the reactant in the second solution has been consumed in a chemical reaction. The solution added from the buret is the titrant. A standard solution is a solution whose concentration is known accurately. The standard solution may be used to titrate a solution of unknown concentration, and by means of appropriate calculations, the concentration of the second solution can be determined. Acid-Base Titrations When a solution of an acid is mixed with a solution of a base the reaction that occurs is called a neutralization reaction. When a strong acid (symbolized HA) and a strong hydroxide base are mixed the reaction can be represented by the partial ionic equation: H+A−(aq) + OH−(aq) A−(aq) + H2O (l) (1) When a diprotic acid (H2A) is neutralized by a strong hydroxide base, the acidic protons are removed in a stepwise fashion : H2A(aq) + OH−(aq) HA−(aq) + H2O (l) (2a) HA− (aq) + OH−(aq) A2−(aq) + H2O (l) (2b) H2A(aq) + 2OH−(aq) 2H2O (l) + A2−(aq) In aqueous media, the overall reaction for the diprotic reaction can be represented in the following manner: 2H+ (aq) + A2− (aq) + 2OH−(aq) 2H2O (l) + A2−(aq) (2c) In both cases, the reactions are rapid and essentially quantitative. Such acid-base neutralization reactions are commonly used in volumetric analysis to determine the concentrations of acidic or basic solutions. The equivalence point (or neutralization point) of the titration is often recognized by the color change of an indicator which is added to the titrated solution. A common laboratory problem is to determine the concentration of acid in a solution. To do this, the acidic solution is titrated with a standard base, usually NaOH of known concentration. However, when solutions of NaOH are prepared, the concentration of the base solution is not accurately known because of the difficulty of handling solid NaOH. In order to establish the -52- concentration of the basic NaOH which is to be used in a titration, the NaOH solution concentration must first be determined by titration with an acid. The acid which is used to titrate the NaOH solution is called a primary standard acid solution. The primary standard is easily handled and solutions of very accurately known concentration can be prepared. The overall procedure can by represented by: primary standard acid [known conc.] secondary standard titrate > base titrate [calculate conc.] -53- > acid [calculate conc.] THE EXPERIMENT Vinegar is an aqueous solution of acetic acid, a simple organic acid. The expanded structural formula for acetic acid is H H C O C O H H Only the hydrogen bonded to the oxygen is acidic. The other three hydrogens bonded to carbon do not react with a base such as NaOH. Acetic acid is a weak acid, which reacts with NaOH according to the equation: CH3CO2H + Na+(OH)− CH3CO2- Na+ + HOH (3) In this experiment, the ultimate goal is to determine the concentration of acetic acid in vinegar purchased from the grocery store by titrating it with a standard NaOH solution (a secondary standard). For the standardization of the NaOH solution, oxalic acid (in the form of its hydrate H2C2O4•2H2O), will be used as a primary standard. Both hydrogens, which are bonded to oxygens in the oxalic acid molecule, are acidic and thus 2 moles of NaOH will be required to neutralize every 1 mole of H2C2O4 during the titration: H2C2O4 + 2 Na+(OH)− 2 Na+ (C2O4)2− + 2 H2O (4) Oxalic acid is also a simple organic acid and is used frequently as a bleaching agent, especially in marine settings. It is very corrosive, so be careful when you handle it. Phenophthalein, which is a suitable indicator for the titration of a weak acid with a strong base, is used in the titration. Phenolphthalein is another weak organic acid. Its color change is from colorless in acidic solution to pink in basic solution upon removal of the acidic proton present in the molecule. A more exhaustive treatment of indicators and their roles during titration will be addressed during the second semester laboratory course (LeChatelier’s Principle and Equilibrium). The molarity of the acetic acid solution can be determined from the volume and molarity of the NaOH used in the vinegar titration. Finally, the weight% of acetic acid in the vinegar solution will be calculated and compared with the 5wt% allowed by Federal Law. Calculations The fundamental equation used in titration calculations is the same as the one previously used for dilutions: VA MA = VB MB -54- The volumes and concentrations pertain to an acid-base neutralization reaction rather than dilution of a concentrated solution. The only adjustment needed for the acid-base titration calculation is for those cases where the reaction stoichiometry is not 1:1. In a 1:1 neutralization such as in equations (1) and (3), the number of moles of acid which transferred H(+) and the number of moles of base (OH-) which accepted them must be the same at the neutralization point: VAMA = VBMB VAMA = molesA = nA ; (5) VBMB = molesB = nB nA = nB At the neutralization point in reactions, such as those in equation (2) and (4), the two moles of H(+) from one mole of the oxalic acid have combined with 2 moles of OH(-) from 2 moles of the base. 2(VH2AMH2A) = VBMB (6) 2 • molesA = molesB 2nA = nB Example 1: Calculate the molarity of KOH if 20.10 mL was used to neutralize 13.62 mL of 0.75-M HCl. 1 KOH (aq) + 1 HCl (aq) 1 KCl (aq) + 1 H2O (l) VHClMHCl = VKOHMKOH MKOH = VHClMHCl (13.62)(0.75) = = 0.51 - M VKOH (20.10 Note: Either mL or L may be used for the volumes (V). -55- Example 2: How many milliliters of 2.1-M NaOH are required to neutralize 42.85 mL of 3.6-M H2SO4? 2 NaOH (aq) + H2SO4 (aq) Na2SO4 (aq) + 2 H2O (l) 2 (VH2SO4MH2SO4) = VNaOHMNaOH VNaOH = 2(VH2SO4MH2SO4) 2(0.04285)(3.6) = = 0.1469 L = 146.9 mL MNaOH 2.1 PROCEDURE CAUTION: Oxalic Acid is corrosive. 1. Preparation of Oxalic Acid Standard Solution The oxalic acid solution has been prepared already for you. Make sure to record the molarity written on the bottle! 2. Standardization of the NaOH Solution a. Rinse one buret twice with 5 mL portion of approximately 0.5-M NaOH solution and discard the rinsing solution. Run at least a portion of each rinse through the stopcock and the tip so that all parts inside the buret will have been exposed to the NaOH solution. b. Using a 50-mL graduated cylinder, measure approximately 35 mL of the sodium hydroxide solution. Add this to the buret and let a little flow thro...
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