CHEM 112 Lab Manual Spring 2024 Content Page(s) Lab Policies Laboratory Safety Rules Pregnancy Notice Equipment Check-in List Review of Proper Equipment Use Academic Honesty & Lab Grades Lab Report Guidelines Making Graphs in Excel Week 1: 1/16 – 1/18 NO LABS FIRST WEEK OF CLASS Week 2: 1/22 – 1/25 Orientation & Check-in Week 3: 1/29 – 2/1 Experiment 1: Organic Structures & Nomenclature Week 4: 2/5 – 2/8 Experiment 2: Gas Chromatography, Intermolecular Forces & Boiling Point Week 5: 2/12 – 2/15 Experiment 3: Net Ionic Equations Week 6: 2/19 – 2/22 Experiment 4: Rate Law for Bleaching Crystal Violet Week 7: 2/26 – 2/29 Experiment 5: Le Chatelier’s Principle Week 8: 3/4 – 3/7 Experiment 6: Luminol Lab Practical Week 9: 3/11 – 3/14 NO LABS – Spring Break Week 10: 3/18 – 3/21 Experiment 7: Spectrophotometric Determination of Cu in a Penny Week 11: 3/25 – 3/28 Experiment 8: Synthesis of an Oxalate Hydrate Week 12: 4/1 – 4/4 Experiment 9: Percent Oxalate Determination by Titration Week 13: 4/8 – 4/11 Experiment 10: Bleach Titration Lab Practical Week 14: 4/15 – 4/18 Experiment 11: Redox Reactions & Equations Week 15: 4/22 – 4/25 LAB MAKE-UP WEEK Week 16: 4/29 – 5/2 Check-out & Lab Clean-up (mandatory) 2 3-5 6 7 8-9 10 11-12 13-15 NA 1-15 16-22 23-27 28-35 36-41 42-50 51-52 NA 53-60 61-62 63 64-66 67-71 NA 72 If you have an excused absence from a lab, it is your responsibility to initiate contact with your instructor to get an OK for making up the lab with another section OR Lab Make-Up Week, then you must contact the instructor of the other section to get an OK from them. Contact concerning lab absences should be made a week in advance for predictable absences and within 24 hours of your lab period for unpredictable absences. Spring 2024 CHEM 112L Lab Sections Section 001 002 003 004 005 Day Monday Monday Wednesday Wednesday Thursday Time 1-2:50 3-4:50 1-2:50 3-4:50 2-3:50 pm pm pm pm pm Room 444 444 444 444 444 Instructor Z. Lee Z. Lee A. Macintosh D. Fulmer D. Fulmer 1 Instructor Email z.lee@moreheadstate.edu z.lee@moreheadstate.edu a.macintosh@moreheadstate.edu d.fulmer@moreheadstate.edu d.fulmer@moreheadstate.edu Lab Policies 1. You must attend each scheduled laboratory session. If you have a documented, excused absence, bring the documentation to your instructor before missing lab if possible and as soon as possible otherwise. Don’t wait until the next lab. Missing 3 lab sessions or not turning in lab reports equivalent to 3 lab sessions will result in an automatic E in Chemistry 112. The smooth, safe operation of lab depends on your hearing prelab instructions and safety cautions. If you are late, your instructor may assess a grade penalty. 2. Your lab grade will be based on the quality of your experimental work based on your written lab report, on the quality of your report in terms of communication and/or a quiz. No lab reports will be accepted for experiments which you did not attend and do conscientious work. Sometimes you may be given a quiz instead of submitting a report (this will be announced ahead of time). 3. All lab reports are due at the beginning of the lab period immediately following the day an experiment is completed; they must be done using a word processor unless otherwise stated . All lab reports are expected to be original and done by individuals, not pairs; this includes any tables, graphs, etc. If your instructor tells you that you may write a report with your partner, then both partners are responsible for all of the report. Having one partner unavailable or data unavailable is not an excuse for no report. 4. When working in the lab, you must follow all the safety rules. When you leave the lab, your work areas must be cleaned and all equipment returned to its original place. The last 4 persons leaving the lab are responsible for its overall cleanliness and orderliness. Failure to follow safety rules or maintain cleanliness and order will result in a 50% penalty for that lab. If you violate safety rules, your instructor may eject you from lab and give you a zero for the day. 5. Lab notebooks. You are to keep a bound lab notebook with data in chronological order. Record data in ink and strike through any mistakes with a single line–to not obliterate the data. 2 Name:_______________________________ Course:___________ (print) Date:___________ LABORATORY SAFETY RULES 1) Responsible behavior is essential. Think about the consequences of your actions or inactions. Make note of the safety precautions written in the lab text and told to you by the laboratory instructor. If you are unsure about a procedure, ASK! 2) WEAR APPROVED EYE PROTECTION AND AN APRON AT ALL TIMES IN THE LAB. Goggles and safety glasses are available in the bookstore. If you wear glasses, be sure your eye protection fits over your glasses. If you must wear contacts, then you must wear goggles. Your shoes and clothing should also provide protection, i.e. wear closed-toe shoes whose "uppers" repel spills. Wear clothing that covers you from the ankles to your throat, with the equivalent of a crew-neck t-shirt on top. Failure to do so will result in expulsion from lab. Confine long hair and loose clothing. Leggings and yoga pants do not count as pants. They fit too closely to repel spills. So no bare midriffs sleeveless shirts, shorts, short skirts, high-heeled or slick-soled shoes. 3) In case of fire, accident, or chemical spill, notify the instructor at once. a) Wet towels can be used to smother small fires. b) In case of a chemical spill on you body or clothing, wash the affected area with large quantities of running water. Be sure to first remove the clothing that has been wetted by chemicals to prevent further reaction with the skin. c) If you should get a chemical in your eye, wash with flowing water from the eye wash for at least 15 minutes. Get medical attention immediately. d) Report all injuries immediately to your instructor. 4) Note the locations of the safety equipment listed below. Make sure you know how to use them. a) Eye Washes b) Safety Shower c) Fire Extinguisher d) Blanket e) Telephone f) Sinks g) Hoods h) Exits 5) EMERGENCY HELP IS AVAILABLE BY DIALING 911. Telephones are available in the C hallway between 445 & 446 and in the middle (B) hall outside 423. 6) Do not eat, drink, or chew on anything in the laboratory. This includes food, gum, tobacco, pens, and other chemicals. Also, you should not handle many lab chemicals. They may absorb through the skin or be ingested if you lick your fingers. Wash thoroughly before you leave the lab, even if you haven't "touched" anything. 7) Avoid breathing fumes of any kind. a) To test the smell of a vapor, "waft" the vapor toward you. Test only those vapors specified. b) Work in a hood if there is a possibility that noxious or poisonous vapors may be produced. If unexpected vapors are produced during an experiment, move the experiment to a hood immediately. 8) Keep your work area clean at all times. Clean up spills and broken glass immediately, whether at your bench or in the common area. Clean your work space, including wet-wiping the bench top and putting away all chemicals and equipment, at the end of the period. 9) Perform no unauthorized experiments. This includes using more reagent than instructed. 3 10) Never work alone in the laboratory. 11) Always pour acids into water when mixing. Otherwise, the acid can violently spatter. "Acid into water is the way that you oughter." 12) Do not force a rubber stopper onto glass tubing or thermometers. Use a split-hole stopper if possible. If not, lubricate the tubing and the stopper with glycerol or water; use paper or cloth toweling to protect your hand; and grasp the glass close to the stopper. 13) Beware of hot iron rings and hot glass. They "look" cool long before they can be handled. 14) Dispose of small quantities of excess non-toxic liquid reagents by flushing them down the sink with copious quantities of water. Dispose of solids and toxic liquids in the waste containers provided. Never use your eyedroppers to remove liquids from a reagent bottle. Always pour reagents into your own clean containers. Never return reagents to the bottle. 15) Spatters are common in general chemistry laboratories. Test tubes containing reacting mixtures or those being heated should never be pointed at anyone. 16) Carefully read the experiment before coming to lab. Review techniques to be used and safety precautions to be taken. 17) Finally, and most importantly, THINK about what you're doing. If you doubt that what you're doing is safe, don't do it. Check with your instructor. Ask questions. Minimization of exposure of all students to the chemicals we work with in lab is an important goal for both the instructor and the students. The experiments we pick, chemicals we choose to use, and safety procedures we emphasize are all a part of attaining that goal. However, be aware that the effects of many chemicals are different for different people. Certainly, we all know about allergies, and have heard about how different medicines can interact when taken together and cause disastrous results. Analogously, certain medical conditions such as asthma, allergies, etc. can be worsened by exposure to certain chemicals. Please inform your instructor about any allergies or medical conditions of which they should be aware. Some chemicals which are not generally considered toxins are reproductive toxins. They may affect the fetus directly or affect the ability of the parent (fathers, too) to produce healthy offspring. Also realize that not all chemicals have been tested to see if they are reproductive toxins. Anyone who is sexually active needs to be diligent in their efforts to minimize their chemical exposure in the lab. If you are or become pregnant, you need to notify your instructor as soon as possible. Notification forms are available in your manual or from your instructor. ........................................................................................................................................................... I have read and understand the Laboratory Safety Rules and have retained a copy for my reference. ________________________________ (signature) ________________ (course & section) __________ (date) Please list any allergies or medical conditions (including pregnancy) of which the instructor should be aware. Please indicate if you wear contacts. ___________________________________________________________________________ Revised: RKC 8/17 4 Lab Attire Rules NOT APPROVED: • • • • • • • No leggings, no jeggings, no skinny jeans (or similar skin tight pant by a different name), no shorts, no capris, no ripped/holey pants, no yoga pants. Legs will be totally covered by a loose fitting, durable material. Just a lab apron or coat is insufficient. No tank tops, no sleeveless tops, no cropped shirts, no plunging necklines, no spaghetti straps, or ripped shirts. No Sandals, no open toe, open back, or open weave shoes, shoes with holes in the top or sides, no Birkenstocks, Mary Janes, or Crocs No hats or caps No head phones or earbuds No “forehead protectors” (your eye protection comfortably positioned on your forehead). The instructor reserves the right to determine inappropriate clothing at any time that is not listed here. If you are curious about a particular piece of clothing, consult your instructor a minimum of 24 hours before lab. APPROVED MUST HAVES: • • • • • You must wear a top that covers you upper torso and shoulders equivalent to short sleeves (a good example is a crew neck t-shirt, that fits properly). You must wear pants that covers your lower torso from your waist all the way to the top of your shoes (a good example is a regular pair of jeans or kaki’s that fit properly) You must wear shoes that cover you whole foot and fit correctly, (your no mesh tennis shoes) You must wear a plastic/vinyl apron or lab coat You must wear colorless safety goggles or safety glasses. Safety glasses must be approved by your instructor. They must have side and top shields. VisorGogs are a nice option for some as they breath better than traditional goggles.. The goggles will indicate some sort of code relating to an ANSI Z87.1 Standard (ANSI Z87.1-1989, ANSI Z87.1-2003, ANSI/ISEA Z87.1 D3-2010, ANSI/ISEA Z87.1-2015 or Z87+ D3). If you do not abide by these rules at any time during your lab period you will be excused from lab with zero credit earned. Time is not available to run back to residence for appropriate attire. When in doubt, simply carry scrubs or pants to wear over shorts while in lab. 5 Notification of Pregnancy I hereby notify my lab instructor that I am pregnant. I give permission to the laboratory supervisor to send to my doctor a list of the chemicals I will be using this semester. I also authorize my doctor to discuss with my instructor or the laboratory supervisor any hazards associated with exposure to these chemicals. I also authorize my doctor to respond in writing to the lab supervisor with recommendations for extra precautions they recommend, including exclusion of specific chemicals from the experiments I perform. Student's Name ________________________________________ (please print) Student's Signature _____________________________________ Student's Social Security # _________________ Student's Date of Birth ____________________ Doctor's Name _________________________________________ Doctor's Address _______________________________________ _______________________________________ _______________________________________ Doctor's Phone # ________________________ Course _______________ Lab Instructor _________________________________ Date _________________ 6 Equipment List for Chemistry 112 Name(s) Drawer # __________ Date _______________ Lab section or lab day/time ____________________ In week one, check your desk to see that you have the following items, clean and unbroken. Your instructor will tell you where to place extra items. Replace any broken or missing equipment. Then keep this page for the last day of lab. 1- 20mL beaker 2- 100 mL beaker 2- 250 mL beakers 2- 250 mL Erlenmeyer flask 2- scoopula 1-test tube clamp 10 test tubes, 13 mm x 100 mm 2- Pasteur pipet (looks like a long medicine dropper), with bulb 2- 50 mL beakers 2- 150 mL beaker 1- 400 mL beaker 2- funnel 1- triangle 1 test tube rack 1- pair crucible tongs 2- glass stirring rods 1- 500 mL plastic bottle for Crystal violet (stained is OK) 1- 100 mL graduated cylinder 1- 10 mL graduated cylinder 2- 250 +/or 125 mL plastic bottles Check the large locker by your drawer for the following: one ring stand one ring one utility clamp one wire gauze one bunsen burner and tubing Keep this page for the last day of lab! ****************************************************** Check out: Student’s signature(s): Date Instructor’s signature 7 Review of Proper Equipment Use Many labs direct you to use YouTube videos and to review the proper use of various pieces of equipment before lab. Below are links to videos associated with some important glassware and questions you should be able to answer about each. Volumetric Pipet Use of a pipet: https://www.youtube.com/watch?v=TffTiRw8cQY 1. Be able to describe the procedure used to rinse the pipet properly. 2. Where should the tip of the pipet compared to the liquid surface while liquid is being drawn into the pipet? While liquid is being let down to the line on the pipet stem? 3. Which of the following shows the proper adjustment of liquid level in a pipet so it will deliver its stated volume? 4. What do you do about any droplets or partial droplets of liquid on the outside of the pipet? 5. How much liquid should be left in the tip of the pipet after it gravity drains? 6. Where should you look to find the precision of the pipet (so you know how many decimal places to record in the volume)? 7. What should be done with the pipet before it is returned to the pipet storage container? (rinse it!) Volumetric Flask, Making and Diluting Solutions Use of a volumetric flask to make a solution: https://www.youtube.com/watch?v=Btp1N5Z2L74 • Video notes: o You can also dissolve a solid in a small beaker and then add it to the flask (but then rinse the beaker into the flask). o He has already rinsed the pipet well. o Invert the flask more times than he did. 1. George needs to make 250.0 mL of an aqueous solution. He weighs the appropriate amount of solute needed to make 250.0 mL of the solution, adds 250.0 mL of water to it, stirs to dissolve the solid and mixes it well. What did he do wrong? 2. What should you do with the volumetric flask before using it? 3. Which of the following shows the proper liquid level in a full volumetric flask that is being used for quantitative purposes? 4. At what points should you mix the solute and solvent by shaking the flask well? Why should you do this both before the liquid is filled to the line and then again after? Why should the flask be inverted during some of this mixing? 5. What should you do with the flask when you finish with it? 8 Buret Use of a buret: https://www.youtube.com/watch?v=9DkB82xLvNE • Video notes: o It says drain to 0 mL or below. Don’t waste your time getting to exactly 0.00 o You must also rinse the buret well when you are finished with it, including the tip. o You do not have to add partial drops in this CHEM 111. o The latter part of this video is about doing a titration (which is where burets are commonly used.) Carrying out a titration: https://www.youtube.com/watch?v=9DkB82xLvNE 1. Be able to describe the procedure used to rinse the buret properly. 2. What should be done about any air in the tip of the buret (below the stopcock)? When should this be done? 3. If a buret is calibrated with lines every 0.1 mL, how many decimal places should be recorded in the volume reading? 4. Why is it desirable to (a) not use a very small volume of titrant during a titration and (b) not use a large enough volume to need to refill the buret during one titration? 5. Sometimes a student misjudges and has to refill the buret during a titration. How do you decide if it is time to refill the buret? What do you do in the next trial to avoid having to refill the buret in it? 6. Your lab manual will often suggest (or assume that you know by now) first doing a small, sloppy, trial titration and scaling the volumes appropriately when you start the titration in which data is being carefully collected. Why not just be careful from the beginning and not do the sloppy titration? 7. What should you do with the buret when you are finished with it? 9 Academic Honesty & Lab Grades A Statement on Academic Dishonesty (written by D. Herron, former Chair, Dept. of Physical Sciences) “If science is to produce new knowledge about the world in which we live, it is essential that scientists adhere to the highest possible ethical standards. If scientists are forgiven for reporting false data, disregarding “bad” experiments, or making up d ata rather than performing experiments, it is impossible to trust any of their results. This goes deeper than abstract principles of justice and honesty; the whole scientific enterprise depends on intellectual honesty, and part of what you must learn is how important absolute honesty is in science. Consequently, any incident of a) falsified data, b) “dry lab” reports, or c) plagiarism of other peoples results will be treated as a serious breach of ethics and handled according. The penalty for such behavior will likely result in a minimum penalty of zero for the experiment at hand, and more serious penalties such as a directed grade of E for the course or a recommendation for expulsion from the university may be considered.” In Chemistry 112 labs, you are to write your own lab reports. This includes making your own graphs and tables as well as writing text. You are also responsible for keeping anyone else from using your graphs, tables and text, so don’t lend them to someone “just to see”. You may talk to your lab partner about your data, etc., but you should not write down things while you are doing it. If you write as you talk, your lab reports will sound too much alike and then your instructor is likely to accuse you of cheating. So, “I worked with him/her” is not an acceptable reason for reports being too much alike. The penalty for copying from someone or someone else copying from you will range from a zero on the report to an E in Chemistry 112. The laboratory grade counts for 15 % of the Chem grade. However you must pass both laboratory and lecture with the minimum grade to pass. This means that not passing either laboratory or lecture will result in a failing grade for the entire class. If you miss a lab with an excused absence you are expected to make the lab up or you will receive a zero. If you miss three labs without them making up or if you do not turn in the reports for the equivalent of three days of lab (which may or not be three lab reports), then you will receive an automatic E in Chemistry 112. I have read and understand the above rules; I agree to abide by them and I understand the penalties. _______________________ Signature _________________ Date Sign this sheet and turn it in to your instructor. 10 Lab Report Guidelines Generic Report Format The format below is general. Specific experiments may dictate some variations in report format, but all should follow this general form. Do not use a cover sheet. Top of first page: • Your name, your partner’s name if you worked with someone (collecting data) • Chem 112 • Lab day, time • Title of the experiment (title of what you did, not a generic title). For example, Analysis of Draino, not Analysis of Household Cleaner. Objective: This is one sentence stating exactly what specific task you did in this laboratory exercise. Procedure and Observations: A detailed description of the work you did. This includes all observations, generally mixed in with the procedure. If you are doing quantitative tasks and follow exactly the procedure in this manual, then just state that you did so; but you must still include observations and enough detail to tell when you saw various observations. Do not skimp on observations. If a change in the procedure is given, you must tell this. Give any procedure which you devise. Results: Data Tables: For data which differ from trial to trial but are collected using the same general procedure, use standard word processor tables or (if you need to do calculations on tabulated data) spreadsheet tables. The tables should be as compact as possible while being easy to understand. Do not break a table over two pages. All data must be clearly labeled. If any data is clearly incorrect (if there is reference book data available or if there are duplicate measurements), then indicate this in some way and then do not use it in calculations. Sample Calculations: If a calculation is based on a chemical reaction equation, write it down. Summarize the calculation with a “word equation” first, then show one calculation with numbers substituted in. Show one of each type of calculation this way; results of repeated calculations of the same type should simply be placed in a table and then the average and standard deviation should be reported for that type of calculation. Discussion and Conclusions: This section states what you found and how well you achieved your objective. Your conclusion should be directly related to the stated objective in your report. Leave out the filler (BS). You should support this only with specific data and observations which are included above in the report. Of course, you should also report specific experimental error(s), how they affected your results, and how the results might have been improved, but none of the “I might have misweighed” sort of errors which you have no evidence about. If possible, find the “correct” answer in a handbook or in a reliable on-line source (properly cite where you find it) and calculate a percent error. 11 Common Things to Watch Out for with Lab Reports 1. The writer should read the report after it is written and revise it as needed. The order of the report may or may not be the order in which you did things in lab. 2. Don’t write a week-by-week report (i.e. Don’t write “Week 1”, etc.) Write a report for one project unless the structure of the experiment is so you are really doing many unrelated activities. 3. Don’t write instructions of how to do the experiment. Write a report of what you did. The admonishment “Someone should be able to do the experiment from your report” means that you should have sufficient detail in your report, not that it should be a set of instructions. 4. Report repetitive data in tables. Describe a procedure done over and over only once and then list the samples for which it was used. 5. Pay attention to significant figures! 6. Organize your report so it is easy for the reader to follow what is done. For example, if step B of a procedure can’t be done until step A has been done, report A first so your reader doesn’t say to themselves “What about A?” and then later “Oh, there’s A!” A second example: data used in more than one table should never be simply referenced as “sample 1”, etc. in a manner that makes a reader have to look back-and-forth between the two tables. Sections of the report should be similarly organized—don’t put the conclusions before the data and analysis, for example. This is reader-unfriendly and shows poor communication skills. 7. Don’t connect the dots on the graph. 8. Look up appropriate handbook data to check your experimental data. Calculate averages from final answers of individual trials, not from raw data. (For example, calculate the molarity of each trial solution, then calculate an average molarity rather than averaging the milliliters used and calculating only one molarity.) If you have more than one answer, calculate an average and a standard deviation. 9. We will assume good lab technique for things like washing glassware (so don’t bother to tell us that you did it). 10. Run the spell checker on your report (though don’t believe everything the spell checkers suggest); in particular, the term absorbance is preferred to absorbancey and HCl is not Hcl. Use subscripts and superscripts where appropriate. 12 Making Graphs in Excel Directions for PC Computers Different computers have different defaults of what menus show, etc.. You might need to figure out how to get to the commands. While Microsoft Office products use the term chart, we may often use the term graph, which is more specific. I sometimes alternate below, depending on whether I am quoting what an Excel menu says or not. Getting the “raw” graph of one set of data: • Type your independent value (x) data in one column and your y-data in the column immediately to its right. If it is already typed but not side-by-side or in the wrong order, it is probably easiest to just cut-and-paste it into the correct order. • Select all of the data by dragging the cursor over it. • Choose Insert. From the icons that appear, choose Scatter and then from the popup menu, choose the icon that has only dots—no lines or curves. The graph should now appear, but you must edit it before turning it in! Required Editing of the “raw” graph: • Click in the middle of the graph to select it. • Make sure the background is white. If it is not, put the cursor somewhere on the graph but away from points and lines. Right click and choose format chart area. No Fill should be selected. • Make sure the graph has a title either above the graph or below the graph. The title should tell what is being graphed. An example would be Graph 1: Volume of Sample vs. Mass of Sample. No units are given in a title. With the graph selected, choose Layout Chart Title>Above Chart. Now you can select the generic title that has been inserted and edit it. • Axes should be labeled, and the label should include a unit unless the value being graphed has no unit (which is rare). With the chart area selected, choose Layout and then Axis Titles>Primary Horizontal Axis Title>Title Below Axis. Now choose Layout>Axis Titles>Primary Vertical Axis and either Vertical title or Horizontal Title (you might try both and see which looks better). Now the axes should have generic titles “Axis Title” shown on the graph. Click on these to edit them. • If there is only one type of data on the graph, you should remove the legend. Choose Layout>Legend>None. • Adding a Line of Best Fit and R 2 o If the points on the graph appear to fall pretty much along a straight line, then add a bestfit line: select Chart Tools>Layout>Trendline>Linear Trendline. Also, after this select Chart Tools>Layout>Trendline>More Trendline Options. Select Display Equation on Chart and Display R-squared value on chart. Editing that Might Be Needed • Zooming in on one part of the data. If Excel defaulted to showing the intersection of the axes at (0,0) and your data is considerably away from the 0 point, you should “zoom in” so the data points are distributed over most of the page. Occasionally you might want to show not only a general graph of the data, but you may want to have a second graph that shows only a section of the data. o To display only a section of the data based on the x values, choose Layout>Axes> Primary Horizontal Axis>More Primary Horizontal Axis Options. In the popup window that appears, for Minimum choose Fixed and then enter the value of the leftmost 13 • • gridline/axis tic mark that you want to appear on the x-axis of the zoomed graph. For Maximum, choose fixed and enter and the value of the rightmost gridline/axis tic mark that you want to appear. You may repeat this process with the y-axis if you desire by choosing Primary Vertical Axis in step one. Showing More Tic Marks on the Axes. Choose Layout>Axes> Primary Horizontal Axis>More Primary Horizontal Axis Options and Uner Major unit chose Fixed and then enter the distance (in values) that you want between the tic marks. If you want to display some tic marks without numbers, you may fill in the minor tic mark entry as well. Changing the number of decimal places shown in the numerical labels on the axes. Layout>Axes> Primary Horizontal Axis>More Primary Horizontal Axis Options. On the left side, choose Number. Then choose number in the right list. Enter the number of ecimal places in the box at the top right. o There are many other options available under Layout>Axes> Primary Horizontal Axis>More Primary Horizontal Axis Options. Directions for Macintosh Computers Different computers have different defaults of what menus show, etc.. You might need to figure out how to get to the commands. While Microsoft Office products use the term chart, we may often use the term graph, which is more specific. I sometimes alternate below, depending on whether I am quoting what an Excel menu says or not. Getting the “raw” graph of one set of data: • Type your independent value (x) data in one column and your y-data in the column immediately to its right. If it is already typed but not side-by-side or in the wrong order, it is probably easiest to just cut-and-paste it into the correct order. • Select all of the data by dragging the cursor over it. • Click on Charts > Scatter>Marked Scatter • Click in the middle of the graph, away from any numbers or symbols. • Chart Layout>Chart Title (choose above or below and then type in a title specific to the graph) • Chart Layout>Axis Titles>Horizontal Axis Title (then choose one and type in the x-axis title, including units) • Chart Layout>Axis Titles>Vertical Axis Title (then choose one and type in the x-axis title, including units) • If there is only one set of data, delete the Legend by clicking in it and selecting return • If the data is linear, add a best fit line by Chart Layout>Trendlines>Linear. Repeat Chart Layout>Trendlines and choose Trendline Options>Options, then click on Display Equation and Display R2 value. For each axis: • If the first point is very far from the number where the axis starts (excess white space), examine where the first point is numerically. Then double click right on the axis. In the box that comes up, choose scale and type in a whole number slightly less that that of the first point’s value on that axis. Double-click one of the points and unclick Shadow if it is on. 14 Putting More Than One Set of Data on a Graph Excel This will most commonly occur with the two sets of data that have the same x values and different y values, this is assumed in the instructions below. • Enter all the x values in one column, and each set of y values in its own column. So, for example, x values in column A, first y set in column B, second y set in column C, etc. • Now select all of the data by dragging over it and then do the steps as you would if you were graphing only one set of data and had the data selected. This should result in a graph where each set of data has a separate set of symbols and Excel know to treat each separately (when a line is added, for instance) • Generally if there is more than one set of data on the graph, you should keep the legend and label it appropriately. o Changing the labels in the Legend: choose Chart Tools>Design>Select Data. On the popup window, under Legend Entries select the name of the series that is being given in the legend (the one you want to re-label). Then choose Edit and under Series Name enter the name you want to appear on the legend. Click OK until you exit. 15 Experiment 1: Organic Structures & Nomenclature Introduction This exercise is designed to introduce you to the naming and structure of organic molecules. As you work through this exercise keep the following basic fact in mind: • IN SIMPLE ORGANIC COMPOUNDS, CARBON SHARES FOUR AND ONLY FOUR PAIRS OF ELECTRONS! Atoms are bonded in organic compounds as follows: • C shares four pairs and has no lone pairs. ALWAYS. • H shares one pair of electrons with any atom. ALWAYS. • O shares two pairs of electrons, and has two unshared pairs. ALMOST ALWAYS. • N shares three pairs of electrons, and has one unshared pair. ALMOST ALWAYS. • F, Cl, Br, and I share one pair of electrons and have three unshared pairs. ALMOST ALWAYS. NOTE: It is incorrect to refer to a pair of electrons that is in a double or a triple bond as “a bond”; that is two pairs of electrons in C=C is refered to as one bond, not two. Organic compounds are represented by chemical formulas. A molecular formula indicates the kinds of atoms and the number of atoms (e.g., C27 H46O), whereas a structural formula simply describes which atoms are bonded to which other atoms and how the atoms are bonded. The structural formulas are often written as some sort of condensed formulas (also called condensed structures). For example, a compound with the molecular formula C10 H22 may be written in the following ways: condensed structural formulas H H H H H H H H H H H C C C C C C C C C C H H H H H H H H H H H or CH3(CH2)8CH3 CH3CH2CH2CH2CH2CH2CH2CH2CH2CH3 or structural formula line (stick) structure Another possibility for C10 H22 is: CH3 CH3 CH3CHCH2CH2CH2CCH3 CH3 Notice that these two C10 H 22 compounds both have the same molecular formula, but the atoms are bonded differently. The first is a straight chain compound and the second is a branched chain compound. These compounds are called skeletal isomers of each other. There are 73 other isomers of C10 H22 ! 16 3-D drawings use a plain line for a pair of electrons in the plane of the paper, a solid wedge with its narrowest end on the central atom to indicate a bond pointing towards the viewer, and a dashed wedge to indicate a bond pointing away from the viewer: C H solid wedge dashed wedge When drawing 3-D diagrams of molecules two rules must be kept in mind: • If there are multiple center atoms, try to put as many bonds (atoms) as possible in the plane of the paper. Pick out a “row” of atoms that goes through the whole molecule and let it be in the plane of the paper; draw it with the correct angles, then add the other atoms to be backward and forward. o Example: For butane H H H H H C C C C H H H H H Select all the carbon atoms and one hydrogen atom from each end and draw it as follows: H C C C C H Next, add the forward and backward atoms as shown in the following structure: H H H H C C C H H C H H H H This is commonly drawn without showing any carbons: H H H H H H H H . • The bonds should be shown at approximately the correct angle in any size molecule. Keep in mind that if there are several tetrahedral arrangements in a row, the forward and backwards pointing bonds in a tetrahedral arrangement on the same atom are often shown to both point upwards or both down, not one up and one down. If there are several tetrahedral arrangements in a row, the up/down arrangement alternates. For example: H H H H H H H H H H H NOT H H H H H The same viewing angle should work for all molecules, and in the second drawing the angle shown for the first two carbons is not the same as that for the last. 17 Name ________________________________ Prelab 1. What is the prefix used for an organic compound where the longest carbon chain contains … a. 2 carbon atoms _______________ b. 4 carbon atoms _______________ c. 6 carbon atoms _______________ d. 8 carbon atoms _______________ 2. What is the name of an alkyl substituent with … a. 1 carbon atom _______________ b. 3 carbon atoms _______________ c. 5 carbon atoms _______________ d. 7 carbon atoms _______________ 3. What is the name of each of the following halogens as a substituent … a. fluorine _______________ b. chlorine _______________ c. bromine _______________ 18 Name ________________________________ Organic Structures & Nomenclature Lab Report Sheet Procedure Complete each of the following tasks and submit your report sheet at the end of the lab period. Use the provided model kits when directed. Note: These kits contain different colored balls for the atoms and holes to connect atoms with bonds. It is not that important if you use atoms of a particular color for particular atoms as long as the ball used has enough holes for each electron pair. The rigid sticks are used for single bonds and the flexible sticks must be used to make the double and triple bonds (one stick per pair of electrons). Part A. Draw Lewis structures of the molecules whose formulas are below and provide the name of each. Remember Lewis structures show lone pairs even on outer atoms. Next, convert each Lewis structure to a line structures. Note: In a line structure carbon and hydrogen atoms are not shown; each end point of a line represents a carbon atom and each carbon is assumed to have enough hydrogen atoms to fill carbon’s octet. Lastly, make models of each using the provided model kit. Using the model as a guide, draw a 3-dimensional picture of each. Be sure to draw your structure showing the correct bond angles! Formula CH3 Cl C2 H5 Cl Lewis Structure Name Line Structure 3-D Structure 19 C3 H8 Part B. The following are isomers of C 7 H16 . 1 2 3 4 Name the molecules represented above. 1. ______________________________ 2. ______________________________ 3. ______________________________ 4. ______________________________ The test of whether two organic compounds are the same or different is whether the two models can be superimposed exactly on each other. This is most easily tested using models. When you try to superimpose the models, you may do any twisting or turning of the single bonds you wish, but you may not break any "bonds” in the models. Which of the molecules represented below is/are the same as #2? __________ Which of the molecules represented below is/are the same as #3? __________ Which of the molecules represented below is/are the same as #4? __________ 20 Part C. There are four isomers with the formula C4 H9 Br. Make models of each one and complete the table below. Note that the –Br is a "bromo" substituent. Line Structure Name #1 #2 #3 #4 How are isomers different than isotopes? 21 Part D. Your instructor will assign you one of the four groups of compounds from the table below (#1, #2, #3, or #4). Write the name of each structure in your group right below the given line structures. Hint: NO 2 is a “nitro” substituent. #1 #2 #3 #4 (a) F F (b) Cl Cl F Cl Cl (c) F Br Br (d) Br (e) 22 Cl Experiment 2: Gas Chromatography, Intermolecular Forces, & Boiling Point Background Mixtures of many types are separated by the use of gas chromatography (GC), also called gas-liquid chromatography (GLC). This type of separation is based upon differences in boiling points as well as the intermolecular attractions (intermolecular forces) between the mixture components and the liquid that coats the walls of the column. The boiling point of a compound is related to the intermolecular attractions between its molecules. Often it is important to consider differences in the mass of molecules instead of simply the types of intermolecular forces, because molecules with larger masses are typically larger and thus experience stronger London forces. In GC, a tiny volume of sample is drawn into a syringe and then injected into the end of a heated tube (called a column) that has a liquid coating on the walls. This liquid coating is called a stationary phase. A gas (called the mobile phase) is flowing through the tube. The inside of the column on most modern instruments is hollow, and the sample and the mobile phase pass through the hollow part. The column is many meters long and has a diameter in the low micrometer range. As the sample moves through the column, the different components in the mixture repeatedly adsorb onto the stationary phase and then come out into the mobile phase. After a number of times of adsorption/desorption, the sample comes out the end of the column and goes into the detector. The lower the boiling point of a component, the less it interacts with the stationary phase and the more quickly it moves down the column and out into the detector. side of column from port to detector liquid coating (a) early (b) later side of column to detector quid coating (b) later As the molecules pass through the detector, there is a signal, resulting in a peak on the chromatogram. Even if all the molecules are identical in composition, some of them pass through the detector a tiny bit sooner than the others, so the peak grows grad ually and then decreases. Ideally, the peak should be symmetrical. If the mixture is separated well then each component has its own peak. The time it takes for the top of the peak to go through the detector is called the amount retention time and is used to tell which of various possible of signal compounds is causing that peak (from a legal viewpoint it is not enough to identify the compound, but if you know what time might be there as in our samples, then you can choose between the possible components). The percent of each component can retention time be calculated by dividing the area of its peak by the sum of the area of all the peaks present and then multiplying by 100. 23 Prelab Name ______________________________ Part 1: Gas Chromatography Watch the following videos and then answer the question below • https://www.youtube.com/watch?v=3AQ55RPVE_A • https://www.youtube.com/watch?v=4Xaa9WdXVTM o Note that the Kahn Academy video discusses a difference in mass of the molecules with the same boiling point, but this is really differences in intermolecular attractions. Be aware that there are some differences in different types of chromatographs, such as types of detector, and whether you have to make a solution of the sample or use it pure. 1. What information is obtained from the retention time of a peak in a chromatogram? 2. What information is calculated from the area of the peaks in a chromatogram? 3. If your sample contains three compounds and there is no solvent, how many peaks should you expect in the chromatogram? 4. What does volatile mean? 5. If a sample is made using 0.145 g of methanol and 0.554 g of ethanol, (a) What percent of the total peak area is expected for the methanol peak? ππππ ππ πππ‘βππππ ππππ ππππ ππ πππ‘βππππ ππππ + ππππ ππ ππ‘βππππ ππππ × 100 Show work. (b) Look up the boiling points of the methanol and the ethanol and tell which comes out first (of the two). Explain. 24 Part 2: Intermolecular Forces & Boiling Point Review organic functional groups and intermolecular forces in your lecture text (sections 23.2 and 11.1) and then answer the following questions. 1. What is the suffix used when naming compounds with each of the following functionalities? a) a hydrocarbon with only carbon-carbon single bonds __________ b) a hydrocarbon with a carbon-carbon double bond __________ c) a hydrocarbon with a carbon-carbon triple bond __________ d) an alcohol __________ e) an ester __________ 2. Provide the Lewis structure for each of the following functional groups (only the functional group, not a specific molecule): a) an alcohol b) an ester 3. For each of the following compounds, draw line structures of two molecules and use dotted lines to indicate what atoms of the two molecules are attracted the most, i.e., use dotted lines to show the strongest intermolecular attraction between the molecules. a) ethanol b) ethyl ethanoate (ethyl acetate) 4. Identify what type(s) of intermolecular forces are present between molecules of: a) a hydrocarbon β‘ London dispersion β‘ dipole-dipole β‘ hydrogen bonding b) an alcohol β‘ London dispersion β‘ dipole-dipole β‘ hydrogen bonding c) an ester β‘ London dispersion β‘ dipole-dipole β‘ hydrogen bonding 5. Considering compounds with similar molar masses but a different functionality, which is likely to have the … a) lowest boiling point β‘ hydrocarbon β‘ alcohol β‘ ester b) highest boiling point β‘ hydrocarbon β‘ alcohol β‘ ester 25 Procedure Part 1: Gas Chromatography Your lab instructor will demonstrate how to do an injection properly using a standard mixture of four known alcohols. You will need the chromatogram of the standard to help identify the components in your unknown mixture (this will be discussed in lab). Each group will be assigned an unknown mixture and have a turn using the GC. When it is your turn, be very careful with the syringe so you do not bend the plunger! The best way to prevent this is to not draw the plunger very far out, as you won’t need to do so. Be sure there are no air bubbles in the liquid before you inject it! While your group is awaiting your turn on the GC (and/or after your turn on the GC) you will work on Part 2. Part 2: Intermolecular Forces & Boiling Point Use the resources provided in lab to complete the following table; for each of the given compounds identify the type of compound then provide the condensed structural formula, the molar mass, and the boiling point. Name Type of Compound butanol (butyl alcohol) β‘ hydrocarbon β‘ alcohol β‘ ester ethyl ethanoate (ethyl acetate) β‘ hydrocarbon β‘ alcohol β‘ ester β‘ hydrocarbon β‘ alcohol β‘ ester β‘ hydrocarbon β‘ alcohol β‘ ester β‘ hydrocarbon β‘ alcohol β‘ ester β‘ hydrocarbon β‘ alcohol β‘ ester β‘ hydrocarbon β‘ alcohol β‘ ester β‘ hydrocarbon β‘ alcohol β‘ ester β‘ hydrocarbon β‘ alcohol β‘ ester β‘ hydrocarbon β‘ alcohol β‘ ester β‘ hydrocarbon β‘ alcohol β‘ ester β‘ hydrocarbon β‘ alcohol β‘ ester β‘ hydrocarbon β‘ alcohol β‘ ester β‘ hydrocarbon β‘ alcohol β‘ ester β‘ hydrocarbon β‘ alcohol β‘ ester heptane octane pentane pentanol propyl ethanoate (propyl acetate) butyl ethanoate (butyl acetate) hexanol hexane methyl ethanoate (methyl acetate) heptanol pentyl ethanoate (pentyl acetate) octanol nonane Condensed Structural Formula 26 Molar Mass (g/mol) Boiling Point (°C) Report Part 1: Gas Chromatography • Provide your chromatographs of the standard mixture and your unknown mixture with the peaks labeled identifying which alcohol correspond to each peak. • Calculate the percentage of each alcohol in your unknown mixture (show your work directly on your unknown chromatograph). Part 2: Intermolecular Forces & Boiling Point • Transfer the name, type of compound, molar mass, and boiling point data found in lab into an Excel spreadsheet. o Organize the data by functional group, i.e., group all of the hydrocarbons together, all of the alcohols together, and all of the esters together. • Graph the boiling point vs. molar mass as three sets of data on one graph (one for hydrocarbons, one for alcohols and one for esters). o All of this data should go on one graph, but as three sets of data, not one set. o If you do not know how to put multiple data sets on one graph, see the graphing instructions at the beginning of the lab manual. • Put a best fit line through each of the data sets (so you should have three lines) and display the equation for each of the three lines. o The unit on the slopes are the (units of y/units of x); in this case it should be the change in boiling point per one g/mol change in the molar mass. o Hint: The three slopes should help you answer the first three questions below; the molar mass increases by ~14 g/mol for each added CH 2 group. • Transfer your table(s) of data and the graph of boiling point vs. molar mass to a Word document. • Answer the questions below in the same Word document and submit a copy of the document along with the items from Part 1. 1. On average, how much is the boiling point of a hydrocarbon changed by adding one CH 2 group (increasing the molar mass by 14 g/mol)? Hint: You must use the slope from your hydrocarbon data. 2. On average, how much is the boiling point of an ester changed by adding one CH 2 group? 3. On average, how much is the boiling point of an alcohol changed by adding one CH 2 group? 4. Within the range of the molar masses of the substances used in this experiment, does the trend of boiling points follow what the textbook would predict (substances whose molecules can hydrogen bond boil at higher temperatures than substances with polar molecules that cannot hydrogen bond, which are in turn higher than substances with nonpolar molecules)? Briefly explain using your data (examine your graph to compare alcohols, esters, and hydrocarbons that have similar molar masses). 27 Experiment 3: Net Ionic Equations Introduction In this exercise you will be combining solutions of ionic compounds, acids, and bases and observing the results. Throughout the process you will make observations of the initial solutions, thoroughly mix the two solutions, observe any immediate changes that occur, and observe the final state of the mixture. Then you will identify the ions that were initially present in the two solutions that were mixed and write the net ionic equation that is consistent with what you observed (or indicate if no reaction occurred). Although you will not be required to write out the molecular equation and the ionic equation for these reactions, it can be helpful in determining the net ionic equation (especially if you are new to or haven’t had much practice writing out net ionic equations previously). You should review section 4.2 in your lecture textbook if you need a review of how to write molecular and ionic equations. You should also have the solubility rules from your lecture textbook (Tables 4.2 and 4.3) memorized before coming to lab so that you can successfully distinguish between products that should be soluble and products that should be insoluble. Example Data Analysis The example below summarizes the kind of data analysis we expect. Note that you need to write the formulas of the substances, not the 1B, etc. Example Solutions mixed: 0.1 M CoSO 4 and 0.1 M NaOH Appearance of initial solutions: The CoSO 4 solution was pink and the NaOH solution was colorless. Appearance after mixing: Upon mixing we could see a blue solid and the liquid was pink. Starting solutions contain: πΆπ 2+ , ππ42− , ππ +, ππ» − Net Ionic Equation: πΆπ 2+ (ππ) + 2 ππ» − (ππ) → πΆπ(ππ»)2 (π ) Example Notes: To determine the net ionic equation, you must be able to predict the products of the reaction and identify any insoluble products (or other weak or non-electrolytes formed, such as water) because these will be part of the net ionic equation along with any ions that are necessary in forming them. Alternatively you can first predict the products and write the reaction in the molecular form, then convert it to the complete ionic equation, and finally cancel the spectator ions to get the net ionic equation. For the above example the following are the corresponding molecular and ionic equations that would help lead you to the net ionic equation shown above. πΆπππ4 (ππ) + 2 ππππ»(ππ) → πΆπ (ππ»)(π ) + ππ2 ππ4 (ππ) πΆπ 2+ (ππ) + ππ42− (ππ) + 2 ππ+ (ππ) + 2 ππ» −(ππ) → πΆπ (ππ»)2 (π ) + 2 ππ+ (ππ) + ππ42− (ππ) 28 Name ______________________________ Prelab Memorize the solubility rules provided in Table 4.2 and 4.3 found in section 4.2 of your lecture ebook. At least for now you should know the main rules (the information on the left of each table), within about the next month it will become important that you have both the main rules and exceptions memorized. You should also review how to predict products of reactions between ionic compounds, predict products of reactions between ionic compounds as well as acids and bases, and how to write reaction equations in the molecular, ionic, and net ionic forms. All of these things can be found in sections 4.2 and 4.3 of your lecture ebook. Utilize this information to answer the following prelab questions. _____ 1. Based on the solubility rules, which one of these compounds is insoluble in water? A) NaCl B) MgBr2 C) FeCl2 D) AgBr E) ZnCl2 _____ 2. Based on the solubility rules, which one of these compounds is insoluble in water? A) Na2 SO4 B) BaSO 4 C) CuSO 4 D) MgSO 4 E) Rb2 SO 4 _____ 3. Based on the solubility rules, which one of these compounds is soluble in water? A) Hg2 Cl2 B) Na2 S C) Ag2 CO 3 D) Ag2 S E) BaCO 3 _____ 4. Based on the solubility rules, which one of these compounds is soluble in water? A) AgBr B) AgCl C) Ag2 CO 3 D) AgNO 3 E) Ag2 S _____ 5. If aqueous solutions of Na2 CO 3 and BaCl2 are mixed, which insoluble precipitate is formed? A) Ba2 CO3 B) BaCO 3 C) NaCl D) NaCl2 E) BaO _____ 6. If aqueous solutions of Mg(C 2 H3 O2 )2 and LiOH are mixed, which insoluble precipitate is formed? A) LiC2 H3 O2 B) Li(C2 H3 O 2 )2 C) MgOH D) Mg(OH)2 E) CH 3 OH _____ 7. If aqueous solutions of Pb(NO 3 )2 and NaCl are mixed, which ions, if any, are spectator ions? Pb(NO 3 )2 (aq) + 2 NaCl(aq) → PbCl2 (s) + 2 NaNO 3 (aq) A) Pb2+(aq), Cl– (aq) D) Na+(aq), Cl– (aq) B) Na+(aq), NO 3 – (aq) E) None are spectator ions. C) Pb2+(aq), NO 3 – (aq) _____ 8. Identify the net ionic equation for the following reaction: 2 NaCl(aq) + Hg2 (NO 3 )2 (aq) → 2 NaNO 3 (aq) + Hg2 Cl2 (s) A) Na+(aq) + NO 3 – (aq) → NaNO 3 (aq) B) Hg2 2+(aq) + 2 Cl– (aq) → Hg2 Cl2 (s) C) 2 Hg+(aq) + 2 Cl– (aq) → 2HgCl(s) D) Hg2 (NO 3 )2 (aq) → Hg2 2+(aq) + 2 NO 3 – (aq) E) Hg2 2+(aq) → Hg2 (s) 29 _____ 9. What is the net ionic equation for the reaction that occurs when aqueous solutions of Pb(NO 3 )2 and NH 4 Cl are mixed? A) Pb(NO 3 )2 (aq) + 2 NH 4 Cl(aq) → 2 NH 4 NO3 (aq) + PbCl2 (s) B) Pb2+(aq) + 2 Cl– (aq) → PbCl2 (s) C) Pb2+(aq) + 2 NO 3 – (aq) + 2 NH 4 +(aq) + 2 Cl– (aq) → 2 NH 4 +(aq) + 2 NO 3 – (aq) + PbCl2 (s) D) NH 4 +(aq) + NO 3 – (aq) → NH 4 NO 3 (s) E) NH 4 +(aq) + NO 3 – (aq) → 2 NO(g) + 2 H 2 O(l) _____ 10. Complete the following reaction and identify the Brønsted acid. NaOH(aq) + HCl(aq) → A) NaH(aq) + HOCl(aq); NaOH is the acid. B) NaH(aq) + HOCl(aq); HCl is the acid. C) NaCl(aq) + H2 O(l); NaOH is the acid. D) NaCl(aq) + H 2 O(l); HCl is the acid. E) NaCl(aq) + H 2 O(l); NaCl is the acid. _____ 11. If aqueous solutions of Ba(OH)2 and HNO 3 are mixed, what products are formed? A) BaN 2 (s) + H 2 O(l) C) Ba(s) + H 2 (g) + NO 2 (g) E) Ba3 N2 (s) + H 2 O(l) B) Ba(NO 3 )2 (aq) + H 2 O(l) D) Ba2 O(s) + NO 2 (g) + H 2 O(l) _____ 12. Which salt is produced by the neutralization of hydrobromic acid with magnesium hydroxide? A) MgBr B) Mg2 Br3 C) MgBr2 D) Mg3 Br2 E) Mg2 Br D) HBr E) HCl _____ 13. Which of the following is a weak acid? A) H 2 SO 4 B) HNO 3 C) HF _____ 14. Which of the following is a strong acid? A) H 3 PO 4 B) HNO 3 C) HF D) CH 3 COOH E) H 2 O _____ 15. Which is the net ionic equation for the reaction between aqueous solutions of lithium hydroxide and hydrobromic acid? LiOH(aq) + HBr(aq) → H 2 O(l) + LiBr(aq) A) LiOH(aq) → Li+(aq) + OH – (aq) B) HBr(aq) → H +(aq) + Br– (aq) C) H +(aq) + OH – (aq) → H 2 O(l) D) Li+(aq) + Br– (aq) → LiBr(aq) E) Li+(aq) + OH – (aq) + H +(aq) + Br– (aq) → H 2 O(l) + LiBr(aq) 30 Procedure PRECAUTIONS: ALMOST ALL THE SOLUTIONS USED IN THIS EXERCISE ARE CORROSIVE. ANY SPILLS MUST BE DILUTED WITH WATER AND WASHED COMPLETELY AWAY IMMEDIATELY. Mixtures containing nickel, copper, or barium must be disposed of in the bottle labeled "Heavy Metal Recovery" At the start of lab your instructor will assign you one of the following sets: (a) (b) (c) (d) (e) (f) (g) (h) (i) (j) 1A, 1C, 2B, 3F, 4D, 5C, 6G, 8I, 11E, 11J, 12J, 13I 1B, 1C, 2A, 3F, 4G, 6F, 8D, 10J, 11I, 12I, 13E, 13J 1A, 1D, 2C, 3B, 4G, 8F, 9G, 10I, 11J, 12J, 13E, 13I 1B, 1D, 2C, 3A, 4G, 8F, 9D, 10J, 11E, 11I, 12I, 13J 1A, 1F, 2D, 3C, 4B, 5I, 9C, 10G, 11E, 11J, 12J, 13I 1B, 1F, 2G, 3C, 4A, 5D, 6I, 10F, 11J, 12J, 13E, 13I 1A, 1G, 2B, 3D, 4C, 5F, 6G, 9J, 11I, 12I, 13E, 13J 1B, 1G, 2A, 3D, 4C, 6F, 8D, 10I, 11E, 11J, 12J, 13I 1A, 1F, 2F, 3B, 4G, 6C, 8G, 9J, 11E, 11I, 12I, 13J 1B, 1G, 2F, 3A, 4D, 8C, 9D, 10I, 11J, 12J, 13E, 13I Each code in your assigned set corresponds to two chemicals (one for the number and one for the letter; see the table below) that you will combine, observe, and identify the reaction that occurs (if any). For example, the code 8J means that you should combine a 0.1 M solution of FeSO 4 with a 0.1 M NaOH solution and carefully observed the result. 1. 0.1 M MgCl2 2. 0.1 M CaCl2 3. 0.1 M SrCl2 4. 0.1 M BaCl2 5. 0.1 M ZnSO 4 6. 0.1 M CoCl2 8. 0.1 M FeSO 4 9. 0.1 M CuSO 4 10. 0.1M MnSO 4 11. 0.1 M HCl 12. 0.1 M CH 3 COOH 13. 0.1 M HNO3 A. 0.1 M Na2SO 4 B. 0.1 M (NH 4)2SO 4 C. 0.1 M Na3PO4 D. 0.1 M Na2CO3 E. NH 3 F. 0.1 M K 3PO4 G. 0.1 M K 2CO3 I. 0.1 M KOH J. 0.1 M NaOH For each of your assigned sets you should obtain about 1 mL of each solution in two separate test tubes. Then write down observations of the initial appearance of each solution. Pour one solution into the other and gently thump the side of the tube near the bottom to mix the solutions. Write down initial observations of the mixed solutions. Set the tube aside for 10 minutes to ensure the reaction is complete and to allow any precipitate to settle. Write down final observations of the mixed solutions. Lastly, write down the net ionic equation for the reaction that occurred (this should be consistent with your observations). It may be helpful to have some scratch paper to first write out the molecular and ionic forms of the reaction before simplifying it to a net ionic equation. 31 Note: You will be assigned some combinations of solutions that will result in an acid-base reactions. There will not be any change in appearance that occurs as a result of these reactions, but that does not mean that there was no reaction. If the mixture does not produce a precipitate you should determine whether or not an acid-base reaction could have happened which would produce the nonelectrolyte water and therefore have a net ionic equation. Please note each of the following tips: 1. The test tubes you use (13 ο΄ 100 mm) hold almost exactly 10 mL. Measure by estimating a fraction of the test tube volume, not by using graduated cylinders, etc. 2. You should never have more than 3 mL total in a test tube so that you can efficiently mix the contents of a test tube without stirring (“thump” the tube). 3. When mixing liquids in this exercise, use around equal volumes of the two liquids, keeping the total volume below 3 mL. Observe what happens when you mix them. Allow them to stand for at least 10 minutes before you decide about the final state. Fill out the included report sheets as you perform the lab; you will be expected to submit your completed report sheets before you leave the lab. See your instructor if you have problems 32 Name: Lab Day and Time: Solutions Mixed: Appearances of Initial Solutions: Appearance After Mixing and Centrifuging (if Needed) Ions Present: Net Ionic Equation: Solutions Mixed: Appearances of Initial Solutions: Appearance After Mixing and Centrifuging (if Needed) Ions Present: Net Ionic Equation: Solutions Mixed: Appearances of Initial Solutions: Appearance After Mixing and Centrifuging (if Needed) Ions Present: Net Ionic Equation: Solutions Mixed: Appearances of Initial Solutions: Appearance After Mixing and Centrifuging (if Needed) Ions Present: Net Ionic Equation: 33 Name: Lab Day and Time: Solutions Mixed: Appearances of Initial Solutions: Appearance After Mixing and Centrifuging (if Needed) Ions Present: Net Ionic Equation: Solutions Mixed: Appearances of Initial Solutions: Appearance After Mixing and Centrifuging (if Needed) Ions Present: Net Ionic Equation: Solutions Mixed: Appearances of Initial Solutions: Appearance After Mixing and Centrifuging (if Needed) Ions Present: Net Ionic Equation: Solutions Mixed: Appearances of Initial Solutions: Appearance After Mixing and Centrifuging (if Needed) Ions Present: Net Ionic Equation: 34 Name: Lab Day and Time: Solutions Mixed: Appearances of Initial Solutions: Appearance After Mixing and Centrifuging (if Needed) Ions Present: Net Ionic Equation: Solutions Mixed: Appearances of Initial Solutions: Appearance After Mixing and Centrifuging (if Needed) Ions Present: Net Ionic Equation: Solutions Mixed: Appearances of Initial Solutions: Appearance After Mixing and Centrifuging (if Needed) Ions Present: Net Ionic Equation: Solutions Mixed: Appearances of Initial Solutions: Appearance After Mixing and Centrifuging (if Needed) Ions Present: Net Ionic Equation: 35 Experiment 4: Rate Law for Bleaching of Crystal Violet by NaOH ***You must bring a flash drive to lab to save your data files.*** Objective The purpose of this experiment is to determine the rate law for the reaction of the compound crystal violet (methyl violet) with ππ» − ion. To accomplish this you will need to find the reaction order with respect to crystal violet, the reaction order with respect hydroxide ion, and the value for the rate constant. Introduction When hydroxide ion is added to a crystal violet solution, the solution becomes colorless as the crystal violet reacts. You can follow the reaction by placing a reaction mixture in a spectrophotometer and observing the change in absorbance with time. When carrying out this type of experiment you first need to find the wavelength where the maximum absorbance occurs (λ max ). It is also important that you find a concentration with an appropriate absorbance for the rate measurements (not too high for the instrument and not so low that the error in the measurements is excessive). This is done by finding the absorbance at λmax of a solution whose concentration is known, and since absorbance is proportional to concentration the necessary scaling factor for the concentration can be determined from the ratio of the absorbances. Once absorbance versus time data has been collected it can be plotted to help determine the order of a reaction with respect to the absorbing substance. Although this analysis is often done with concentrations, absorbance can be substituted for concentration since they are proportional. Therefore, when a plot of the absorbance versus time is linear the reaction is zero order, when a plot of ln(absorbance) versus time is linear the reaction is first order, and when a plot of 1/absorbance versus time is linear the reaction is second order. The plots below show you each case for data from a zero, first, and second order reaction; notice that plots with a mismatched order have curvature, e.g., when first and second order data are plotted as absorbance versus time the result is a curve whereas the zero order data is linear. 36 Prelab Consider the following data collected for the decomposition of substance A: Time (min) 0 60.0 96.4 157.5 • • • • • [A] (M) 0.50 0.40 0.35 0.28 Prepare the three plots necessary to test whether the data is zero, first, or second order, i.e., [A] vs. t, ln[A] vs. t, and 1/[A] vs. t. Add a trendline to each plot and show the equation and R 2 value on each plot. Provide a copy of your plots. Identify whether the reaction is zero, first, or second order. Report the rate constant for this reaction (be sure to include units with the value). 37 Procedure The stock solution of crystal violet that is provided for you is an appropriate concentration for the absorbance measurements, so you will not need to dilute the crystal violet solution. Make note of the concentration of the crystal violet solution provided. Set up and calibrate your spectrophotometer: β‘ Plug in the LabQuest 2 device (LQ2). β‘ Attach the LQ2 to the cube spectrometer β‘ Power on the LQ2. After an initial startup screen, a screen with a red stripe that shows USB:Abs should appear (this is the home screen). If you don’t see this ask for help. β‘ At the upper right hand corner there should be a gray box with Mode:Full Spectrum, which is what you need for this part. If you don’t see this ask for help. β‘ Choose Sensors at the top left of the screen. β‘ Choose Calibrate -> USB: Spectrometer near the bottom of the gray menu. A new screen should now appear, indicating the time necessary for warm up. β‘ While the spectrometer is warming up, place a clean cuvette filled about 2/3 full with DI water (a blank) into the cell holder of the spectrometer. β‘ Once the spectrometer is warmed up and your blank is in place, choose Finish Calibration. β‘ When the calibration is complete, choose OK to go to the original screen. Obtain the full absorbance spectrum of the crystal violet solution from 400-800 nm and save the spectrum: β‘ β‘ β‘ β‘ Empty your cuvette and shake out the excess water. Fill the cuvette about 2/3 full with the crystal violet solution. Place the cuvette into the spectrometer. On the LQ2 screen, tap the green triangle at the bottom left. Data acquisition should start and the screen should change to show a full spectrum of your sample. β‘ When the black line with circles stops changing (this should happen almost instantly), tap the small box at the left bottom with the red square to stop data collection. β‘ Save this file. o Select File in the upper left corner and then Save. o Touch the field at the top with the stylus and then change ‘untitled’ to your last name and a number (e.g. Miller1, Miller2, etc.; be sure you know what each file name corresponds to). o Tap Done and then Save in the bottom right corner. β‘ Empty the cuvette, rinse with DI water, and shake out the excess water so it is ready for your first kinetics trial. Use the full absorbance spectrum to determine the wavelength of maximum absorbance, λ max , in this wavelength range. Set up the spectrophotometer to collect absorbance versus time data at the λ max : β‘ In the upper left corner, choose the icon that looks like a meter; this will take you to the screen with the red stripe. β‘ In the upper right corner, choose the mode square. β‘ In the drop down menu at the top, choose time based. β‘ Tap the field next to Duration, change the number to 1800, and tap Done. Then make sure the unit is set to s for seconds. β‘ Choose OK. β‘ It will ask about your data and go ahead and select Discard since you already saved it. β‘ This will take you back to the screen with the red stripe. β‘ Tap the red stripe and select Change Wavelength. β‘ Input the λmax into the Selected Wavelength field and then tap OK. 38 Now, you are ready to perform your first kinetics trial. Accurately pipet a small volume of the diluted crystal violet stock solution into a small, dry beaker (2-3 mL is sufficient). Pipet the same number of milliliters of 0.5 M NaOH into another small, dry beaker. Rapidly pour one of the beaker’s contents into the other, quickly but carefully swirl the contents to mix, and then quickly fill your cuvette about 2/3 full with the mixture using a transfer pipet, immediately place the cuvette in the spectrophot ometer and tap the green triangle at the bottom left of the screen to start collecting data. If the initial absorbance is less than 0.5 (you can read this from the live plot that is displayed or look at the absorbance reading shown near the top right of the screen), then you did not do the mixing, transfer, and start of data collection quick enough and you must start the trial over. Once you have successfully started a trial that has an initial absorbance of 0.5 or greater you must allow the data collection to continue until the absorbance drops below one tenth of its initial absorbance, e.g., if the initial absorbance is 0.5 then you must continue data collection until the absorbance drops to 0.05 or if the initial absorbance is 0.7 then you need to continue until the absorbance drops to 0.07, etc. Tap the button with the red square to stop the collection early if the initial absorbance was too low or once the absorbance has dropped to below one tenth of the initial absorbance. If you have to start a new trial it will ask you about discarding the data after you tap the green triangle to start a new collection, in this case you should tap Discard because you don’t want to add your new trial to the previous one. Once you have successfully collected data that meets these criteria stop the collection and save the data. Next you need to perform three more kinetics trials with diluted concentrations of NaOH (something less than 0.5 M), saving each data set after it is collected. You get to choose the diluted concentrations of NaOH that you will use, but your lowest should be somewhere in the range of 0.05-0.1 M and your four NaOH concentrations (including the 0.5 M) should be equally spaced (or close to equally spaced). Follow the same steps as the previous trial for each of these trials: tap the green triangle button to st art data collection once you have inserted the cuvette, discard the previous data that is already saved, tap the red square button to stop data collection, and save the data. Note: All of the trials must meet the same absorbance criteria as the first (initial absorbance of at least 0.5 and final absorbance that is lower than one tenth of the initial absorbance). Lastly, you need to run a duplicate of one of your previous trials and save the data; if there is a trial you feel you may have made an error on then you should duplicate that trial. Follow the same steps as before. When you are done you will have five trials where four will have different NaOH concentrations and one will be a repeat. Export your saved absorbance vs. wavelength data set and your five absorbance vs. time data sets to a USB drive: β‘ β‘ β‘ β‘ β‘ β‘ β‘ β‘ β‘ Unplug the spectrometer from the USB port on the LQ2. Plug your USB drive into the USB port on the LQ2. Choose File -> Open. In the upper left hand corner choose the farthest left icon (this will show you the LQ2 file directory instead of your USB file directory). Choose the file you want to save (you have to open the file in order to be able to save it to your USB). This will open the file, but it will not go to the screen that shows the graph (while you could go back to the graph, it is not necessary to export the file). Choose File -> Export (NOT Save). Type a name for your file (it does not automatically pull the name that you gave it on LQ2). Tap Done, then OK. Now the data has been transferred as a text file to your USB drive. Repeat the File -> Open, File -> Export process for each of your data files. Either in lab or as soon as you get to a computer, check that you can import your data into Excel. DO NOT wait until the next week to check, otherwise if you cannot open your data then you may not have time to get it so you can complete your lab report. Turn off the LQ2, unplug everything and put it back how you found it. 39 Data Analysis β‘ The end goal is to determine the rate law for the reaction, including the reaction order with respect to crystal violet, the reaction order with respect to hydroxide ion, and the rate constant for the reaction, i.e., to find m, n, and k in rate = k[crystal violet]m[OHβ»]n . β‘ Find the reaction order with respect to crystal violet, m. o Each kinetics trial needs to be analyzed to determine if the reaction is 0, 1 st, or 2nd order with respect to crystal violet. o For each kinetics trial, make a plot of absorbance vs. time, a plot of ln(absorbance) vs. time, and a plot of 1/absorbance vs. time to test for 0, 1 st, or 2nd order, respectively. βͺ You should end up with three plots for each of your five kinetics trials, giving you a total of fifteen plots. βͺ Note: Normally concentration is used to do this, but since concentration and absorbance are directly proportional the absorbance can be used in place of concentration; the same graph will be a straight line if we use absorbance instead of concentration. β‘ Find the apparent rate constant, k’, for each trial, where k’ = k[OHβ»]n . o You can’t find the actual rate constant, k, directly because the reaction order of the hydroxide ion is not yet known, therefore first you must find the apparent rate constant, k’. o Look at your plots for each run, and decide which is the straightest for the majority of runs. This indicates whether the order with respect to crystal violet is 0, 1 st, or 2nd order. o When you have decided on the most linear set, use that set to determine k’ for each trial. Hint: The rate constant is always equal to the magnitude of the slope, so k’ is the absolute value of the slope of the linear plots. o Note: Since you are using absorbance values instead of concentration, the rate constant will be in terms of the absorbance instead of the concentration. β‘ Find the reaction order with respect to [OHβ»], n, and the rate constant, k. o Make a table of k’ and the corresponding [OHβ»] values. o Use the values from the above table to make a second table with ln(k’) and ln[OHβ»]. o Make a plot of ln(k’) vs. ln[OHβ»], and add a trendline and its equation to the plot to find n and k. βͺ Recall that k’ = k[OHβ»]n . βͺ By taking the ln of the above, it follows that ln(k’) = n ln[OHβ»] + ln(k). βͺ Thus, n is the slope of the plot and k can be calculated from the y-intercept of the plot (k = ey-intercept). 40 Report Your report must include each of the following items in this order. Your instructor will notify you of any additional report expectations. • • • • • • • • • • Solution data: concentration of crystal violet solution, concentrations of hydroxide ion for each trial, and sample calculation for the dilutions. The absorbance vs. wavelength plot and a statement of the λ max determined from it. Kinetic plots for each of the five trials. o These should be presented on five different pages, with the three plots for a given trial on its own page so they can be compared easily. o Each page must be clearly labeled with the trial’s hydroxide concentration. o Each page should include a trendline and its equation on only the most linear plot for that trial. An explicit statement of the order of the crystal violet and how it was determined. A table of k’ vs. [OHβ»] and a table of ln(k’) vs. ln[OHβ»]. The graph of ln(k’) vs. ln[OHβ»] with the trendline and its equation. An explicit statement of the order of the hydroxide ion and how it was determined. The calculation of numerical value of k. A conclusion stating the rate law in as explicit a form as you can (meaning with the actual values for k, m, and n). A conclusion about the uncertainty in your measurements. Hint: The k’ for the runs that you duplicated should help in estimating this uncertainty. 41 Experiment 5: Le Chatelier’s Principle Purpose This exercise is intended to give you first-hand experience observing chemical equilibria, how changing intensive variables can affect the direction of the reaction, and how the latter relates to Le Chatelier’s Principle (LCP). Introduction LCP states that if you start with a system at equilibrium and change one of its intensive properties, then the equilibrium will shift to offset that change in intensive property. As a result of that shift, we may see other changes. In this experiment, you will often use these other observed changes to deduce which way the equilibrium shifted. In some chemical reactions, the product will have a different color than the reactant. In this case, if the equilibrium is disturbed, a resulting color that is more product-like indicates a higher concentration of product and thus that the equilibrium shifted to the right. If the color becomes more reactant-like, then there is a higher concentration of reactant and the equilibrium must have shifted left. Similar changes in equilibrium position are indicated by solids forming or dissolving. Experiment Hints • • • • • • • Look down through the top of the test tube to observe color changes; observing this way will minimize visual changes due only to dilution. To see that this is true, consider if you doubled the volume of a solution by adding water and nothing else happened. Then looking from the top of the tube would be looking through twice as much solution that is half as dark, which would look the same color intensity as the original. If an equilibrium shifts due to a dilution, it is NOT due to the change in concentration of water! In this case, you will need to consider the total solute concentration and its changes as the intensive property. If a reaction is endothermic, heat may be regarded as a reactant and if a reaction is exothermic, heat is a product. For this situation, an increase in temperature increases the “concentration” of the heat. Write careful notes about what you see for each part of each of the experiments on the provided report sheets. At each step, including before mixing reagents: o Report the colors observed. Your observations of changes during dropwise additions should be summarized rather than reported “drop by drop”. o Clearly indicate what your observations imply about species present and what the changes in species present imply about the direction of the equilibrium shift of the main reaction equation (ME). o Trace changes from the added reagent to the changes in intensive properties at each step of the way until you reach the LCP predicted shift in the main equilibrium. o Do not skip steps! Note that LCP cannot be applied until an equilibrium has been established and sometimes you must mix things to establish the equilibrium before disturbing it. Be sure it is absolutely clear which are observations and which are deductions based on the observations and the equation. Be careful with each and every section; proper use of terminology is extremely important! Do not say things like “LCP predicts if the reactant concentration increases then the equilibrium will shift to produce more product”. LCP does not say anything about making product—it simply says the equilibrium would shift to decrease the concentration of the added reactant. 42 Example Report Format Main chemical equilibrium equation: Useful reaction equation: π΅π2+ (ππ) + πΆππ42− (ππ) → π΅ππΆππ4 (π ) yellow yellow ppt. 2 πΆππ42− (ππ) + 2 π»+ (ππ) → πΆπ2 π72− (ππ) + π»2 π(π) yellow orange A. Addition of Na2 CrO 4 to BaCl2 : • Observations: o Establishing Equilibrium: The yellow sodium chromate and the colorless barium chloride solutions mixed to result in a yellow solution with a small amount of yellow precipitate. o As more sodium chromate solution was added to barium chloride solution, more yellow precipitate was formed in a still yellow solution. • Interpretation (direction of equilibrium shift): o Since the product in the main equilibrium (ME) is a precipitate, the ME must have shifted right. • LCP Explanation: o LCP would predict this because the concentration of chromate ion was increased due to added reagent, so the equilibrium shifted right to decrease the chromate ion molarity. B. Addition of 6 M HCl to the above mixture: • Observations: o As the colorless HCl was added to the above (A) mixture, the solid dissolved and the solution turned from yellow to orange. • Interpretation: o Because the solid dissolved and something orange formed, the main equilibrium must have shifted left. • LCP Explanation: o As HCl was added, the [H +] reacted with the CrO 4 2-, so its concentration decreased. This decrease makes the main reaction equilibrium shift to the left to try to increase the [CrO 4 2-], and a left shift results in less BaCrO 4 precipitate. Note that in the above: • In part A, there was no need to discuss the useful reaction equation as it was not useful in this case. Only use the “useful equation” if you need it. • Reactants are always described. • The equilibrium shift of the main equilibrium (ME) is explained with LCP. • LCP and equilibrium shifts are connected with observations • LCP explanations are often worded as “The equilibrium shifted (fill in) in order to (increase, decrease) the concentration of (fill in). 43 Prelab Name: ______________________________ Review section 15.5 of your lecture textbook which introduces Le Chatelier’s principle. You should be able to predict what will happen as a result of stress applied to a reaction at equilibrium. _____ 1. For the following reaction at equilibrium, which choice gives a change that will shift the position of equilibrium to favor formation of more products? 2 NOBr(g) ↔ 2 NO(g) + Br2 (g) A) B) C) D) E) ΔHºrxn = 30 kJ/mol Increase the total pressure by decreasing the volume Add more NO Remove Br2 Lower the temperature Remove NOBr selectively _____ 2. For the following reaction at equilibrium in a reaction vessel, which one of these changes would cause the I 2 concentration to increase? 2 NOI(g) ↔ 2 NO(g) + I 2 (g) A) B) C) D) E) ΔHºrxn = 45.3 kJ/mol Add more NOI Add more NO Increase the pressure Compress the gas mixture into a smaller volume Decrease the temperature _____ 3. The following reaction is at equilibrium in a sealed container. N2 (g) + 3 H 2 (g) ↔ 2 NH 3 (g) ΔH°rxn < 0 Which, if any, of the following actions will increase the value of the equilibrium constant, K c? A) B) C) D) E) Adding more NH 3 Adding more N 2 Increasing the pressure Lowering the temperature Adding a catalyst _____ 4. The reaction system POBr3 (g) ↔ POBr(g) + Br2 (g) is at equilibrium. Which of the following statements describes the behavior of the system if POBr is added to the container? A) POBr will be consumed in order to establish a new equilibrium. B) The partial pressures of POBr3 and POBr will remain steady while the partial pressure of bromine increases. C) The partial pressure of bromine will increase while the partial pressure of POBr decreases. D) The partial pressure of bromine remains steady while the partial pressures of POBr 3 and POBr increase. E) The forward reaction will proceed to establish equilibrium. 44 _____ 5. At 450°C, tert-butyl alcohol decomposes into water and isobutene. (CH 3 )3 COH(g) ↔ (CH 3 )2 CCH2 (g) + H 2 O(g) A reaction vessel contains these compounds at equilibrium. What will happen if the volume of the container is reduced by 50% at constant temperature? A) B) C) D) E) The forward reaction will proceed in order to reestablish equilibrium. The reverse reaction will proceed in order to reestablish equilibrium. No change occurs. The equilibrium constant will increase. The equilibrium constant will decrease. _____ 6. When the substances in the equation below are at equilibrium, at pressure, P, and temperature, T, the equilibrium can be shifted to favor the products by CuO(s) + H 2 (g) ↔ H 2 O(g) + Cu(s) ΔHºrxn = –2.0 kJ/mol A) B) C) D) E) increasing the pressure by means of a moving piston at constant T. increasing the pressure by adding an inert gas such as nitrogen. decreasing the temperature. allowing some gases to escape at constant P and T. adding a more copper. _____ 7. The reaction 2 H 2 O2 (g) ↔ 2 H 2 O(g) + O 2 (g) is exothermic, ΔHºrxn = –210 kJ/mol. Which one of the following is correct? A) B) C) D) E) K P at 800 K is smaller than K P at 1200 K. Temperature does not affect K P K P depends only on the pressure. K P at 1200 K is smaller than K P at 800 K. K P depends on total pressure as well as temperature. 45 Name: ______________________________ Procedure You will establish an equilibrium for three reactions, make changes to disturb the equilibria, observe the impact of the disturbances, deduce what shift in the equilibrium occurred, and explain how LCP explains the observations. The specific steps you need to follow and the questions you are expected to answer for each reaction are outlined in the report form below. Fill out the included report form as you perform the lab; it is due at the end of the lab session. NaOH, HCl, HNO 3 and NH 3 ARE CORROSIVE. THEY MUST BE WASHED OFF IMMEDIATLEY WITH LARGE QUANTITIES OF WATER. THEY ARE VERY DAMAGING TO EYES! KEEP YOUR SAFETY EYEWARE ON AT ALL TIMES. Part 1 Bromothymol blue is an acid base indicator. Acid-base indicators are large organic molecules that can gain or lose hydrogen ions to form substances that have different colors. The protonated form is abbreviated below as HIn and the deprotonated form is abbreviated as Inβ». Main chemical equilibrium equation: HIn(aq) ↔ Inβ»(aq) + H 3 O +(aq) yellow blue Sometimes useful reaction equation: OHβ»(aq) + H 3 O +(aq) ↔ 2 H2 O(l) 1. Establish an equilibrium by adding 4-6 drops of bromothymol blue to ~3 mL of pH 7 buffer solution in a small test tube. a. Observations: 46 2. Now add 1 M HCl dropwise, until you see a distinct change, to the tube from step 1, to disturb the equilibrium. a. Observations: b. Which direction did the ME shift? What intensive variable in ME changed due to what you did? Did it increase or decrease? c. LCP Explanation of how the ME shifted due to change in intensive variable in ME: 3. Add 1 M NaOH dropwise until you see a distinct change. a. Observations: b. Which direction did the ME shift? What intensive variable in ME changed that led to the ME shift? Did it increase or decrease? How did the change that you made lead to the change in intensive variable in ME? c. LCP explanation of how the ME shifted due to change in intensive variable in ME 47 Part 2 Main chemical equilibrium equation: Co2+(aq) + 4 Cl−(aq) ↔ CoCl4 2-(aq) pale pink bright blue 1. Establish an equilibrium by adding 12 M HCl, dropwise, to 1 mL of 1 M CoCl2 until the solution appears purple. Be patient! a. Observations: 2. Now visualize cooling the sample and answer the following: a. If the οH rxn is positive I can treat heat as a ______________ and the equilibrium would shift ___________ to make more heat, therefore, if I cool the sample I would see: b. If the οH rxn is negative I can treat heat as a ______________ and the equilibrium would shift ___________ to make more heat, therefore, if I cool the sample I would see: 3. Now place the tube in an ice bath and allow it to stand until no further change occurs. a. Observations: οHrxn is ________________. I can tell because: Carefully and thoroughly explain how you can use the color changes to decide whether οHrxn is positive or negative. (Hint: what does the equilibrium shift tell you about whether heat is a reactant or product?) 48 Part 3 Main chemical equilibrium equation: Mg2+(aq) + 2 OH −(aq) ↔ Mg(OH)2 (s) Useful Reaction Equation 1: NH 3 (aq) + H 2 O(l) ↔ NH4 +(aq) + OH −(aq) Useful Reaction Equation 2: H 3 O+(aq) + OH −(aq) ↔ 2 H 2 O(l) Note: You will not need any more than one of the useful reactions at a time. 1. Establish an equilibrium by adding 1 drop of 1 M NH3 to 1 mL of 0.1 M MgSO 4 in a test tube. Observe carefully. If a very light precipitate does not form then add another drop of NH 3 and continue until a precipitate is seen (it may be only a little). 2. Add 1 M NH 4 Cl dropwise with shaking until no further change occurs. a. Observations: b. Which direction did the ME shift? What intensive variable in ME changed that led to the ME shift? Did it increase or decrease? How did the change that you made lead to the change in intensive variable in ME? c. LCP Explanation of how the ME shifted due to change in intensive variable in ME: 49 3. Add 6 M NaOH dropwise with shaking until no further change occurs. (Do not add more than 6 drops) a. Observations: b. Which direction did the ME shift? What intensive variable in ME changed that led to the ME shift? Did it increase or decrease? c. LCP Explanation of how the ME shifted due to change in intensive variable in ME: 4. Add 6 M HNO 3 dropwise with shaking until no further change occurs. a. Observations: b. Which direction did the ME shift? What intensive variable in ME changed that led to the ME shift? Did it increase or decrease? How did the change that you made lead to the change in intensive variable in ME? c. LCP Explanation of how the ME shifted due to change in intensive variable in ME: 50 Experiment 6: Chemiluminescence – Oxidation of Luminol1 Lab Practical Introduction Chemiluminescence is the production of visible light by a chemical reaction without production of heat. It occurs when a reaction produces a molecule in an electronically excited state and that molecule then releases a photon of visible light as it returns to the ground state. Commercially available lightsticks are an example of chemiluminescence. A solution containing luminol is sometimes used in combination with a solution of hydrogen peroxide by crime scene investigators to detect the presence of hidden blood spatters. In this lab you will prepare a solution containing luminol and mix it with a hydrogen peroxide solution and observe the resulting chemiluminescence. Prelab Completion of the prelab is your “ticket” to get into lab. You will not be allowed to work without having these answered, and if you have to answer them at the beginning of the lab period, you will not be allowed extra time on the lab itself. Your instructor may choose to give you a quiz over the proper use of quantitative glassware and/or calculations similar to those needed to prepare solutions A and B at the start of lab. 1. Show the calculation of how much ammonium carbonate monohydrate, (NH 4 )2 CO3 βH2 O, is needed to make 100.0 mL of 2.5 M ammonium carbonate, (NH 4 )2 CO3 . Be sure to write your work clear including units and the correct number of significant figures. 2. Show the calculation of how many milliliters of 3.0% H 2 O 2 is needed to make 500.0 mL of 1.0% H2O2. Note: These calculations are similar to those you will need to do in lab; in lab you will need to do the calculations for how to make solutions A and B without notes or other assistance. You can practice in advance and ask you instructor to check your work during office hours. It would be a good idea to review how to make solutions from solids and how to dilute solutions, including the calculations associated with these processes, before coming to lab (see section 4.5 in your lecture textbook). 51 Hazards All chemicals (including solids) used in this experiment are harmful if taken internally. A solution of hydrogen peroxide is a topical antiseptic and cleaning agent. Rinse it off if you get it on you. Solution Preparation Hints One mole of a hydrate provides one mole of the contained salt, but also brings along some water molecules. Thus, when determining the amount of hydrate needed, the mass of the contained water molecules must be included in the molar mass calculation. As you may recall from the oscillating reaction in CHEM 111, when you are preparing a solution with more than one solute you simply calculate how many grams of each solute you need to make the desired volume (in this case 250.0 mL) of the given concentration for each solute as though it was the only one present, then put all the solutes together before adding solvent. You should dissolve the solutes in a beaker first (just be sure you don’t exceed 250 mL). This will allow you to more easily stir the solution to get the solutes to dissolve and also not clog a funnel or the neck of the flask with solid. After they dissolve, transfer the solution to the flask, rinse the beaker and pour it into the flask and dilute to the line. Procedure At the start of lab you will be directed to either calculate the amount of each solute necessary to prepare solutions A and B (see solution information below) or to take a quiz with similar calculations. After submitting your calculations or quiz prepare solutions A and B. Once you have solutions A and B notify your instructor you are ready to carry out the reaction. The apparatus you need to mix the solutions will already be set up for you. This will be done with the room lights dimmed to better see the chemiluminescence. Slowly pour solutions A and B simultaneously into the funnel and carefully observe what happens. Afterwards, flush the tygon tubing and sink with lots of water. Solution A: Make a 250.0 mL solution that is: • 0.0377 M sodium carbonate • 1.13×10-3 M luminol, C8 H7 O2 N 3 • 0.286 M sodium hydrogen carbonate • 4.38 × 10-3 M ammonium carbonate o Using ammonium carbonate monohydrate, (NH 4 )2 CO3 βH2 O • 1.60 × 10-3 M copper sulfate o Using copper sulfate pentahydrate, CuSO 4 β5 H 2 O Solution B: Make 250.0 mL of a 0.15% H 2 O2 solution using the provided 1.875 % solution of H 2 O2 . Report Turn in anything requested by your instructor before you leave lab; this may include your calculations if you did not turn them before making the solutions as well as observations. 1 B.Z. Shakhashiri Chemical Demonstrations: A Handbook for Teachers of Chemistry Vol 1. University of Wisconsin Press: Madison, WI. 52 Experiment 7: Spectrophotometric Determination of Copper in a Penny Objective The purpose of this experiment is to review how to prepare solutions and how to obtain and use absorption spectra of solutions. You will also review how to use a volumetric flask and a volumetric pipet to prepare a standard solution. You will use a spectrometer to obtain absorption spectra of your standard solution and an unknown solution. This experiment will prepare you for the upcoming Luminol Lab Practical where you will need to use volumetric flasks and volumetric pipets to makes solutions. Background One of the most direct ways to measure concentration, especially for large numbers of samples, is spectrophotometry. A basic spectrophotometer is shown in schematic diagram below. The light path is shown by the long thick gray arrow. The white light from a normal light bulb passes into the monochromator which discards all the light except that for a single wavelength which passes through it. This sounds like a light filter and that is just how a monochromator works. Indeed, glass optical filters can be used for this purpose if a fairly wide range of wavelengths can be used. This light then passes through the sample cell and hits the surface of a detector. The detector’s electrical signal then goes to the electronics (possibly including a computer) where it is converted to an absorbance. We can summarize all this by saying that monochromatic light is passed through the sample and a reading related to the amount of light absorbed is given. The absorbance readings we use are calculated as: absorbance = –log(I sample/I reference). The number is larger as the sample absorbs more light. The reason we use absorbance to represent the amount of light absorbed by the sample is that absorbance should be proportional to the concentration of the species absorbing light, that is: πππ πππππππ = π × πππππππ‘πππ‘πππ where k is a constant that depends on the nature of the species absorbing light, the wavelength used, and the thickness of the sample. The size of the constant, k, is determined experimentally by making a number of standards, which are solutions that contain different known concentrations of the species that absorbs light. We then measure the absorbance of these solutions and make a graph of absorbance vs. concentration which is called a calibration line. The unknown is then run and the calibration line can be used to find the unknown concentration from its absorbance. When doing spectrophotometry, we normally work at a wavelength where the species has maximum absorbance (referred to as the π πππ₯ ), that is, where the species absorbs the most light. The π πππ₯ is 53 determined by making a solution of the colored species and measuring its absorbance as a function of wavelength of visible light. There are spectrophotometers that do this; usually these are diode array instruments and while they work in a way that is similar to the simple spectrophotometer shown above, they produce a whole spectrum (absorbance at all wavelengths) instead of a single reading of absorbance at a single wavelength. Shown below are the spectra of 3 different solutions of the same dye, superimposed on each other. Two points should be made about these spectra. The main point is that no matter what the concentration is, the wavelength of the absorption maximum and general shape of the spectra are the same, just with a different absorbance. This is what is expected if the only change is the amount of the colored material. No matter how much variation of concentration there is, the shape of the spectrum remains the same with just changes in height. Second, we can look at the spectrum and see that the absorbance maximum is at 631 nm and the absorbance is around 1.6 for a 10 mg/L standard. Also note that the absorbance is around 1/2 this value in the 5 mg/L standard and 1/4 this value for the 2.5 mg/L standard. This is enough information to let us know that absorbance measurements for solutions of this substance should be made at a wavelength of 631 nm. We should also be aware that a maximum absorbance of 0.8 is to be used for optimal accurate measurements, so standards should be made that cover the range 0 – 5 mg/L. Once standards are made and absorbances are measured, the concentration and absorbance data are tabulated and plotted. Below is a table of example data along with the resulting calibration line. The graph itself could then be used to determine the concentration of an unknown solution from its absorbance, but it is more accurate to use the calibration line equation to determine the unknown's concentration by plugging in its absorbance in for y and solving for x which is the concentration. 54 Prelab Name: ______________________________ You must bring a penny in 1983 or later to lab. Go to youtube.com to watch the videos on reading the meniscus, the use of a pipet, volumetric flask. After watching Youtube, complete the following questions. These will be due at the beginning of lab. • Reading the Meniscus: https://www.youtube.com/watch?v=QCzaPy_XOqg • Use of a volumetric flask to make a solution: https://www.youtube.com/watch?v=Btp1N5Z2L74 o Notes: βͺ He has already rinsed the pipet well. Invert it more times than he did. • Use of a pipet: https://www.youtube.com/watch?v=TffTiRw8cQY • Review how to do dilution calculations in your textbook (pp.158-9) Using a Volumetric Flask to Make a Solution 1. George needs to make 250.0 mL of an aqueous solution. He weighs the appropriate amount of solute out and then adds 250.0 mL of water to it, stirs to dissolve the solid and mixes it well. What did he do wrong? 2. What should you do with the volumetric flask before placing solute or solute solution into it? 3. Which of the following shows the proper liquid level in a full volumetric flask that is being used for quantitative purposes? 4. a) At what points should you mix the solute and solvent by shaking the flask well? b) Why should you do this both before the liquid is filled to the line and then again after? c) Why should the flask be inverted during some of this mixing? 5. What should you do with the flask when you finish with it? 55 Using a Pipet 1. Briefly describe the procedure used to rinse the pipet properly. 2. Where should the tip of the pipet be compared to the liquid surface … (a) while liquid is being drawn into the pipet? (b) while liquid is being let down to the line on the pipet stem? 3. Which of the following shows the proper adjustment of liquid level in a pipet so it will deliver its stated volume? 4. What do you do about any droplets or partial droplets of liquid on the outside of the pipet? 5. Should any liquid should be left in the tip of the pipet after you have used it to transfer the labeled volume of liquid or should you blow it out? 6. Where should you look to find the precision of the pipet (so you know how many decimal places to record in the volume)? 7. What should be done with the pipet before it is returned to the pipet storage container? You will need a USB drive to save you data when you are in lab, so be sure to bring one with you to lab. 56 Procedure Each student (or pair) will dissolve a penny in dilute nitric acid, and then boil this to remove some of the excess nitric acid. The solution will then be made basic and diluted to some definite volume in a volumetric flask. Each student (or pair) will then measure the absorbance of this penny sample. The analyte is Cu2+ and the brightly colored species is the royal blue complex ion of copper with ammonia formed by the reaction: Cu2+ + 4 NH 3 → Cu(NH 3) 4 2+ pale aqua colorless royal blue . Each student (or pair) in the class will also need to make a standard Cu2+ solution by diluting a lab stock solution (found in a large plastic bottle on the center bench) after adding the standard amount of ammonia to it. The absorbance of the standards will be measured and recorded (for the class data) for the absorbance vs. concentration graph. Then students will be able to determine the amount of copper in their penny from the class absorbance vs. concentration graph that they will construct after measuring the absorbance of their digested penny solution. NITRIC ACID (HNO3 ) AND 15 M AMMONIA ARE CORROSIVE. AVOID SKIN CONTACT; WASH AWAY ANY ACID OR AMMONIA WITH LARGE AMOUNTS OF WATER. 15 M AMMONIA (NH3 ) HAS A VERY STRONG ODOR AND MUST BE KEPT IN THE HOOD AT ALL TIMES. IT MAY NOT BE REMOVED FOR ANY REASON. DO NOT CARRY OPEN CONTAINERS OF 15 M AMMONIA IN THE LAB. USE EVERY REASONABLE PRECAUTION TO PREVENT RELEASING AMMONIA VAPOR IN THE LAB. WASH ANY AFFECTED AREA WITH LARGE QUANTITIES OF WATER. Procedure for preparing penny solution: Weigh and record the mass of a post-1983 penny. Put the penny into a 100-150 mL beaker and add 30-40 mL of 6 M HNO 3. Cover the top with a watch glass, and gently warm the mixture to dissolve (digest) the penny. [CAUTION: IN HOOD]. After the penny is dissolved, boil away some of the excess acid, cool the sample and add the entire solution to a 250 mL volumetric flask. Measure out 50 mL of 15 M NH 3 (the ammonia does not have to be pipeted) and use a medicine dropper and small portions of it to rinse the beaker three times, putting the rinse into the flask. Put the rest of the ammonia into the flask, and rinse the penny beaker with a small portion of DI water and put it into the flask (do not go over the line). Fill the flask to the line with DI water and mix it well. Procedure for preparing standards: Each student will make 100.0 mL of a standard of assigned concentration in the range of about 20–800 mg Cu2+/L by mixing the (pipeted) volume of stock appropriate for their assigned concentration with 25 mL of 15 M ammonia in a 100 mL volumetric flask. The mixture should be shaken vigorously. They should then record the spectrum of the 500 mg Cu 2+/L in NH 3 solution available on a diode array spectrophotometer, noting the wavelength at which the absorbance is highest. This wavelength of highest absorbance depends only on the identity of the copper + ammonia complex present and not on concentration or amount of other colorless materials that are present. 57 Collecting absorption spectra. You will collect absorption spectra for each of your five standards (test tubes 1-4 and the stock solution) as well as an unknown provided by your instructor. Follow the instructions below to collect and save your absorption spectra. 1. First set up your spectrometer using the following steps. a. Plug in the LabQuest 2 device (LQ2). b. Attach the LQ2 to the cube spectrometer c. Power on the LQ2. After an initial startup screen, a screen with a red stripe that shows USB:Abs should appear (this is the home screen). If you don’t see this ask for help. d. At the upper right hand corner there should be a gray box with Mode:Full Spectrum, which is what you need for this experiment. If you don’t see this ask for help. e. Use the stylus attached to the cord (you likely need to pull it out of its storage slot on the back of the spectrometer) to choose Sensors at the top left of the screen. f. Choose Calibrate -> USB: Spectrometer near the bottom of the gray menu. A new screen should now appear, indicating the time necessary for warm up. g. While the spectrometer is warming up, place a clean cuvette filled about 2/3 full with DI water (a blank) into the cell holder of the spectrometer. Note: i. The spectrometer must be placed on a flat surface to hold the cuvette in place. ii. Some of the cuvettes are stained or have other imperfections from prior use, so it is best to use the same cuvette for the entire experiment because this blanking process will account for these things in your spectra. iii. When you put the liquid into a cuvette be sure there are not air bubbles; this can happen when you put the liquid in too quickly. iv. Only touch the cuvette on the frosted sides and wipe the clear sides with a Kimwipe to remove moisture and fingerprints just before inserting it into the spectrophotometer. You do not need a new Kimwipe everytime you do this, and be sure to keep loose Kimwipes out of the hood! v. The clear sides of the cuvette should be oriented so the light beam is travelling through them. h. Once the spectrometer is warmed up and your blank is in place, choose Finish Calibration. i. When the calibration is complete, choose OK to go to the original screen. 2. Collect an absorbance vs. wavelength spectrum for each of your five standards and your unknown using the following steps. a. Empty your cuvette and rinse it with a small amount of the solution you will be measuring. b. Fill the cuvette about 2/3 full with the solution to be measured. c. Place the cuvette into the spectrometer. d. On the LQ2 screen, tap the green triangle at the bottom left. Data acquisition should start and the screen should change to show a full spectrum of your sample. e. When the black line with circles stops changing (this should happen almost instantly), tap the small box at the left bottom with the red square to stop data collection. f. Save this file. i. Select File in the upper left corner and then Save. ii. Touch the field at the top with the stylus and then change ‘untitled’ to your last name and a number (e.g. Miller1, Miller2, etc.; be sure you know what each file name corresponds to). iii. Tap Done and then Save in the bottom right corner. 58 g. Repeat this process for each standard and your unknown. Note: When you start a new data collection it will first ask if you want to discard the old data even if you already saved the data; this will not eliminate your saved data it will just start a new data collection. 3. Follow the steps below to export your data in text format so that you can process it in Excel. a. Unplug the spectrometer from the USB port on the LQ2. b. Plug your USB drive into the USB port on the LQ2. c. Choose File -> Open. d. In the upper left hand corner choose the farthest left icon (this will show you the LQ2 file directory instead of your USB file directory). e. Choose the file you want to save (you have to open the file in order to be able to save it to your USB). This will open the file, but it will not go to the screen that shows the graph (while you could go back to the graph, it is not necessary to export the file). f. Choose File -> Export (NOT Save). g. Type a name for your file (it does not automatically pull the name that you gave it on LQ2). h. Tap Done, then OK. Now the data has been transferred as a text file to your USB drive. i. Repeat the File -> Open, File -> Export process for each of your data files. 4. Either in lab or as soon as you get to a computer, check that you can import your data into Excel. DO NOT wait until the next week to check, otherwise if you cannot open your data then you may not have time to get it so you can complete your lab report. 5. Turn off the LQ2, unplug everything and put it back how you found it. 59 Report You should report your observations of colors and changes which occur during this experiment. Tabulate the class standards data, make a calibration line from this data, and get an equation and R 2 value for the best straight line through the standards data. You may discard any data point you like in making the curve but you must state which points were discarded and why in your final discussion. Use the equation and show how to find the mg Cu in the penny and the % Cu in the penny. Finally conclude how reliable this result is, using your results, and average of % Cu. You should also compare your result with the stated value from the U.S. Department of the Treasury, U.S. mint (usmint.gov) and calculate a % error. Don't hesitate to include a critique of the class data for the calibration line if it looks awful or suspicious to you. Also state which specific data points need to be omitted (if any), and why you would consider omitting them. *Your instructor will let you know of any further expectations. 60 Experiment 8: Synthesis of an Oxalate Hydrate Introduction One of the lab activities of organic, inorganic and pharmaceutical chemists is the synthesis of new compounds. When they start a project, they have a target compound in mind, and after they isolate a product they must find out if they have the desired product. This analysis may be done several ways, but determining the mass percent composition is one step. In this experiment, you will perform the synthesis step of this process, specifically, you will make a hydrate containing potassium, iron(III) and oxalate ions. Your synthesized compound will have the generic formula K x Fey (C2 O4 )zβQH 2O, where x, y, z, and Q are unknown values. In the next experiment you will do a titration to determine the percent of oxalate in your synthesized compound. Procedure • First be aware of the following: o You must rinse all sink surfaces well when pouring solutions containing iron down the drain, as upon standing in air the iron will react to form rust stains. o The final product will decompose upon long exposures to light, so store it in your desk drawer when you are not using it. • Synthesis: o In a flask, mix 30.0 mL of 1.95 M potassium oxalate and 10.0 mL of 1.5 M iron(III) chloride and filter off the resulting solid. Recrystallization (purification): o Add about 5 mL hot water to an Erlenmeyer flask containing the crystals and heat as you stir vigorously. o If the crystals did not all dissolve, add about 0.5-1 mL of hot water and continue to stir vigorously while heating. βͺ Repeat this process until the crystals dissolve. βͺ Caution: If there is a tiny amount of crystals that will not all dissolve do not keep adding water as it is likely to be an impurity and adding too much water will cut your yield! If this happens try to get the impurity out before the next step. o After all the crystals dissolve remove the flask from the heat, cool the solution. βͺ Run water on the outside of the flask, but if it is too hot to hold with your hands you will need to first let it cool in the air before running water over it. βͺ Next put the flask in an ice-water bath or add 10 mL of ethanol to ensure the solution is cooled enough for maximum crystallization. o Filter off the crystals formed. o Spread the crystals on a piece of filter paper and leave them to dry in your lab drawer. • 61 Report You do not need to prepare a lab report for your synthesis reaction, instead you need to complet the following literature search activity: Literature Use: Finding a Cited Journal in the American Chemical Society (ACS) Web Editions The journal format of ACS publications (other than Biochemistry) follow one of the following formats, depending on which journal is used: • Author 1; Author 2; etc. Title of Article. Journal Abbreviation Year, Volume, Inclusive Pagination • Author 1; Author 2; etc. Journal Abbreviation Year, Volume, Inclusive Pagination. For example one citation is: • Caruso, R. A.; Susha, A.; Caruso, F. Multilayered Titania, Silica, and Laponite Nanoparticle Coatings on Polystyrene Colloidal Templates and Resulting Inorganic Hollow Spheres. Chem. Mater. 2001, 13, 400– 409. o For this article, the following information would be needed to look it up using the ACS Web Editions: βͺ Journal Name Abbreviation: Chem. Mater. βͺ Volume: 13 βͺ Page: 400. • If each issue starts with page one, the issue number may be in parentheses after the volume number. You will be assigned one or two of the citations below to look up. Go to MSU library home web page http://www.moreheadstate.edu/library/. Under Other Search Options, choose ACS Web Editions, then enter your MSU ID number and password. At right, toward the top, choose the Citation Tab and enter the volume number and page number for your assigned article. Go to the PDF version of the full article and print page one to turn in next week. Don’t print only the abstract—do full page 1. 1. Organometallics, 2004, 23 (15), pp 3562– 3583 3. Organometallics, 2011, 30 (20), pp 5338– 5343 5. J. Org. Chem., 2012, 77 (18), pp 7804–7814 7. J. Org. Chem., 2007, 72 (2), pp 313–322 9. Anal. Chem., 2013, 85 (24), pp 11677– 11680 11. J. Am. Chem. Soc., 2012, 134 (44), pp 18197–18200 13. J. Agric. Food Chem., 2014, 62 (51), pp 12418–12427 15. J. Phys. Chem. B, 2014, 118 (51), pp 14913– 14921 17. J. Am. Chem. Soc., 2006, 128 (48), pp 15354–15355 19. J. Am. Chem. Soc., 2014, 136 (51), pp 17692–17701 21. J. Med. Chem., 2016, 59 (23), pp 10435– 10450 62 2. J. Nat. Prod., 2013, 76 (9), pp 1541–1547 4. Anal. Chem., 2014, 86 (24), pp 11981– 11985 6. J. Am. Chem. Soc., 2012, 134 (51), pp 20581–20584 8. J. Nat. Prod., 2014, 77 (6), pp 1275–1279 10. J. Agric. Food Chem., 2017, 65 (2), pp 364– 372 12. J. Med. Chem., 2016, 59 (18), pp 8149–8167 14. J. Am. Chem. Soc., 2011, 133 (2), pp 200– 202 16. J. Med. Chem., 2007, 50 (22), pp 5269–5280 18. J. Nat. Prod., 2013, 76 (9), pp 1541–1547 20. J. Med. Chem., 2016, 59 (23), pp 10343– 10382 22. J. Am. Chem. Soc., 2013, 135 (51), pp 19079–19082 Experiment 9: Percent Oxalate Determination by Titration You will do a titration of your oxalate hydrate from Experiment 9 using potassium permanganate in order to determine the percentage of oxalate in your oxalate hydrate sample. The oxalate ion in your hydrate reacts according to the following equation: 5 πΆ2 π42− (ππ) + 2 πππ4− (ππ) + 16 π» + (ππ) → 10 πΆπ2 (π) + 2 ππ2+ (ππ) + 8 π»2 π(π) The slightly pink endpoint is reached when the first drop of excess permanganate is added. Procedure Grind about 1 gram of your product in a mortar and pestle until it forms a powder. Accurately weigh 0.15-0.20 grams of it into a titration flask and add 30 mL of DI water. Add 10 mL of 3 M sulfuric acid and swirl the flask contents until the solid dissolves; if a brownish color develops, add more acid. Use a buret to add about 10 mL of potassium permanganate to the titration flask (be sure you write down your initial buret volume because you will need it to determine the total amount of titrant added once the titration is complete). Add about 1-2 mL of 0.2 M MnCl2 and wait until the color of the KMnO4 disappears (the initial reaction is slow, but the manganese(II) formed catalyzes the reaction). Note: Mn2+ is a product of the reaction, so it is OK to add a little of it, it will not mess up the stoichiometry of the titration. After the color disappears, continue the titration until the faint pink endpoint color persists for 20-30 seconds. At this point all of the oxalate ion has reacted, so record the final buret reading. Repeat this titration two more times for a total of three trials (or repeat additional times for any that you mess up). Use the total volume of titrant per trial to calculate the mass of oxalate in your titrated hydrate sample and then calculate the mass percent of oxalate in your hydrate. 63 Experiment 10: Bleach Titration Lab Practical Objective The purpose of this experiment is to determine the concentration of sodium hypochlorite in an unknown bleach solution. Introduction Two of the most common types of reactions are acid -base reactions and oxidation-reduction (redox) reactions. It is not surprising that such reactions are the basis for many analytical procedures used to determine the amounts of substances in samples. An oxidation/reduction titration involves the addition of an oxidizing agent to a solution of a reducing agent or vice versa. One of these solutions has a known concentration while the other is unknown. The solutions are added until the stoichiometric amounts of oxidizing and reducing agents have been combined. The oxidizing agent in liquid bleach is sodium hypochlorite. To determine its concentration, it will be added to a solution of potassium iodide to produce iodine, which is then titrated with a solution of sodium thiosulfate. Thus, there are two reactions which will take place. First: (1) 2 HCl(aq) + 2 KI(aq) + NaOCl(aq) → I 2 (aq) + NaCl(aq) + H 2 O(l) + 2 KCl(aq) After the iodine is made in reaction (1), it is reduced back to the iodide ion by the sodium thiosulfate as shown in the following equation (2). (2) I 2 (aq) + 2 Na2 S2 O3 (aq) → 2 NaI(aq) + Na2 S4 O6 (aq) Prelab You MUST do the prelab to start lab. Start early and if you do not know how, see your instructor for help. Your instructor may choose to not let you in lab at all if you do not have the prelab done correctly at the start of lab. Have the following worked out carefully with proper unit cancellation shown. These will be checked before you enter the lab! DO NOT use C1 V1 = C2 V2 (or M1 V1 = M2 V2 ) as that is for dilutions, not titrations. Although this equation works in a limited number of titrations, it will not work for many of them. This problem’s numbers are not the numbers you put in your report — it is simply an example so you can be sure you know how to do the calculations in lab. You will have to do the calculations from scratch in the lab without the prelab as a crutch, and you will need to turn in your calculations from the titrations you do in lab before you leave. Part a is a scale-up calculation, and part b is a standard titration calculation. a. Suppose you did an exploratory titration and 2.00 mL of the bleach took 4.60 mL of Na 2 S2 O3 (aq). How many mL of bleach would you use to end up needing 25-40 mL of Na2 S2 O 3 (aq)? Round this to the nearest mL due to pipet sizes available. Show work. Note: There is a range of correct answers, yours just needs to fall in the correct range based on the given criteria. b. What is the molarity of NaOCl in a solution of bleach if 10.00 mL requires 22.35 mL of 0.02191 M Na2 S2 O 3 to reach the endpoint? Show work. Hint: Determine the amount of I 2 needed in the second reaction, then use that to determine the amount of NaOCl in the first reaction. If you don’t remember how to do titration calculations you may want to review section 4.6 from your lecture textbook. Note: The correct answer is 0.02448 M. 64 Safety Liquid bleach is corrosive and must be handled with care. Any spill on your body or in your eyes must be washed with water immediately. Suggestions • • • • You should approach this lab like the acid-base titration lab except for the fact that you should do these titrations slowly after you add the starch. Don’t get in a hurry to get done. The color change after the addition of the starch is quick and if you add titrant too quickly you will miss it. Be sure to add the starch when the solution is a bright yellow like the outside of a lemon, do not wait until it is light yellow or you will likely already be past the endpoint. Note that it is the total amount of titrant per trial that is used to do the calculation — so take an initial reading before beginning the titration, add titrant, add starch, finish the titration and take a final reading. The difference in the two readings is the total volume of titrant used in the titration and to be used in your calculation. Lab Practical Overview You will not be allowed to use any of your own copies of the lab or any additional notes. Your instructor will provide you new copies of the lab instructions and blank paper for your results and calculations. Each of you will be given a vial of unknown bleach solution. This unknown bleach solution must be quantitatively transferred to a 100 mL volumetric flask and diluted to the mark (meaning you need to transfer the entirety of the unknown solution into the volumetric flask; it is important that you don’t leave any behind or lose any of it). The diluted solution in the flask is the solution whose NaOCl concentration you will report, not the original unknown solution. You should begin by carrying out a rapid course titration to determine the amount of diluted unknown bleach solution that will require about 20-40 mL of titrant (thiosulfate) after doing a scale-up. You only do the scale-up once! However, be sure that you are not using more than 30 mL of your diluted bleach solution for each titration, otherwise you will not have enough to do the minimum requirement of 3 fine titrations. You need to create a data table of your results for each titration (including the course titration), show work for your scale-up calculation, and show work for the calculation of the concentration of NaOCl in your bleach unknown from your first fine titration on the provided data sheet that must be turned in at the end of lab. Before each titration you should: A. Have your buret filled with the titrant (sodium thiosulfate, Na2 S2 O3 ) and ready to go with the initial volume reading written down. B. Transfer your needed amount of diluted bleach solution into your Erlenmeyer flask using a pipet. Note: You should only use a small volume for your course titration, then calculate how much you will use for the for the fine titrations (the volume of diluted bleach solution should be large enough that it requires about 20-40 mL of titrant, but not more than 30 mL of your bleach solution otherwise you will not have enough for the minimum requirement of 3 fine titrations). C. Have around 20 mL of DI water measured (it doesn’t need to be exact) and ready to add to the Erlenmeyer flask with your measured diluted bleach solution. D. Have about 1-1.5 g KI measured (also does not need to be exact) and ready to add to the Erlenmeyer flask. E. Have 10 mL of 1 M H 2 SO4 measured and ready to add to the Erlenmeyer flask. 65 Procedure 1. Quantitatively transfer your unknown bleach solution into a 100-mL volumetric flask and dilute to the mark with DI water. 2. Pipet your needed volume of the diluted bleach solution into a 250-mL Erlenmeyer flask. 3. Add about 20 mL of DI water to the Erlenmeyer flask and stir. 4. Add about 1-1.5 g of potassium iodide to the Erlenmeyer flask and stir until it is dissolved. 5. Add 10 mL of 1 M H 2 SO4 ; this addition can be done with a graduated cylinder. The resulting solution should be an orange, dark red, or brown color (this depends on your unknown bleach concentration). 6. Titrate the sodium thiosulfate into the Erlenmeyer flask until the solution turns bright yellow. Be sure you take an initial volume reading of the titrant before beginning the titration! 7. Add a full Pasteur pipet of starch solution to the Erlenmeyer flask. Upon the addition of starch, your solution should turn a gray, blue, or almost black color (this depends on how close you are to the endpoint and your total volume). If the solution is colorless after add the starch, then you overshot the endpoint and will have to start the titration again. 8. Add titrant slowly until the solution just turns colorless. Take the final volume reading of the titrant. *You will have adequate precision when the volumes of titrant used for three trials in a row are within 0.2 mL. *If you violate safety procedures or do not clean up well, your instructor may count off. Report Prepare a table with all of the relevant titration data and results. You should calculate the molarity for each of the three good fine titration trials you did, then average those. Make a neat, well-labeled copy of your scale-up calculation and one sample calculation of your unknown’s molarity. Be sure to write your name on your report and also record your unknown vial number and the volume of diluted bleach solution you used for your fine titrations near your data table. Staple all scratch paper to the back of this. Your instructor will initially grade only your average molarity of NaOCl, but if it is “way off” they might let you redo the calculations for a penalty. 66 Experiment 11: Redox Reactions & Equations For this experiment there is a station set up for each reaction. Instead of working in your typical groups of two you should pair up into groups of four. It does not matter what order you do the stations, but some will be demonstrations that your instructor will do for you first. Carefully write down your own observations (each person) for each reaction; be sure to write down reactant descriptions, immediate changes observed when combined, and changes observed after some time has passed . After your group has completed the reactions you need to go through the process shown in the redox example below to determine the redox reaction for the process. You will need to write and balance your own net ionic equation for each reaction. As shown in the example you should look for half-reactions involving the chemicals you mix to get started on the equation writing (use the table of standard reduction potentials at the end of this experiment). Your instructor will give you hints if you need some. Many of the equations will be relatively simple and not involve that many different species—look for simple equations first! Once you determine the reaction equation answer any additional questions concerning the process (this applies for reaction #1 and #3). You can work with others as you complete these tasks, but ultimately you need to come up with and write out your own answers. You will be penalized if your caught simply copying answers from others; in a previous semester a number of students received zeros on this lab for sharing equations. If discussing a process with other students does not help you come up with the answer ask your instructor, DON’T copy answers from other students. Redox Example The reaction of manganese(II) chloride with sodium bismuthate (NaBiO 3 ) in acidic media. Lab observations: Manganese (II) chloride tetrahydrate is a pale pink solid. When it is mixed in water, a colorless solution results. Addition of the yellow sodium bismuthate with stirring results in a pale purple solution with no solid present. What is the redox reaction for this process? • • What is present? o ππ2+ , πΆπ − , ππ+ , π΅ππ3− Look at the available half reactions (use the table at the end of this experiment) and find half reactions that contain your beginning components. Suppose the following are what we find. o Available half-reactions with ππ2+ βͺ πππ4− (ππ) + 8 π»+ (ππ) + 5 π − → ππ2+ (ππ) + 4 π»2 π(π ) βͺ πππ2 (π ) + 4 π» + (ππ) + 2 π − → ππ2+ (ππ) + 2 π»2 π(π ) βͺ ππ2+ (ππ) + 2 π − → ππ(π ) 1.51 V 1.23 V -1.17 V o Available half-reactions with πΆπ − βͺ πΆπ 2 (π) + 2 π − → 2 πΆπ − (ππ) 1.36 V o Available half-reactions with ππ+ βͺ ππ+ (ππ) + π − → ππ(π ) -2.71 V o Available half-reactions with BiO 3 βͺ π΅ππ3− (ππ) + 2 π − + 6 π»+ (ππ) → π΅π 3+ (ππ) + 3 π»2 π(π ) 2.03 V 67 • Find the best oxidizing agent and best reducting agent, consistent with experimental observations. o If we blindly choose the most positive and most negative voltages, we’d be using βͺ π΅ππ3− (ππ) + 2 π − + 6 π»+ (ππ) → π΅π 3+ (ππ) + 3 π»2 π(π) 2.03 V +( − ) ( ) βͺ ππ ππ + π → ππ π -2.71 V o We need one oxidation and one reduction, so we need to reverse the most negative one: βͺ π΅ππ3− (ππ) + 2 π − + 6 π»+ (ππ) → π΅π 3+ (ππ) + 3 π»2 π(π) 2.03 V + − βͺ ππ(π ) → ππ (ππ) + π 2.71 V o Adding these two together gives Na(s) as a reactant, and we have no Na (s) present, so we can’t use that half-reaction. o The next-most negative voltage is the last listed reaction with manganese, but reversing it would give Mn(s) as a reactant, and that is not present. o The third-most-negative one is βͺ πππ2 (π ) + 4 π» + (ππ) + 2 π − → ππ2+ (ππ) + 2 π»2 π(π ) 1.23 V o Reversing it results in reactants that are present, but it also shows the formation of solid product, which did not occur. o The fourth-most-negative choice would be βͺ πππ4− (ππ) + 8 π»+ (ππ) + 5 π − → ππ2+ (ππ) + 4 π»2 π(π ) 1.51 V o Reversing it gives βͺ ππ2+ (ππ) + 4 π»2 π(π ) → πππ4− (ππ) + 8 π»+ (ππ) + 5 π − -1.51 V 2+ o ππ and water are reactants, (and those are present) and it gives only soluble products, which matches our observations in lab. o Thus, if we take the best voltages consistent with what is present and what is observed, we would use: βͺ Oxidation: • 4 π»2 π(π ) + ππ2+ (ππ) → πππ4− (ππ) + 8 π»+ (ππ) + 5 π − βͺ Reduction: • 2 π − + 6 π» + (ππ) + πππ΅ππ3 (π ) → π΅π 3+ (ππ) + 3 π»2 π(π ) + ππ + (ππ) o The reaction is balanced by scaling the two half-cell reactions to involve the same number of electrons (multiplying the oxidation reaction by the number of electrons in the reduction step and vice versa): βͺ 8 π»2 π(π ) + 2 ππ2+ (ππ) → 2 πππ4− (ππ) + 16 π» + (ππ) + 10 π − βͺ 10 π − + 30 π»+ (ππ) + 5 πππ΅ππ3 (π ) → 5 π΅π 3+ (ππ) + 15 π»2 π(π ) + 5 ππ+ (ππ) o Adding these two reactions eliminates the electrons terms and yields the balanced reaction: βͺ 14 π»+ (ππ) + 2 ππ2+ (ππ) + 5 πππ΅ππ3 (π ) → 7 π»2 π(π ) + 2 πππ4− (ππ) + 5 π΅π 3+ (ππ) + 5 ππ+ (ππ) 68 Name: ______________________________ Procedure 1. a) Penny plus HCl b) Penny plus nitric acid Note: You need to prepare the HCl part early in the lab period (at least within the first hour) and leave it until close to the end of the lab period to take your final observations (it will not be done reacting, but it will be far enough to understand what is happening). Bring at least two pennies to lab; they need to be post 1983. Put your name on a label and place it on a small beaker. Use a triangular file to scrape the edge of a post 1983 penny (it seems to work best to scrape in multiple places on opposite edges of the penny) or cut the penny in half if you wish (and then use both halves). Part a. Pour about 50 mL of 6 M hydrochloric acid into a small beaker. Gently slide the penny down the side of the beaker into the hydrochloric acid; don’t have your face close when you do this, in case it splashes! Observe for a few minutes, then place the beaker at the back of a hood and put a watch glass on top of it to keep dust out. Make sure there is a sign on this hood that says “Keep heat and flames away from this hood”. Come back and take your final observations before the end of the lab period. Part b. Repeat part a but with HNO 3 instead of HCl. This one will react quicker than the HCl so you don’t have to wait as long to take your final observations. Disposal: Pour the liquids down the drain and flush with lots of water. Do not lose what remains of the penny in the drain! It is easiest to pour the beaker contents through a funnel that will catch what remains of the penny and then rinse it so you don’t get remaining acid on your hands. The remains of the penny can be disposed of in the trash. Question: What caused the difference in product in (a) and (b)—consider Eo ’s of reactions in your answer. What does the difference in (a) and (b) tell you about the strength of the HCl and HNO 3 as oxidizing agents? Why did I specify a post 1983 penny? What would happen you used a pre-1980 penny in the HCl? In the HNO 3 ? 2. Al and CuCl2 Place about a pea size amount of CuCl2 in a vial and add just enough water to dissolve it (stir with stirring rod to get solid to dissolve). Lightly crumple a piece of Al foil or powdered Al (whichever is available) so it will go down into the vial. Push the foil down into the solution of CuCl2 with the stirring rod and observe carefully for a few minutes, then check it every few minutes until it quits changing appearance. Note: A gas is sometimes evolved when these are mixed. This is a side-reaction and can be ignored. Keep this equation simple. Disposal: Remove any solid left (pour through filter paper if necessary), rinse it with water and put it in the trash can. The solution leftover can go down the sink (run the water several minutes afterward). 69 3. “Elephant toothpaste” CAUTION--30% HYDROGEN PEROXIDE BURNS SKIN QUICKLY! DON’T GET IT ON YOU. WASH UP ALL SPILLS IMMEDIATELY (THIS INCLUDES ANY THAT DRIPS DOWN THE SIDES OF THE BOTTLE.) While wearing gloves, pour about 5 mL of 30% hydrogen peroxide into a 50 mL graduated cylinder. Observe it carefully. Bubbles indicate the H 2 O 2 is decomposing. Add 2 mL of dish detergent so any bubbles forming will result in foam; this is just so the bubbles will be easier to see, and otherwise does not influence the reaction. After observing again, place the graduated cylinder into the sink and add a pea size amount of KI, which is a catalyst. A catalyst speeds up a reaction but is not itself consumed; you may see a brown/orange color from the catalyst—you don’t need to explain this with your reaction equation. Hint for the reaction equation: This is a decomposition reaction, so the only reactant in the final equation is the hydrogen peroxide. However, each half-reaction can have other stuff that could cancel. Look for one half-reaction in which hydrogen peroxide is oxidized and another in which it is reduced. Disposal: When the reaction is over, be sure to clean up glassware (inside and outside) while wearing gloves. Make sure to rinse the outside of the gloves well and leave them for the next person. All products from this reaction can be rinsed down the sink. Run water for a few minutes after putting them in the sink. Questions: If the hydrogen peroxide is decomposing, what might be forming? Come up with several possible reactions that would explain your observations and use half-cell potentials to determine which is most likely. Note that you don’t need to explain the catalysis. 4. Zn in copper(II) sulfate Add a small piece of zinc to a solution of copper(II) sulfate in a vial. Don’t use much of the solution as you won’t need it to see the reaction. Check it every five minutes. Keep this equation simple. Disposal: Use tongs or forceps to fish the wire out of the vial, holding it over the sink while you do so. Rinse the wire off with water from a squeeze water bottle. Place the solution from the vial in heavy metal waste container. Rinse the sink very thoroughly, even if you don’t think you got anything in it. 5. Sodium and water (demo). Your instructor will do this as a demonstration for you. He/she will first place a small piece of sodium metal into the water. After that reaction has stopped, he/she will add some phenolphthalein indicator to help you figure out the products of the reaction. This equation is not too bad when you take all the evidence into account. 6. Iron(II) Ion and Permanganate Ion Dissolve a small (~ ¼ of a pea) sized amount of iron(II) ammonium sulfate in a minimum amount of water. Add 2-3 drops of the provided sulfuric acid and then add the provided KMnO 4 solution dropwise. 70 7. Thermite (demo) Your instructor will do this as a demonstration for you. You should observe the Al and Fe 2O3 reactants before they are mixed as well as what occurs during the reaction and what the products look like. One product will be molten elemental iron. There is no gaseous product. This demonstration will be done outside, so you will all have to do it at the same time. Keep this equation simple, but remember, solids do not form of just one charge of ion. Some Standard Reduction Potentials at 25 oC* Half-Reaction E° (V) K+(aq) + e- -----> K(s) -2.93 Na+(aq) -2.71 + e- -----> Na(s) Mg2+(aq) + 2 e- -----> Mg(s) -2.37 Al3+(aq) -1.66 +3 e- -----> Al(s) Mn2+(aq) + 2 e- -----> Mn(s) -1.18 2 H 2 O + 2 e- -----> H2(g) + 2 OH -(aq) Zn2+(aq) + 2 e- -----> Zn(s) -0.83 -0.76 Fe2+(aq) + 2 e- -----> Fe(s) -0.44 Ni2+(aq) e- + 2 -----> Ni(s) Sn2+(aq) + 2 e- -----> Sn(s) -0.44 -0.23 Pb2+(aq) + 2 e- -----> Pb(s) -0.13 2 H +(aq) + 2 e- -----> H2(g) 0.00 Sn4+(aq) +0.13 +2 e- -----> Sn2+(aq) Cu2+(aq) + e- -----> Cu+(aq) +0.13 SO 4 2-(aq) + 4 H+(aq) + 2 e- -----> SO 2(g) + 2 H2 O +0.20 Cu2+(aq) + 2 -----> Cu(s) I 2(s) + 2 e- -----> 2 I -(aq) +0.34 +0.53 MnO4 -(aq) + 2 H 2 O + 3 e- -----> MnO2(s) + 4 OH -(aq) +0.59 O2(g) + 2 H +(aq) + 2 e- -----> H2 O2(aq) Fe3+(aq) + e- -----> Fe2+(aq) +0.68 +0.77 NO 3 -(aq) + 4 H+(aq) + 3 e- -----> NO (g) + 2 H2 O +0.96 Br2(l) + 2 e- -----> 2 Br-(aq) +1.07 H +(aq) +1.23 O2(g) + 4 e- +4 e- -----> 2 H2 O MnO2(s) + 4 H+(aq) + 2 e- -----> Mn2+(aq) + 2 H 2 O Cl2(g) + 2 e- -----> 2 Cl- +1.23 +1.36 (aq) MnO4 -(aq) + 8 H +(aq) + 5 e- -----> Mn2+(aq) + 4 H2 O H2 O2(aq) + 2 H+(aq) + 2 e- -----> 2 H2 O +1.51 +1.77 O3(g) + 2 H +(aq) + 2 e- -----> O2(g) + H2 O +2.07 71 Exam and Checkout In the first hour of the last lab period, you will take an exam which is covers material from both CHEM 111 and CHEM 112. This is a standardized ACS exam that will count as one lab grade. Additionally, it will give you practice for the final exam in lecture which will also be a standardized ACS exam, but only over CHEM 112 material. In the second part of the lab period, you will check and clean the equipment in your drawer and do an additional lab cleaning job assigned by your instructor. Failure to attend Lab Check Out and Exam will result in the failure of both 112 Lab and Lecture. 72
