Chemistry 227, 228 and 229General Chemistry Labs 2015-2016 Academic Year Lab Coordinators: Dr. Eric Sheagley
[email protected] Dr. Dean Atkinson
[email protected] 1 Table of Contents General Chemistry Laboratory ....................................................................................................... 3 Grading Criteria .............................................................................................................................. 4 Laboratory Safety Rules and Procedures ........................................................................................ 5 Keeping a Lab Notebook ................................................................................................................ 9 Report Guidelines ......................................................................................................................... 14 General Chemistry Lab Report Checklist ..................................................................................... 19 CH227 LABS ................................................................................................................................. 21 Lab: Scientific Measurements: Precision and Accuracy .............................................................. 21 Pre-Lab: Who has the same solid that I have? ............................................................................. 26 Pre-Lab: How much sugar is in a can of coke? ........................................................................... 33 Pre-Lab: A Cycle of Copper Reactions ....................................................................................... 43 Pre-lab: Which Alkali Metal Carbonate? ..................................................................................... 50 Pre-lab: Using Conductivity to Find an Equivalence Point ......................................................... 57 Prelab: Atomic Emission Spectra ................................................................................................ 65 Pre-Lab: Determining the Concentration of a Solution: Beer’s Law ........................................ 72 CH228 LABS ................................................................................................................................. 84 Pre-Lab: Enthalpy of Neutralization of Phosphoric Acid ............................................................ 84 Pre-Lab: Hess’s Law .................................................................................................................... 91 Deriving the Gas Laws Using Computer Simulations .................................................................. 97 Pre-Lab: Decomposition of Hydrogen Peroxide........................................................................ 102 Pre-Lab: Vapor Pressure and Heat of Vaporization .................................................................. 108 Pre-Lab: Using Freezing-Point Depression to Find Molecular Weight..................................... 116 Pre-Lab: The Rate and Order of a Chemical Reaction .............................................................. 123 Pre-Lab: Chemical Equilibrium: Finding a Constant, Kc .......................................................... 130 Le Chatelier’s Principle in a Cobalt Complex ............................................................................ 138 CH229 LABS ............................................................................................................................... 142 Pre-lab: Acid Rain....................................................................................................................... 142 Pre-lab: Acid Ionization Constant, Ka......................................................................................... 149 Pre-lab: Titration of a Diprotic Acid: Identifying an Unknown ................................................. 158 Pre-lab: Buffers ........................................................................................................................... 167 Pre-lab: Determination of the Ksp of Calcium Hydroxide .......................................................... 173 Pre-lab: Thermodynamics of the Solubility of Potassium Nitrate .............................................. 180 Pre-lab: Redox Titration: Analysis of a Commercial Bleach ..................................................... 186 Pre-lab: Synthesis of Acetaminophen ......................................................................................... 191 Pre-lab: Electrochemistry: Galvanic Cells and the Nernst Equation .......................................... 199 2 General Chemistry Laboratory SYLLABUS – 2015-2016 Prelab Exercises: Prelab instructions are includ ed in the lab packet. You should answ er any questions presented and prepare for the w eeks lab before your lab m eeting. Prelabs are d ue at the beginning of the lab period . Materials: You will need chemical splash safety goggles. These are available from the chemistry stockroom (Room 280 SRTC) or at the campus bookstore. You will need a bound carbonless copy notebook (not loose paper) for recording data. You are responsible for all laboratory equipment checked out to you. If you break glassware, you will pay the replacement cost of the glassware. D ress for lab: You m u st w ear shoes that cover your entire foot, includ ing the heel. They should fit up near your ankle; leather is preferred but any non -porous m aterial is okay. Short shorts and short skirts are not allow ed . Your clothing m ust cover your torso and legs d ow n to your knees. You w ill also be required to w ear a provid ed lab coat w hile w orking in the lab. Grading: The laboratory is grad ed on a Pass/ N o Pass basis. An average of 75% of all points available in the lab is required to pass. Late Work: Laboratory reports are d ue at the beginning of the lab period follow ing com pletion of the experim ent. Lab reports should be typed . Late reports w ill be d ocked 5 points per d ay late. Attendance: Attendance in this lab is mandatory. YOU MUST ATTEN D ALL SCH EDULED LABORATORY MEETIN GS. If you are not able to attend lab you m ust notify your laboratory instructor as soon as possible. St udent s are responsible for complet ing t he lab report for t he missed lab. Data can be obtained from a lab partner or the lab TA. The m ad e up w ork should be clearly labeled and ind icate the origin of the d ata reported . Reports are d ue the class m eeting follow ing the syllabus d ead line. In addit ion t o complet ing t he make-up lab y ou must make up t he missed lab t ime. The m ake-up laboratory w ill not be the sam e lab you m issed but w ill be a unique activity that w ill take place d uring w eek 10 of the quarter, d uring the regularly sched uled lab period . FAILURE TO DO BOTH WILL RESULT IN A N O PASS GRADE. If you m iss tw o or m ore labs your grad e w ill be a N O PASS. N OTE: If you are m ore than 15 m inutes late to lab you w ill be m arked late. Tw o late arrivals d uring the term w ill be counted as a m issed lab. In ad d ition, late stud ents m ay be assigned to lab clean up d uties at the conclusion of the lab period . If you are chronically late you w ill be given a N O PASS at the lab coord inators d iscretion. Plagiarism : Experim ents w ill be d one in groups sharing the com pu ter for d ata analysis and acquisition. You m ay com pare d ata w ith other groups, but the content of your lab reports MUST be w ritten individually. It w ill be consid ered an act of plagiarism if you 3 borrow tables or graphs from another stud ent (learning how to properly create a table or graph is an im portant skill, learn how to do it on your ow n!). You cannot paraphrase the internet, your book or any other source w ithout the proper reference. Ad d itionally, it w ill be consid ered an act of plagiarism if you borrow d ata w ithout prior approval from your TA. There are ad d itional resources online to help you avoid plagiarism . Please be sure to check http://www.lib.pdx.edu/instruction/survivalguide/writeandcitemain.htm or http://web.pdx.edu/~b5mg/plagweb.html, and feel free to discuss the issue with your TA or the lab coordinator. Depending on the severity of the offense(s), you will receive, at a minimum, a zero score for the report. Additionally, a report may be made to the Office of Student Affairs. Grading Criteria Unless otherwise noted in the course schedule, every lab report is worth 120 points, including the prelab, notebook and technique. Each lab report will be graded according to the following point distribution: Prelab: 20 points Abstract: 10 points Introduction: 10 points Data: 10 points Results: 15 points Discussion: 15 points In addition to the above points each lab meeting will have an additional 40 points assigned on the following basis: Notebook: 20 points These points are awarded by the TA based upon the quality of your lab notebook. Your TA will be looking to see that you are including a title, a statement of purpose, the procedures, data tables and that all data is present. Lab technique: 20 points The basis for assigning these points includes (but is not limited to) general lab technique and methods, safety, general mannerism in lab and cleanliness. Both of these criteria will be evaluated by your TA during each lab meeting. At the end of each lab you must check out with your TA so that he or she can assess your lab notebook and verify that you have cleaned your work area. Grading: Your grade will be assigned based on the percentage of total points scored in the class approximating the following scale (Note: this scale is subject to change based on class performance): Grade Score A ≥ 90% B ≥ 80% C ≥ 70% D ≥ 55% F < 55% More than one absence will result in a grade of F for the class. Do not copy your partners, friends, old lab reports. That is plagiarism! 4 For this reason it is necessary to wear proper PPE. The most important aspect of having the goggles fit comfortably is the proper adjustment of the strap length. • Clothing – Dress appropriately for laboratory work. leather is preferred but any non-porous material is okay.Laboratory Safety Rules and Procedures Safety Rules The guidelines below are established for your and your classmates’ personal safety. They should fit up near your ankle. You should not wear contact lenses in a chemical laboratory. be certain to remove gloves and wash hands before ingesting food or drink. The PPE for student labs consist of goggles. Some chemicals may dissolve “soft” contact lenses. Proper PPE is required for all students or they will be asked to leave the lab •Goggles – Goggles must be worn whenever any experimental work is being done in the laboratory to protect the eyes against splashes. • Eating. Spills and other accidents can occur when least expected. Chemical vapors may become trapped behind the lenses and cause eye damage. Adjust the strap length so that the goggles fit comfortably securely and are not too tight. • Gloves – Gloves should be worn to protect the hands from chemicals. In addition. this prevents things such as door handles from getting contaminated. 5 . including the heel. If you take a break. • Personal Protective Equipment (PPE) is used to protect you from serious injuries or illnesses resulting from contact with chemical hazards in the laboratory. Your clothing must cover your torso and legs down to your knees. Wash your hands after finishing lab work and refrain from quick trips to the hall to drink or eat during lab. gloves and clothing. These are for sale in the bookstore and stockroom. Failure to adhere to the guidelines below will result in a loss of Lab Technique points. drinking and smoking are prohibited in the laboratory at ALL times. If you find that your goggles tend to fog. Gloves are provided through your student fees and are located in the student labs. Only indirect-vented goggles are allowed in the student labs and should be worn at all times when any chemical is being used in the lab. you can pick-up anti-fog tissue from the stockroom. For health and safety reasons it is important to always remove at least one glove when leaving the student laboratory. • Headphones may not be worn in lab. You must wear shoes that cover your entire foot. you are required to wear a department provided lab coat while working in the lab. • Never work alone in the laboratory or in the absence of the instructor. then notify your lab instructor or stockroom personnel. safety shower. stand under the safety shower and flood the affected area with water. Leave the building by the nearest exit and meet your TA on the field next to Hoffmann Hall. • All chemical waste must be disposed of in proper containers. • If a mercury thermometer is broken. doors that lead to other labs may be the best choice. should be kept in the designated area at the entrance to the lab. roll the person in the fire blanket to extinguish the flames. It is good to know the nearest exits of your lab on the first day of class. • In case of fire or accident. Your book bag. etc. Laboratory Procedures and Protocol General Etiquette: • Leave all equipment and work areas as you would wish to find them.. so that mercury is not spread. call the instructor at once. Putting chemicals into the wrong 6 . • If a person’s clothing catches fire. Proper disposal of chemicals is important for student safety and proper disposal. • Spilled chemicals must be cleaned up immediately. Notify students around you. Remove clothing to minimize contamination with the chemical. This would also be the meeting place in the event of an earthquake or other emergency. • In case of a chemical spill on the body or clothing. If acids or bases are spilled on the floor or bench. first aid kit. then dilute with water. eyewash fountain and all exits. • Keep your lab bench area neat and free of spilled chemicals. ask the instructor for assistance. Most other chemicals can be sponged off with water. • If evacuation of the lab is necessary. not at your bench. leave through any door that is safe. If the material is corrosive or flammable. fire blanket. do not attempt to clean up yourself. The stockroom is equipped for proper clean up and disposal of mercury. • Small fires may be extinguished by wet towels. fire extinguisher. Notify the instructor or stockroom personnel if ANY blood is spilled in the lab so that proper clean up and disposal procedures may be followed.Safety Procedures • Know location of safety equipment. • Avoid contact with blood or bodily fluids. neutralize with sodium bicarbonate. or not obstructed. coat. ii.dirty glassware is harder to clean later. Be sure that the lids or stoppers are replaced. Gloves – this is ONLY for used gloves. glass. There are several locations for very specific waste. • Clean your bench and equipment Clean all your glassware. • Clean the common areas before you leave the lab. Waste jars for each experiment will be provided in the lab. They will be labeled specifying which contents should be placed inside. iv. v. • We recommend always picking up bottles by the label. • Do not put stoppers or lids from reagents down on the lab bench. or gloves. Rinse well with water. Contaminated paper waste – this is ONLY for paper towels used for clean-up of chemical spills. Make sure the chemical name and concentration match what is required by the experiment! • Do not take the reagents to your bench. Do not put anything down the sink unless you are explicitly told to dispose of it this way. They may become contaminated. It is important that you replace the lids to the waste containers. as it is frequently oily. Wash with water and detergent scrubbing with a brush as necessary. Normal trash – this is for all other trash that is not chemically contaminated. Handling Chemicals: Obtaining reagents: • Read the label CAREFULLY. Remember to wear gloves while working with reagents. The Chemicals are organized by experiment in secondary containment bins. The water and gas should be turned off and your equipment drawer locked. Broken glass – this is ONLY for broken glassware. make sure it is placed in a secondary container.containers can lead to injury from unexpected chemical reactions. When done with the waste jar. Point deductions for the entire class will be imposed if the instructor or stockroom is not satisfied. i. Mixing waste can also make it more difficult or expensive for PSU to dispose of them. • Return any special equipment to its proper location or the stockroom. Chemical waste – these containers are ONLY for chemical waste generated in the lab. Do not dry glassware with compressed air. 7 . Your instructor will provide specific disposal guidelines when needed. If all students do this. Only chemicals should go into waste jars. then any unnoticed spills when pouring will not cause possible problems for the next user. iii. READ THE LABELS. They are each specifically labeled for each lab and waste type. Following these guidelines assists us in lowering the environmental impact of the labs. • All chemicals should be treated as potentially hazardous and toxic. Treat it as waste and dispose of it properly.• Do not place your own pipet. acid should be added slowly to water with continuous stirring. Laboratory Procedures • Never pipet any liquid directly by mouth! Use a rubber bulb to draw liquid into the pipet. This process is strongly exothermic. • Use the fume hood for all procedures that involve poisonous or objectionable gases or vapors. explosive spattering. notify your instructor. • To prepare a dilute acid solution from concentrated acid. never weigh directly onto the weighing pan. and adding water to acid may result in a dangerous. You can always go back for more. • Never weigh hot chemicals or equipment. • When weighing chemicals on the balances. always use a test tube holder and be certain never to point the open end of the test tube toward yourself or another person. • When heating a test tube. • Any chemicals that come in contact with your skin should be immediately washed with soap and copious amounts of water. • Do not put any excess reagent back in the reagent jar. If you are unclear how to clean a spill. Please pour on the conservative side to minimize waste and cost of labs. and keep all glass pieces wrapped in a towel while applying gentle pressure with a twisting motion. hold the tubing in your hand close to the hole. Weigh into a weighing boat or beaker. When smelling a chemical. Never taste a chemical or solution. The balances you are using are precision pieces of equipment and costs up to $4000. gently fan the vapors toward your nose. lubricate the glass tubing with a drop of glycerin. Any spills on the balances MUST be cleaned up immediately. or spatulas into the reagent jar. Pour a small amount into a beaker and measure from that. dropper. 8 . • Never use an open flame and flammable liquids at the same time. • Handling glass tubing or thermometers: to insert glass tubing into a rubber stopper. there are certain principles that should be followed. High on this list are the following: Use a notebook with pre-numbered pages Record entries in ink Keep entries reasonably neat and organized Never tear pages out of your lab notebook (other than the carbonless copy pages) What Kind of Notebook Should I Use? For this class you must use a notebook with carbonless copy pages. So. There are also certain conventions for lab notebooks that are universally followed. 9 .Keeping a Lab Notebook In keeping a lab notebook. and working notes • Prepare data tables in advance . These boil down to being clear and complete in your entries in your lab notebook. In order to be clear. (This applies to notebooks in learning laboratories: Your lab instructor may want to look at what you did in order to understand your results. etc.) The word “clear” here is crucial. Descriptions of procedures must be clear and concise. to the point. This boils down to clear descriptions of what you did and what you observed as a result. It is a working tool. Nothing should be recorded on odd scraps of paper. manual calculation. leave a few pages for a Table of Contents • Each lab should have a brief introduction and description of procedure • Generally use only the right hand page for most text • Use facing left page for working graphs. That is the proper place for all lab planning and observations. clearly labeled. You should record all your work in your lab notebook. it needs to be clear. General Guidelines • Write your name on outside front of notebook • Use black ink. and a reference for other researchers who might want to read your notebook and reproduce your work.with columns for calculated results and notes • Working graphs done in lab notebook to monitor progress Usage and Structure The overriding principle for a lab notebook is to record in it all the pertinent information about your lab work. fine-tipped ball-point pen (this will photocopy clearly) • At the front of the notebook. data must be recorded in well-thought-out tables. This is often the case. what conclusions you have reached. not 9. use the left-hand page to do the calculations. you would describe what results you got. Title: With your lab notebook laid open. Your procedures should be of sufficient detail that you. can independently perform the lab activity without looking at the lab manual. etc. This section is where you should have pre-prepared tables for data collection. I subtracted wrong! We put in 10. Results and Discussion: You might want to include a final section that is labeled Results and Discussion. If unexpected results occur later. write down what you actually do and what you observe. One example of how to use the left-hand page: if your work requires simple calculations using your measurements. Purpose: Below the title. record observations and data or measurements. In general.Structure for your Lab Notebooks: For each lab in this class you should have the following sections in your lab notebook: Title Purpose Procedure and Observations It is also often helpful to include a Result section Note: When preparing your notebook for lab only write on the right hand page. ideas for continuing work. and the date. It will be these ideas that you present in your lab reports. such as the molar mass of an unknown sample. (“Oh. As the name suggests. 10 . your purpose will be what you are attempting to find or solve for. The left-hand page is reserved for recording scratch work.5 like we thought!”) Better to discover the error after the fact than never to discover it at all. on the right hand page write down the title of the experiment. In general. you will use the right-hand page for all your writing. In this section. Don’t use this space until you need to. In the left hand column write your procedure and in the right column next to the procedure. Procedure and Observations: This next section will be labeled Procedure and Observations. sometimes you can look back at your scratch work and discover the error.5 grams of copper sulfate. write the purpose of the experiment in one or two sentences. the student. Set up this section by dividing the page into a right and left column. This should be done before leaving the lab. This section serves to remind you and notify the reader what the experiment is about. 11 .An example of a prepared notebook follows. 12 . It is not expected that you write perfect prose in your notebook – it is a first draft.” Adapted courtesy of Keith James and Jonathan Frankel. such as the Introduction before each experiment. Also. you should strive to write as logically and clearly as possible. However. this would be expanded a bit and made grammatically correct.0 M HCl were added to the clear reaction mixture. with many entries being written while an experiment is in progress (your observations) it is understood that many entries will be brief – but still record crucial observations.↓ a few secs later→ clr soln. Just do the best you can. This immediately resulted in a crimson solution. this is a working document. “10 mL of 1. leaving a clear solution. to get into the habit. pcpt. It is also a good idea to write in the third person passive voice.Writing Style in the Lab Notebook For certain entries in your lab notebook.” When written into a lab report or journal article. and so that in many cases you can copy entries from your lab notebook into your reports without the need for major revisions/rewrite. as a working document. Example Notebook entry: “Added 10 mL of 1M HCl – solution turned red instantly. 13 . and a red precipitate formed a few seconds later. This includes referencing your TA. your lab group. 14 .” Also bad: “Lab groups poured 50 mL of hydrochloric acid into a flask” This is not the correct form of 3rd person. in fact.” Uses future tense. This is done with the aim to identify articles that need to be read in full. Then do it again until it is right. you will find instructions on what type of report will be required. You may discuss the calculations and analysis with your lab mates or TA. however. It is expected that you will complete each experiment and do the necessary calculations and analysis during the scheduled lab period each week. groups of students. it isn’t. the abstract needs to be brief. which can be found on D2L. For most experiments this term. Abstract: This is a condensed version of your lab report. It includes Joe’s name. You will enjoy writing reports more if you take pride in what you hand in. The reports are due by the beginning of class the week following completion of the experiment.Report Guidelines At the end of every experiment in this manual. passive voice. Other experiments will require a formal report. also. Good: “50 mL of 1. Abstracts are. It should not contain personal pronouns such as. The sections are presented here in the order they should appear in your lab report. After you write your report. your written lab report should be your own individual work!! The lab report sections should be complete but CONCISE. Also bad: “Joe Shmoe poured 50 mL of hydrochloric acid into a flask. your report should be 2-3 pages long. Fix it. “I”. Below is a description of what should be included in each section. A library search of the literature generally involves reading abstracts. or the class as a whole. It must also be in the past tense. A lab report must be written in the third person. It is a stand-alone document. and eliminate many others whose abstract makes it clear that they are not relevant to the study at hand. So. If is doesn’t sound right. For your work. your lab drawer. Some experiments require only a worksheet. often published separately from the articles they describe. Writing Style You will write your reports using a formal scientific writing style. but complete. there is one more thing to do before you print it and hand it in: Proofread it! Read it out loud. Also bad: “We are going to put 50 mL of acid into the flask. this section should usually be between two and three sentences long.0 M HCl were poured into a 125 mL Erlenmeyer flask” Bad: “I poured 50 mL of hydrochloric acid into a flask. “we” or “he” neither should it contain proper names of persons. “we”.” This is not the correct form of 3rd person. such as gravimetric analysis. you should talk about it in the introduction. The purpose of this experiment was to determine the color of the sky. Remember to keep this section brief. you should consider writing this part of the lab report after you have finished the remaining sections. What did you do? In the introduction.There are three questions that should be answered in any good abstract 1. Introduction: Here. Be sure to include any mathematical equations that are new to this experiment when appropriate. How can this purpose be achieved? Provide any relevant background to put the experiment in context. Doing so will better help you identify the major results and allow you to include the appropriate percent errors when possible. Chemical equations should be handled the same way. These equations should be on their own line. What did you do? 2. Don’t just write the equations. Your Abstract answers the question. you should look back to your purpose in your lab notebook. Your Introduction will often include some explanation of the theory behind the experiment. 15 . stating the type of methods used. provide information as to why they are relevant. What did you find? Even though it sequentially appears first. but part of the paragraph itself as they come up. The rest of your Introduction section should flow from the purpose. This is where you will talk about the key concepts of the experiment. if you use a law in the lab’s procedure to determine something. Be sure to include the units for your numerical results. then the first question would be answered by stating that the molar mass of an unknown compound was experimentally determined. You will write your procedure in your lab notebook. there is no need to rewrite it here. For the second and third questions. you want to address WHY you did this experiment. rather than trying to explain the procedure. Which laws are being used to help you fulfill the purpose? What are those laws? Why are they important to this experiment? As a general rule. To answer the first question. How did you do it? 3. For example. You should use your purpose statement as a guideline for the rest of the introduction. Your introduction should follow a logical pattern from your statement of purpose to your data section without relying on a statement of the procedure. Your introduction begins with a statement of the purpose of the experiment. you want to be brief. your purpose statement is what you will do. You will find that as you write the report that you will be repeating yourself a bit. If your purpose was to determine the molar mass of an unknown solid compound. (In published papers. Figures should not be divided across page boundaries Remove gridlines. including units. If the slope or intercept is necessary for other parts of the experiment. When you present your data in a table it is necessary to take the following into account. If this information is pertinent.980 3. Table 1: Mass and volume measurements when a portion of an unknown solid was dissolved to make 10. Be sure to refer the reader to view the tables in the text.Data: This section is where your experimental data belong. interpretation and discussion of your results belong. it should be included in the caption.042 2. Tables should not be divided across page boundaries For a simple example. Tables should include descriptive column headings. Please prepare graphs using the following guidelines. including units. then place the values in the caption with proper units.0 mL of aqueous solution. Each axis should be clearly labeled. Number tables sequentially as they appear (Table 1. data should be presented in the clearest format possible. be sure to cite the figures in the text. For a simple example.). titles and equations from the graph. Trial 1 2 3 Mass of unknown Sample (g) Mass of solution (g) 3. This section is not where the calculations.).356 Graphs When a table does not provide a clear picture of the data. when appropriate. In your writing. Insert a caption below the graph that briefly explains what the graph is presenting. a data section is usually not included.) Tables Whenever possible.021 12. 16 . Pick the format that best displays the data and stick to that. Number figures sequentially as they appear (Figure 1. this is a class so this section will be included.964 11. usually in the form of a table. Table 2….128 12. but. see Figure 1. Construct a descriptive table caption and place it above the table. see Table 1. a graphical presentation of data is necessary. Figure 2…. In this section you would also include observations and descriptions of other pertinent events. Do not present the exact same information as both a table and a graph. You do not need to write out your calculations here. including your significant results. This DOES NOT mean to include detailed procedures or that you need to re-explain your calculations in words. The Introduction states the purpose.04 0. If something did go wrong.6 0. It DOES mean that a general description of the experiment can be useful in explaining your results and putting them in context. If this is the case. This is where you get to present your thinking process.12 Concentration (mM) Figure 1. You will want to draw everything together in this section. and should always display the significant result of the experiment. for example.8 0.4 0. Results: The results section is where you should report all of your results. It is possible that nothing actually went wrong. A calibration curve for the absorbance at 470 nm of aqueous Allura Red solutions as a function of the concentration.1 0. you will need to dig deeper to explain away any error. this is also where the answers get written up. For any labs that have questions to answer. This section will usually contain either tables or figures.Absorbance at 470 nm 0.06 0. Any calculations that you do should be attached at the end of the report. For your error analysis you should be able to look at whether your results were greater than or less than (in magnitude) the expected results. The Results section gives the final results. In this section you should also discuss error analysis. Anything that you are calculating.08 0. it would not be plausible that the primary error source was loss of heat to the atmosphere. You should be able to identify plausible sources of error from that.86 mM -1.02 0. A best fit line was rendered resulting in a slope of 5. Your Discussion section should bring everything from the results section to your final result along a similar pathway. Try to investigate what might have caused your results to vary. Try to see what could explain away the direction of your error. The Data section shows the collected data. This happens when your percent error is very small. Discussion: In this section.2 0 0 0. should be stated here. The discussion is one of the most important parts of the lab report! It is your chance to show WHAT YOUR RESULTS ARE and that you UNDERSTAND what you did in the lab. If you measure too much heat. like your lab partner 17 . Your introduction provided the purpose and followed a logical pathway to measuring Data. you will discuss interpretations of the experimental results. 18 . then that should go here. Just keep it in third person passive voice When answering any additional questions that are posed at the end of the lab. There are many possible methods that might be used to solve a particular problem or determine a particular characteristic of a chemical. too. along with an explanation of how you attempted to correct for the error. Adapted courtesy of Keith James and Jonathan Frankel.forgot to write down the exact molarity of your reagent. make sure that you approach the question from the context of the lab. Each experiment is designed to show you how a specific problem can be solved using some set of theories and laws. Use that information to answer additional questions. 25 mL was added…. we. _____What did you do? (Identify the rationale behind the investigation)? _____How did you do it (summarize the procedure. Table captions above and figure captions below.25 mL was added)? _____ Did you make sure that you did not start a sentence with a number? _____ All subjects and verbs are in agreement? _____ Did you make sure that there are no run-on sentences or fragments? Abstract The abstract is a condensed summary of the report's findings. Tables and figures are not broken over multiple pages _____Are the axes on your graphs formatted properly with labels? _____Are all graphs and tables accompanied by a written description relating the same information to the reader? 19 . in the context of this lab.25 mL was added…. and self-contained and. not . approximately three sentences long. not 0. without using specific steps)? _____Present the important findings numerically including error statistics? Introduction The introduction will provide the reader information on what you are doing why you did it and critical background information necessary in understanding the methods and results of your experiment. Abstracts are often written last. chemical or mathematical.General Chemistry Lab Report Checklist General _____ Have you listed your name.) _____ Is proper tense is maintained within sections? _____ Have you correctly written your chemical formula and names correctly? _____ Were correct subscripts. a descriptive lab title and date? _____ Did you use spellchecker? _____ Is your report written in passive third person voice (you did not use the words I.25mL was added)? _____ Did you check significant figures? _____ Do your numbers include leading zeros (0. without results _____Are your data tables properly formatted? _____Are your figures and tables numbered sequentially and referred to in the text. included? Data This section should give only the data and observations from the lab. _____Did you include a statement of purpose? _____Is there sufficient background so that the reader can understand what you did? _____Are necessary equations. etc. They should be clear. they. and symbols are used? _____ Did you separate the numbers from their units (0. partner's name. concise. superscripts. you will discuss interpretations of the experimental results. tables or key results? _____Have you discussed sources of error or ambiguities in the data? _____Did you confirm all relationships that were stated in purpose or abstract? _____Do your conclusions clearly contribute to the understanding of the overall problem? _____Are all of your calculations attached at the end? 20 . with an explanation as to how that source caused your results to differ from the accepted or expected values.Results: This section is for any significant results that you calculated. Your readers must easily find your results in order to evaluate and interpret them. _____Units? Significant Figures? _____Is a straight forward presentation of the results of your experiment included in either a table or in text? _____Can your key results be understood by a reader without reliance on figures and tables? Discussion: In this section. You will need to present plausible sources of error. You should have a paragraph here that clearly states your results in addition to the tables or figures. The results here should reflect what you stated in your purpose. You should bring everything together in this section following a logical progression similar to the introduction. cite tables or figures. _____Can your key results and discussion be understood by a reader without reliance on figures and tables? _____Are key results highlighted and carefully explained? _____Did you make logical deductions based on the results (usually questions are given in the lab manual to help this)? _____Have you referenced any figures. It will be necessary to describe your results. while the roadside scale has the ability to measure large masses to the nearest 20 pounds. To arrive at the 155. has only 2 significant figures. think about measuring your weight on a bathroom scale as compared to a roadside scale.14 g/mL. 7. For instance. you had to 21 . you measured the density of zinc to be 7. Because the value 7. reproducibly. the accepted density of zinc is 7. The greater the level of precision. You know your weight greater than 155 pounds but less than 156 pounds. suppose that when weighing yourself in the example above. This shows that there is an uncertainty in the last digit of the measurement. For instance. Precision also deals with how small of a change can accurately be measured. your values are quite close to each other. The value 140 lb. In general. Precision is a measure of the reproducibility of a measurement. the needle points somewhere between 155 and 156 pounds. when making a measurement.6 pounds but on the roadside scale. Because of the scale used above. in a series of trials. The manipulation of measurements with a know precision (number of significant figures) is included in your text book. on the other hand. For example. Just be sure not to report more significant digits than the instrument used for measuring will allow. One student may experimentally determine it to be 7.27 g/mL and another may determine it to be 7. This suggests that your precision is good. 155. Scientists represent this precision in the equipment by putting writing “± error” after the measurement. In those cases. The bathroom scale has the ability to measure relatively small masses to the nearest 0. On a bathroom scale your mass might read 155. The more significant figures that are reported the more precise our measurement. a scale the measures the weight of large trucks. In this class. the greater the number of digits in that measurement that are significant (Signficant Figures). the one and the four.1 lb. For example. Measurements made of relatively small masses will be more precise with the bathroom scale. accuracy cannot be measured.1 pounds. it is customary to include all the known values of a measurement plus one digit that is estimated. we will use significant figures to indicate the level of precision with which a measurement has been made instead of writing the error using the “± error “notation. your mass may be given as 140 pounds.20 g/mL.20 g/mL is closer to the accepted value. Assuming that the bathroom scale is not digital. You will find that in some cases there is no accepted value to compare a result to. you will be able to measure a mass to the nearest 0. The limited precision of the roadside scale would create a level of uncertainty in measuring small masses. it is considered to be a more accurate measurement.CH227 LABS Lab: Scientific Measurements: Precision and Accuracy Background Accuracy is a measure of how close a measurement is to the correct value.6 ± 0.28 g/mL and 7. how close a set of measurements are to one another.1 pounds.6 lb has four significant figures. This uncertainty will restrict the conclusions that can be made from the measurement so a roadside scale would not be the instrument of choice for measuring a person’s weight. the first scale’s measurement would be written as 155. meaning that all of the digits in the value are known with a relatively high degree of certainty. If.26 g/mL. The roadside scale measurement would be written as 140 ± 20 lb.25 g/mL. we can only have a relatively high degree of certainty of the value to the tens place.6 pounds reported above. the graduation is every 1 lb. one only needs to look at the graduations on the instrument used to make a measurement.estimate the 0. report your value to your TA. In the example above. Your TA will lead a quick discussion about your results before allowing you to continue. estimate the value to one decimal place more than the level of graduation.6 pounds (the estimated value in the measurement). Procedures You will be graded on the number of digits used. Your TA will have a large graduated cylinder filled with some amount of a liquid on display. You will be limited to ±1/10 of the smallest known graduation. Circle the best response. Without consulting anyone in your class and without sharing your value. Take out a 25-ml graduated cylinder. Since the graduations are every 0. The 25-mL cylinder is calibrated to every ______ mL.1 lb Every measurement you make in this lab must include the proper number of significant figures. determine the volume of the liquid and record the value here__________ Before continuing. Use a pen to record all measurements!!! 1) Measurements using a graduated cylinder. In general. Therefore. one always records the volume as the level at the TOP / BOTTOM of the meniscus. Check to see whether your graduated cylinder is calibrated to every 1 mL or every 0. the measurement is reported to the 0. Use the following info to estimate the the measured value: 22 . the presence of units on your measurements. and explanations. To determine this. When measuring volume using a graduated cylinder.5 mL it is difficult to report your measured volume beyond one decimal place.5 mL. with at least two labeled graduations (as in the previous drawing).0 mL.7 mL.4 x 0.4) of the distance between the graduations 12. 12.2 mL = 12.5 mL.5 mL + 0. Place about 10-20 ml of water into your 25-mL cylinder and make a sketch of it below.With certainty.5 and 13. 0.7 mL should be recorded.5 mL. you can see the liquid is slightly above 12.5 mL. Be as accurate in your drawing as possible. Let’s approximate that the liquid is 4/10 (0.5 mL = 0. The bottom of the meniscus sits just above 12. Since each graduation is 0. Report your measurement here (Don’t forget units!): ________________ 23 . Add this to your certain digits: 12.2 mL. Use the balance to measure the mass of the weight. Add approximately 25 ml of water to your beaker.1 g versus 120. Mass _________________ Based on your observations. always record every digit the balance displays. Measure the volume of this water using the beaker and record it below. With a digital balance. 3) Measuring mass Obtain something with mass. Record your units. Always use the smallest graduations on your instrument when taking a measurement Length (cm) = ______________________ Length (mm) = ______________________ Obtain a meter stick and record the length in meters (if the meter stick is graduated to the mm. 4) Comparison of a beaker and a graduated cylinder Take out a 100 or 150 -ml beaker and examine its graduations. be sure to add the appropriate number of digits in your response).2) Measuring lengths Now find a small plastic ruler and measure your pencil. how many decimal places does this balance report? Is there a difference in reporting 120. using the graduations on the beaker. Complete the following statements: The beaker is graduated every _____ mL. _______________ mL (using a beaker) 24 .10 g? Explain how the meaning of the measurement changes. Length (m) = ______________________ Does your measurement have a different precision when you use a meter stick? Explain. The measurement should be recorded to every ______ mL. Suppose a 10-mL graduated cylinder has an uncertainty of about ± 0. Thinking about what you learned earlier (looking at the graduations).How many significant figures can you report based on the graduations? ________ sig figs (circle the digit above that was estimated). Always use units. 25 . If it is filled with water to the mark (10-mL). the volume measurement using the beaker or the graduated cylinder? Given a choice. Remember to estimate one digit more than the level of graduation (image B is correct. record a measurement for the volume of water. which glassware would you use to measure volumes more precisely? Explain.01 mL. Now take the contents of your beaker (the 25 mL of water) and pour it into your 100 or 150 -ml graduated cylinder. Use the correct number of significant figures in your answer. the volume should be reported as (include the necessary decimal places): _______________ mL Provide a measurement for each of the following pieces of glassware filled to the level indicated by the arrow. _______________ mL(using a 100 or 150 -ml graduated cylinder) Which is more precise. do not modify). One way of determining the risk of using chemicals is to read the Material Safety Data Sheet (MSDS). Look up the MSDS for both Hydrochloric Acid and Sodium Hydroxide. Part B Prepare your lab notebook for the experiment. aluminum chloride. Additionally. including ammonium iodide. 26 . Read through the MSDS and determine the steps that need to be taken in case of accidental skin exposure (your most common risk in the lab). This includes stating the purpose of the experiment and summarizing the procedure in a bulleted-list format (be sure to include space for observations). What steps need to be taken if there is skin exposure? Note particularly the danger of getting NaOH in your eyes and be sure to wear goggles at all times in the lab! 3. silver nitrate. Some of the chemicals you will use this year are hazardous.Pre-Lab: Who has the same solid that I have? Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. These documents provide a wealth of information regarding the safety risks of each compound. search the internet for two interesting factoids regarding the chemical compound you choose. 2. lithium carbonate. At the start of your lab. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. In this lab. sodium acetate. calcium carbonate. Do a web search with the key words “MSDS and Lead Nitrate”. and potassium iodide. there are many possible unknown compounds. Pick one of the listed compounds and read the MSDS on it. Your goal is to identify other students in class who have the same compound that you have. Guided by your TA. You should identify the people in your lab section that had the same substance and then run some confirmatory tests to verify that the solids are the same. You will be given a sample of an inorganic solid and will determine your sample’s properties. dissolves most salts. In this experiment you will be trying to determine who else in your lab section has the same unknown compound as you. and its reactivity. Others are less able to form regular patterns so their solids are less geometric. This colored light is a characteristic signature of the element. acidic or basic. it’s difficult to move these ions apart. 2) Conductivity of aqueous solutions: When dissolved in water. its relative melting point. You will see as you progress throughout the year in chemistry that compounds can be classified in many different ways (ionic or molecular. Comparisons of different samples may be made by doing side-by-side analysis using the same techniques. describe. We know that water is clear. metals or nonmetals. 3) Crystalline or amorphic: As a result of the types of bonds in the compound. its appearance in a flame. it is important to be able to accurately determine.. Think about how you might classify water on the basis of easily observable properties. To enable classification. Crystalline compounds are hard and brittle because the ions are locked tightly into place by their electronic interactions. the electrical conductivity of the substance and its aqueous solution. while others have lower melting points. such as: the solid’s relative solubility. has a density of 1 g/mL and is composed entirely of hydrogen and oxygen atoms in a definite ratio (this list could go on and on). most discoveries are not made by just one person. These emitted colors can be used to 27 . This will be done while learning some basic techniques that are used for the analysis of chemical compounds. and when they do move apart. boils at 100°C. some compounds dissociate into ions. Upon further heating. your lab section will determine a method for sharing (reporting) your observations. freezes at 0°C. some compounds may decompose into simpler compounds or burn. like a flame."Who has the same solid that I have?" Science is generally a cooperative collaborative affair. some substances form very regularly shaped crystals.). These dissolved ions move through the solution and thus conduct electricity. the acidity/basicity of the compound’s aqueous solution. It is important for scientists to be able to communicate their data and be prepared to share data or samples. 4) Flame test: Some atoms emit characteristic visible colors when excited by an energy source. which is a consequence of the electronic structure of the element. colorless. As a result. the whole crystal typically breaks. and compare the chemical and physical properties of compounds. Some compounds have melting points greater than 200°C.. This is a list of some of the physical and chemical properties that you will investigate during this lab exercise: 1) Melting points: A substance’s melting point temperature will depend on the bonding type or intramolecular forces in the sample. 2011 28 . This lab is based upon the journal article "Who Has the Same Substance that I Have?": A Blueprint for Collaborative Learning Activities. n. basic or neutral aqueous solutions: Some substances will make a solution acidic or basic when they dissociate into ions when dissolved in water. 24 Aug.identify the elemental composition of a substance. Some ions have the ability to act as acids in solution while others act as bases. When a substance is dissolved in water. potassium produces a violet color while lithium will emit a vibrant red. Coppola. 5) Acidic. For instance. Brian P. Web.d. Lawton. 1120 and “Identification of Ionic and Molecular Compounds”. these properties can easily be tested using pH paper. Journal of Chemical Education 1995 72 (12). 6) Reactivities: Each compound has a characteristic reactivity that may or may not be easily elucidated. reactivity patterns may become visible. http://tinyurl. n. Richard G. gas or change in color. Reactions are usually visualized by looking for the formation of a solid. By mixing an aqueous solution of the unknown with an aqueous solution containing another compound.com/3jf6oq6.p. oxidizer. corrosive. health and environmental hazard corrosive corrosive Toxic none health and environmental hazard none none toxic 29 .Equipment Information Each bin should contain: 1 – plastic well plate Chemical Safety Information Same Solid Chemical Lead nitrate Hydrochloric acid Sodium hydroxide Ammonium iodide Sodium acetate Lead acetate Calcium carbonate Sucrose Lithium carbonate Hazards toxic. Determine if the aqueous solution is acidic. Add 5 drops of 0. Determine conductivity in the solid state a. Pour a little of your unknown aqueous solution into one well of the microplate. Mix with your scoopula. Using a magnifying glass. insoluble means that the sample is cloudy or that there is undissolved solid left in the test tube). use the conductivity meter to check the conductivity of deionized water and record your observations. Do not discard the contents of the microplate! 5. Clean your test tube. Obtain a small (pea-sized) sample of your assigned unknown. Add 5 drops of 1M hydrochloric acid (HCl) to the first well. The paper is normally orange. e. Dip your scoopula in your solution and wipe it on a piece of pH paper. and Rrecord your observations. c.) b. examine the sample and record your observations in your lab notebook 2. c. It will turn red if the solution is acidic or blue if basic. Waft a watermoistened piece of pH paper over the third well to see if a gas (ammonia) is produced. Record your observations as S = soluble or IN = insoluble (soluble means that a clear solution has formed. b. Add 2 cm or approximately 1 fingers width of deionized water to the test tube. Record your observations 3. Add ½ of your “pea-sized” sample of the unknown to a small test tube. (The symbol M represents molarity. 30 . d. a. Record your observations 6. Using a dropper. the greater the concentration. [Be sure the probes are dry!] b. As a control. Do not discard the contents of the test tube! 4. d. such as “turned cloudy” or “no change”. basic or neutral. b. b. (The wells should not be full. Determine the conductivity of the solution made in step 3 a. Physical characteristics a. Record your observations. b. multi-well plate and dropper as directed by your TA. The greater the molarity. Using the sample obtained previously.PROCEDURES 1. equally divide your solution amongst three wells (including the one you already used) in your multi-well microplate. a unit of concentration. f.0 M sodium hydroxide (NaOH) to the third well. test for conductivity. Determine the relative solubility of each unknown a. and the pH paper would turn blue in the presence of ammonia gas. a.1 M lead(II) nitrate (Pb(NO3)2) to the second well. Add 5 drops of 1. test for electrical conductivity using the conductivity meter supplied by touching the probes to the sample.) c. Note: aqueous solutions of ammonia are basic. Determine the reactivity of your unknown. Record your observations. Record your data as indicated by your TA and identify all people in the lab section having the same substance. Using a scoopula. sodium acetate. Run some confirmatory tests to verify that you have the same compound as the other groups. c. based on your collective observations. citric acid. f. DATA Guided by your TA. Monitor the time required to melt. d. calcium carbonate. 10. lithium carbonate. Observe the color of the flame. All remaining solutions and solids must be placed in the properly labeled waste jar. c. Record your results. Touch the wet toothpick into the previously unused half of your unknown solid to pick up a small quantity of the solid. heat no longer than 20 s). aluminum chloride. Possible unknowns include: ammonium iodide. you will construct a table to report your results and observations. obtain a small portion (1/2 of the amount remaining from step 1) of your unknown solid. Flame test (done in the fume hood) a. sugar. Soak a toothpick in water for 1 minute. 11. Carefully light the gas burner (unless it is still burning from the flame tests) b.don’t heat any substance longer than 20 seconds!) e. b. Place the portion of the toothpick with the solid on it in the hot part of the flame. 8. potassium iodide. For the TA The most important practical aspect of setting up this experiment is to ensure that the identification is based on the experimental data that are collected by the students. Determine the relative melting point for each unknown compound (done in the fume hood) a. silver nitrate. 9. some of which are hazardous chemicals. Note: some substances will not show a positive flame test (no color change). Clean the scoopula as directed by your TA. Please discuss contamination and how to avoid contaminating the stock solutions and unknowns. 31 . Your lab area should be wiped clean and all glassware and equipment should be placed in your lab drawer. Clean up. (Any substance that will melt under normal lab conditions will do so quickly .7. Choose two tests that makes your unknown unique from the other substances and do a side-by-side comparison to verify your conclusion. based upon whether or not the substance melted and how long it took (remember. Carefully heat the sample on the edge of the scoopula 1 inch from the very top of the flame (not the hottest part of the flame). Who has the same solid that I have? Lab Report: Your report for this lab should include the following sections: Abstract: Your first-draft abstract for this lab should be written as part of a post-lab discussion led by your TA. Provide some error analysis. Indicate how your results for your solid caused you to identify its match. How did your solid’s properties differ from the others? Describe how/why you chose the confirmatory tests and the corresponding results. 32 . For instance. Data: Include the complete set of class data. You will refine this later on your own. Summarize the results obtained for your unknown and the matching unknown(s) in the confirmatory tests. Introduction: Describe what you did (tested and compared several unknown compounds) and provide a bit of insight as to what techniques were used. Explain why you did this experiment (to match the properties of your unknown with other unknown substances). Results: Write a paragraph explaining the results of this experiment Write a sentence or two (or a table) summarizing the matching and nonmatching characteristics to indicate which other solid matched your groups unknown (specifically list which results were similar and which results were different). what sort of weaknesses do you see in the procedures or the way the data were reported that may have caused some ambiguity? Same Solid Lab Report: Subm it you r report on tim e and to your TA in the d ropbox on D2L. Be sure to indicate which unknown number you tested and the matching unknown number(s) Discussion: Write a paragraph that discusses the following points: Discuss how you matched your unknown sample with the other(s) in your lab section. Pay attention to the directions above about formatting tables. Be sure to include your unknown and the data from the confirmatory test. ) Part B Prepare your notebook for the lab. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. 33 . How many milliliters are in a 12 ounce can of soda? 4. what is the density of the solution? 2.25 g/ml would you expect this object to float on water? 3. What is the purpose of constructing a calibration curve? (If you’re not sure.5 g and a volume of 100. A solution has a mass of 109. At the start of your lab. This includes stating the purpose of the experiment.Pre-Lab: How much sugar is in a can of coke? Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. An object has a density of 0. you might want to watch the “weblet” online presentation for this lab before you answer.0 ml. Calibration curves are constructed using known quantities. Determine the amount of sugar (in grams) in a can of coke 2. When comparing Coke and diet Coke. Coke can be represented as a solution of sugar dissolved in water. . so does the density of the resulting solution. you will use this relationship to determine the amount of sugar in Coke. Whenever you are asked to produce a graph from laboratory data (either by hand or using a computer program) all of the following criteria must be met: 1. The relationship between the amount of dissolved sugar and the density of sugar water solutions will be determined using a calibration curve. Learn how to make solutions quantitatively 3. The sugar makes Coke more dense than diet Coke. we will introduce what must be included in any scientific graph.How Much Sugar is in a Can of Coke? GOALS: 1. in which case you substitute a figure caption below the figure for the title) 2. Intensive properties like density are independent of the amount of substance and thus the density of two different solutions can be compared without needing to have the same volume of the two solutions. SCIENTIFIC GRAPHS: This experiment will also serve to introduce you to scientific graphing. called standards. you will prepare a graph of density vs. mass of dissolved sugar . This makes it possible to determine the mass of sugar in Coke by comparing it to solutions with known concentrations of sugar. After obtaining the mass of these standards using an instrument called a balance and calculating the density of each solution. In this case. Calibration curves allow you to determine the content of an unknown by comparing it to observations made on the standards with known values of the property being measured. As the amount of sugar dissolved in a given volume of water increases. Learn how to make and use a calibration curve INTRODUCTION: If you were to measure out identical volumes of Coke and diet Coke. You will then determine the mathematical relationship between the two quantities. To understand why. M D Density = Mass Volume V Density is a convenient quantity because it is independent of the volume used (scientists describe properties like this as intensive). you would find that the two liquids have different masses. The graphed data must take up the full space of the graph 34 . Here. it is found that Coke is more dense than its sugar free relative. This difference in the mass of the two liquids is best discussed by looking at the mass per unit volume (or density) of the two liquids. The main difference between the two “solutions” is the presence of the dissolved sugar in Coke that is absent in diet Coke. All graphs should have a title (except when included in a report or other scientific writing. a molecular view of the two substances is useful. To a first approximation. you will prepare standard solutions of known volume with a known amount of dissolved sugar. Once the relationship between density and sugar content is determined. Both axes must be labeled with a name and units 3. A linear relationship exists between mass and volume.009 50 Mass (g) 40 30 20 10 0 0 10 20 30 40 50 60 Volume (ml) Figure 2: The relationship between Mass and Volume for Water. Varying volumes of water were massed. The equation should be included in the caption below the graph. Figure 1 shows an example of an acceptable scientific graph of raw data.edu/labwrite/res/gt/gt-menu. the line should be shown on the graph. A best fit line was calculated using Microsoft Excel yielding the following equation: y = 1.0015x + 0. Graphing Using Microsoft Excel: An excellent tutorial on graphing with MS-Excel can be found at the following website: http://www. Figure 2 demonstrates the proper way to represent a linear fit on a graph. Gridlines should only be included if they enhance the understanding of the graph.ncsu. When a “best-fit” line to the data is computed and used.009.html 35 . 5. “independent” – for example the graph below is Mass (in grams) vs. The independent variable is the x-axis and the dependent variable is the y-axis. and the graph is referred to as “dependent” vs.4. The Relationship Between Mass and Volume for Water 60 50 Mass (g) 40 30 20 10 0 0 10 20 30 40 50 60 Volume (ml) Figure 1: The relationship between Mass and Volume for Water The Relationship Between Mass and Volume for Water 60 y = 1.0015x + 0. Volume (in mL) 6. With the program open. If it appears correctly. click next. Click next. enter the data to be graphed in the cells. 6. A preview of your chart will appear. Select the ‘display equation on chart’ button and click ok 36 . 3. Select ‘linear’ as your regression type 3. Click and drag the mouse to highlight all the data to be graphed. Select the ‘options’ tab in the popup window 4. Click on the chart wizard icon 4. With the graph selected. Enter a chart title and the axis labels and click finish 7.This is a list of the basic steps necessary to graph data and do a linear regression (the generation of a “best-fit” line) using Excel: Basic Graphing: 1. With the chart selected you can also access the title and axis labels by selecting ‘Chart’ then ‘chart options’ from the drop down menu Adding a Linear Trendline to a Graph: 1. select ‘Chart’ then ‘add trendline’ from the drop down menu. 2. Enter x data in one column followed by y data in an adjacent column. Choose XY (Scatter) for the chart type and the unconnected points icon for the Chart sub-type 5. 2. Equipment Information Each bin should contain: 1 . Use the correct one for each stage of your experiment b. Parafilm is reusable.50mL volumetric flask Notes: a. c. Use the 50mL volumetric flask for the final coke measurement. not the 100mL flask from your drawer. There are 2 waste streams for this lab. Chemical Safety Information Sugar In Coke Chemical Sugar Coke Hazards none none 37 . Once the sugar has completely dissolved. 2. Add water carefully to the fill line. Put the stopper in the flask or cover the top of the flask with Parafilm and invert it ten times to ensure that the solution is thoroughly mixed. Swirl the flask to dissolve the sugar. Weigh the empty flask and record the mass. (The difference between this mass and the first one is the mass of the solution – the volume is 50 ml. 5. Be sure that you add an appropriate caption to the table. the mass of sugar used for each solution is found by subtracting the mass of the empty stoppered flask from the mass of the stoppered flask containing sugar.) 3. Volumetric flasks are designed to accurately contain a specific volume.) 4. Weigh out the desired mass of sugar in a weigh boat.PROCEDURE: Calibration Curve: You will make one sample without sugar and five sugar water solutions to start. Each solution should have a different amount of dissolved sugar covering a range from about 1 – 8 g of sugar per 50 mL of solution volume.) To accurately know the mass of sugar used in each of your five standard solutions. (This mass does not need to be recorded. In later experiments. Weigh the flask containing the solution and record the mass. a table is supplied for you. if you carefully followed these instructions. Weigh the empty flask and record the mass. For the sample without sugar: 1. When filled to the marked line. if you carefully followed these instructions. it is best to bring the fluid to the line carefully by using a wash bottle or eyedropper to assure that the flask is not overfilled (causing you to have to start over). Add the sugar to the flask. 3. (The bubble should run all the way down the neck of the flask to the stopper each time you invert the flask. Weigh the flask containing the solution and record the mass. To make the solutions in a quantitative manner. the flask accurately holds the stated volume (these devices are called TC for “to contain”). 2. add water carefully to the fill line.) 6. Add water to the flask until it is approximately half way to the fill line. Do not shake the flask. (The difference between the two masses is the mass of sugar. they must be prepared in volumetric flasks. (The difference between this mass and the first one is the mass of the solution – the volume is 50 ml. follow these steps: 1. you will be expected to produce your own data tables for your notebook and the Results section of your lab reports. In this experiment. When putting the last bit of solvent into volumetric flasks.) As noted above. 38 . See the report guidelines section of the manual for a discussion on this. The mass of the solution is found by subtracting the mass of the empty stoppered flask from the mass of the stoppered flask containing the solution. Weigh the flask containing the sugar and record the mass. Volumetric flasks are marked with a fill line. Below is an example of an acceptable table to present the data from this experiment. clean 50 mL volumetric flask before carefully filling the flask to the fill line with the flat Coke provided. Put the used Coke in the provided waste jar.Mass of Empty Flask + Stopper (g) Pure H2O Mass of Flask + Stopper With sugar (g) ------ Mass of Mass of Flask + Sugar Stopper (g) With solution (g) ------ Mass of Solution (g) Density of Solution (g/ml) Dissolved sugar per mL solution (g/ml) 0. In order to find the mass of sugar in one can of Coke. Alternatively. then drawing a vertical line down to the x-axis (mass of sugar per mL solution) you can graphically determine the amount of sugar dissolved in each mL of Coke. Calculate the percent error in your determined value.0 Flask 1 Flask 2 Flask 3 Flask 4 Flask 5 Using the data in the above table. Graphs must be prepared using the computer. you can invert the relationship given by the calibration line equation to solve for “x” from the observed density (“y”). if needed. Determine the amount of sugar in a can of Coke: Weigh and record the mass of a dry. Both methods should give the same result for mass of sugar per mL of Coke. RESULTS: When a linear relationship exists between two quantities (density and amount of sugar) it is only necessary to measure one of the quantities (density) and know the relationship (found from your calibration curve) before the other quantity (amount of sugar) can be determined. but could be controlled in making the standards). One liter contains 33. you will need to consider the volume of a can of Coke (12 ounces). This graph represents the relationship between the density of the sugar water solution (something that can be measured) and the amount of dissolved sugar in the solution (something that cannot be measured directly.8 fluid ounces. construct a graph of density of solution (y) vs. Report the equation for the line on the graph. By finding the density of coke (y-axis) and drawing a line to your calibration curve. dissolved mass of sugar per mL of solution (x) and fit the data to a linear relationship as described above. Your TA will assist you with this. 39 . Determine the density of the Coke. based on nutritional information given on the label on a can of Coke. Weigh and record the mass of the flask containing Coke. 0 Flask 1 Flask 2 Flask 3 Flask 4 Flask 5 40 .Sugar in Coke Lab Report Sheet: Name_____________________________ Date______________ Lab Section______________ Provide a brief statement of the purpose of this activity and explain the idea behind a calibration curve. Data: Mass of Empty Flask + Stopper (g) Pure H2O Mass of Flask + Stopper With sugar (g) ------ Mass of Mass of Flask + Sugar Stopper (g) With solution (g) ------ Mass of Solution (g) Density of Solution (g/ml) Dissolved sugar per mL solution (g/ml) 0. Results: Copy and paste your calibration curve in the space below. Make sure you have formatted it as described in the Report Guidelines located in the front of this manual. Report the amount of sugar in 1 mL and 1 can of coke, showing any necessary calculations. Present the percent error for your amount of sugar in a can of coke with respect to the number given on the label, showing any necessary calculations. 41 Based upon if your value is greater or less than the labeled sugar content, provide any valid sources of error Sugar in Coke Lab Report: There is not a form al lab report for this lab. Com plete the above pages using the Microsoft version of this file that is available for d ow nload on the lab D2L page. Once the w orksheet is com plete, subm it the w orksheet on tim e and to your TA in the d ropbox on D2L. 42 Pre-Lab: A Cycle of Copper Reactions Part A Answer the following questions in your lab notebook (be sure to show your work for any calculations): 1. In chapter 4 of your text, read about the d ifferent types of reactions. What causes a precipitate to form w hen certain com binations of aqueous salt solutions are m ixed ? 2. Locate the solubility table in chapter 4 and sum m arize w hich ions are generally soluble and w hich are generally insoluble. 3. Look up the MSDS for nitrogen dioxide gas. What dangers does it present and what steps need to be taken to avoid exposure? 4. This is one of the most dangerous experiments during this term because of the risk of exposure to dangerous chemical substances. For example, concentrated nitric acid is a particularly nasty solution; what happens when it comes in contact with skin? What safety precautions need to be taken to avoid exposure? All of the acids and bases used in this experiment are also potentially dangerous and should all be handled carefully. Part B Prepare your notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. At the start of your lab, remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. 43 A Cycle of Copper Reactions GOALS: 1. Cycle solid copper through a series of chemical forms via aqueous-phase reactions 2. Learn about and identify different types of aqueous reaction types 3. Calculate percent recovered copper after all of the transformations INTRODUCTION: This experiment will cycle elemental copper through a series of five reactions summarized below: Cu(NO3) NaOH Cu(OH) 2 2 HNO3 heat Cu Zn, HCl CuSO4 H2SO CuO 4 The cycle will both begin and end with pure elemental copper. At different stages of the cycle, copper will be present in different chemical forms. At times copper will be present in solid compounds and other times in ionic form. Each chemical change that copper undergoes is observable as a change in the physical properties of the solution (or precipitate). As you perform each reaction, be certain to observe and record your observations of all physical changes. At this point in the term, you should have been introduced (in the lecture class) to three different types of aqueous reactions: precipitation reactions, acid-base reactions, and oxidationreduction (or redox) reactions (in addition, the text may have discussed gas-forming reactions). In precipitation reactions, soluble cations and anions combine to form an insoluble compound that leaves the solution as a solid precipitate. In acid-base reactions, an acid and base react to produce water and a salt. Redox reactions involve the transfer of electrons. As you go through the series of reactions you should be able to classify each reaction (with the exception of reaction 3) as one of the three above-described types of aqueous reactions. Reaction 1: The first reaction proceeds according to the following balanced chemical equation: 4 HNO3 (aq) + Cu (s) Cu(NO3)2 (aq) + 2 H2O (l) + 2 NO2 (g) 44 dried and weighed. The result of this reaction is that copper changes from its elemental state (charge = 0) to an aqueous. Reaction 2: The second reaction then converts the aqueous Cu2+ into the solid copper (II) hydroxide (Cu(OH)2) through a precipitation reaction with sodium hydroxide according to the following balanced chemical equation: Cu(NO3)2 (aq) + 2 NaOH (aq) Cu(OH)2 (s) + 2 NaNO3 (aq) Reaction 3: The third reaction takes advantage of the fact that Cu(OH)2 is thermally unstable. zinc and copper exchange physical (and oxidation) states in and out of acidic solution. When heated.In this first reaction. elemental copper metal is reacted with concentrated aqueous nitric acid solution. washed. 45 . Cu(OH)2 (s) + heat CuO (s) + H2O (l) Reaction 4: When solid CuO is reacted with sulfuric acid. the copper is returned to solution as an ion (Cu2+) according to the following acid-base metathesis / double displacement reaction equation. The solid copper can then be collected. Some copper is bound to be lost in all of the chemical transformations. CuSO4 (aq) + Zn (s) ZnSO4 (aq) + Cu (s) Here. so the percent recovery (mass copper remaining/initial mass x 100%) is expected to be less than 100%. Hydrochloric acid is then used to dissolve any excess zinc. Cu(OH)2 decomposes (breaks down into smaller molecules) into copper (II) oxide and water according to the following decomposition reaction equation. CuO (s) + H2SO4 (aq) CuSO4 (aq) + H2O (l) Reaction 5: The cycle of reactions is completed in this reaction where elemental copper is regenerated according to the following oxidation-reduction reaction equation. Thhis reaction changes copper from its ionic state (Cu2+) to its elemental state by exchanging electrons between zinc and copper. ionic state (Cu2+) in an oxidation–reduction reaction. Chemical Safety Information A cycle of copper reactions Chemical Nitric acid Sodium hydroxide Sulfuric acid Hydrochloric acid Zinc Copper Hazards oxidizer. Then. If any concentrated acid is spilled. Concentrated acids are being used this week. clean up the solid waste using the hand broom and dustpan found near the broken glass container. use spill neutralizer to neutralize the spill. corrosive corrosive corrosive corrosive flammable. b. Notes: a. environmental hazard 46 . environmental hazard flammable. Restrict their use to the hoods.Equipment Information There are no bins this week. ask your TA or repeat the addition of 1 mL of sulfuric acid as necessary. Once the conversion is complete. add 20 mL of distilled water to the flask.0 M H2SO4 (sulfuric acid) to the flask and swirl the mixture for 1 minute. Avoid getting it on your skin or clothing. heat the flask a little longer and/or ask your TA to take a look. Once you are sure that all the gas has been removed in the fume hood. Reaction 3: With occasional stirring. you may return to your workbench. Reaction 2: While stirring with a glass rod. slowly add 20 mL of 6. Be sure to record all observations. Be sure to record the actual amount used to the nearest milligram.0 M NaOH to the flask. 47 . Add 50 mL of nearly boiling hot water to your reaction mixture. Once the CuO has resettled (give it about 5 minutes to settle). If the black solid has still not dissolved. Once the reaction is complete and the gas has dissipated. heat ~ 200 mL of distilled water. Reaction 4: Carefully add 5 mL of 6.0 mL of concentrated nitric acid. All of the black CuO should dissolve and be gone at this point. DO THIS NEXT STEP IN THE HOOD!! NO2 gas is toxic! In the fume hood. carefully decant the supernatant liquid. Be sure to record all observations. Be sure to record all observations. removing as much as possible without losing the desired product (CuO). Place the copper at the bottom of a 250 mL Erlenmeyer flask. swirl the flask until all of the copper has dissolved. If the conversion does not appear to be complete (not all of the blue Cu(OH)2 has disappeared). add the nitric acid to the flask containing the copper. slowly heat the flask on a hot plate until the solution just begins to boil. Do not let the solution boil vigorously. In a clean beaker. The nitric acid should completely cover the copper.5 g of copper. carefully measure out 5. Be sure to record all observations. Do not breathe vapors. Weigh out about 0. In a graduated cylinder. If there is still black solid in your reaction mixture. Remaining in the hood. If you do get any on your skin or clothing. remove the flask from the hot plate and allow the CuO to settle. wash it off immediately with running cold water. add an additional 1 mL aliquot of the sulfuric acid and swirl the mixture for an additional minute. At this point you should notice that the blue Cu(OH)2 has been converted to black CuO.PROCEDURE: Be sure to discard all waste in the waste jars as directed by your TA Reaction 1: Caution: Concentrated nitric acid is hazardous. If you do get any on your skin or clothing. Decant the water after each wash. Gently heat the copper on a hot plate to evaporate any remaining methanol and dry the copper. Stir until the supernatant liquid is colorless (not blue).Reaction 5: Caution: Concentrated hydrochloric acid is hazardous. but do not boil the solution. 615. Allow the copper to settle and decant the methanol. so do not breathe the vapors. the flask may be returned to the workbench. the solution can be murky and grey. 52. there is some possibility of producing more nitrogen dioxide gas. Remain in the hood and add 5 mL of distilled water followed by 10 mL of concentrated hydrochloric acid.0 g of 30-mesh zinc or zinc powder. Wash the copper with an additional 5 mL of methanol. RESULTS: Once the mass of recovered copper is known (difference between the pre-weighed beaker plus copper and beaker alone). Be sure to record all observations and the final mass of the copper. wash it off immediately with water. Hydrogen gas is generated which is extremely flammable. Do not breathe vapors. 48 . J. Ed. Condike. Once the hydrogen gas evolution has completely stopped. Avoid getting it on your skin or clothing. Chem. remove the copper from the hot plate and allow the beaker to cool before determining the mass of the recovered copper. Transfer the solid copper to a clean pre-weighed beaker. Once dry. Add (all at once) 1. DO THIS NEXT STEP IN THE HOOD. You may warm the mixture on a hot plate to speed up the reaction. Decant the supernatant liquid. the percent recovery can be calculated from the following formula: Percent recovery = (mass of copper recovered/initial mass of copper)*100% REFERENCE: 1. more acid (in 1 mL aliquots) can be added. Using a wash bottle to wash the copper metal into the dish can facilitate the transfer. Additionally. Once the evolution of hydrogen gas has become very slow. 1975. GF. Wash the copper at least twice with about 5 mL of distilled water each time. Zn (s) + 2 HCl (aq) ZnCl2 (aq) + H2 (g) If the hydrogen gas evolution stops before all of the solid zinc has been removed. remove the flask from heat and carefully decant the liquid. The hydrochloric acid removes any excess zinc according to the following balanced chemical equation. There should be no open flame in the room. and if so. and if so. Include the initial and final mass of copper Results: Report which of the observations that suggested that a chemical reaction took place (a precipitate formed. how you did it and what your results (percent recovery) were. (the black solid was…). what does the blue represent and how would it change your % recovery? Did you lose any black solid while decanting. Introduction: Begin with a statement of the purpose of the experiment (observe various reactions involving copper). what were you discarding and how would it change your recovery? Was there any solid Zn remaining at the end. there was a formation of a gas….) and relate those observations to the components of the balanced chemical reaction. 49 . Report your percent recovery of copper Attach your calculations to the back of your report Discussion: Discuss the experiment (reaction types involving copper) Discuss your percent recovery and possible sources for loss of copper during the reaction cycle. the solution’s color changed. Copper Cycle Lab Report: Subm it you r report on tim e and to your TA in the d ropbox on D2L.Copper Cycle Lab Report: Your report for this lab should include the following sections: Abstract: Keep it to three sentences and be sure to discuss what you did. Provide any relevant background and key concepts (there are different types of chemical reactions) Include useful chemical equations (include each balanced equation and what type of reaction it is). how would that change your recovery? Was your copper completely dry when you massed it at the end of the lab. and how would it change your recovery? Note that some of the items above increase calculated recovery and some decrease it – identify the direction of the error in each case. Here are some things you can think about: Did any of your discarded waste have a blue color. if so. Data: Include each balanced chemical reaction and your observations for each reaction. This includes stating the purpose of the experiment. H ow m any m oles of alkali m etal carbonate w ere reacted ? 6.723 g of barium carbonate from 2. What is the m olar m ass of the alkali m etal carbonate? Hint: remember the units of molar mass are g/mol. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection.53 g of lithium carbonate w ith excess hyd rochloric acid (balanced chem ical reaction is given below )? Li2CO3(s) + 2HCl(aq) --> 2LiCl(aq) + H2O(l) + CO2(g) 3. Look up the MSDS for barium chloride. The reactants w ere M 2CO 3 and barium chlorid e. Write the balanced chem ical reaction for this synthesis. 7. H ow m any m oles of barium carbonate w ere prod uced ? 5. In the synthesis of barium carbonate from an alkali m etal carbonate (M 2CO 3 w here M is one of the alkali m etals) a stud ent generated 3.001 g of their alkali m etal carbonate. w hat is the m olar m ass of M? Which alkali m etal is closest to this m olar m ass? 8. At the start of your lab. 4. The chem ical form ula for the alkali m etal carbonate is M 2CO 3. What is the law of conservation of mass? 2. 50 . How toxic is this substance? What steps need to be taken if there is skin exposure or accidental ingestion? Part B Prepare your notebook for the lab. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA.Pre-lab: Which Alkali Metal Carbonate? Part A Answer the following questions in your lab notebook (be sure to show your work for any calculations): 1. H ow m any m oles of CO 2 can be prod uced by the com plete reaction of 1. and a third method to confirm them. Your task is to determine which alkali metal carbonate composes the sample. At the end of this experiment you will prepare a report giving your experimental results. In natural deposits. 51 . Based on your experience you will be able to recommend which procedure you would use if you had time and resources for only one technique.Which Alkali Metal Carbonate? The Problem In a search for a good cleaning formulation (as in laundry detergent or a degreaser for metal parts) alkali metal carbonates are found to be useful. Imagine that you are an analytical chemist and have received a sample of a pure alkali metal carbonate from a newly-discovered deposit. This will include the identification of your alkali metal carbonate. error discussion for both methods. You will do an experiment to determine the atomic weight for the alkali metal in the carbonate you have and thus which alkali metal is present. these carbonates may occur as crystals of a single alkali metal carbonate (such as lithium carbonate) or as amorphous solids with several of the alkali metal carbonates co-deposited. You will use two different methods to ensure confidence in your results. You will also evaluate sources of error as you compare your experimental values with the expected value for the atomic weight of the alkali metal. and a rationale for the preferred method you would recommend to the lab. Equipment Information There are no bins this week. c. Notes: a. Be sure to use the correct one for each stage of your procedure. Dry the barium carbonate (BaCO3) using the filter flask. There are two waste streams this week. Use two pieces of filter paper for filtering. Chemical Safety Information Which alkali metal carbonate? Chemical Potassium carbonate Lithium carbonate Sodium carbonate Barium chloride Hydrochloric acid Hazards toxic toxic toxic toxic corrosive 52 . b. Heat the BaCO3 formed to “digest” the precipitate (causing the precipitate to form larger aggregates). The reaction involves your aqueous carbonate reacting with barium chloride (BaCl2) in a precipitation reaction. but massed to mg accuracy and recorded in your notebook) sample of your carbonate to a 250 mL beaker. To precipitate the barium carbonate. Na2CO3. Wash the precipitate with water using the vacuum to pull the water through the filter and allow the solid to dry for 10 minutes with the vacuum still on. 8. You will isolate and weigh it. add 20. 4. PROCEDURE: Be sure to discard all waste in the waste jars as directed by your TA 1. 2. The product is an insoluble barium carbonate. Dispose of in the proper waste container (see TA if you are unsure of the proper procedure). Carefully remove the solid and filter paper from the funnel and place your product on a pre-weighed watch glass. Weigh the combined dry solid and filter paper and record the mass in your notebook.0 M BaCl2 solution to the sample and stir until well-mixed. Be sure to record the masses in your notebook. Filter the barium carbonate using the filter paper in a Buchner funnel using vacuum filtration as demonstrated by your TA. Determine the molar mass of the unknown metal carbonate (not the metal itself) and compare it to the molar mass of the possible alkali metal carbonates (Li2CO3.Gravimetric analysis (Method 1) In this part of the experiment you will perform a synthesis and use reaction stoichiometry to identify your unknown alkali metal carbonate. and hence were present in the 1. CAUTION: The filtrate (solution left after filtration to isolate barium carbonate) contains excess Ba2+. 5. K2CO3). 6. Calculate the mass of barium carbonate by difference (removing mass of watch glass and filters). How many moles of barium carbonate is this? Use stoichiometry to determine how many moles of M2CO3 reacted to produce the barium carbonate. Add a 0. Add 50 mL of water and stir until the carbonate is completely dissolved. Allow the solid to dry until near the end of the lab period.0 mL of 1. This involves boiling the solution for 5 minutes with little agitation.5 g (approximately. 7. What alkali metal carbonate is your sample most likely to be? Where are the errors most likely to enter into the experiment? 53 . 3. CAUTION: Use gloves to handle the barium compounds. It may be necessary to put the watch glass on a hot plate (on LOW heat for 10 minutes) to speed up the drying process. Analysis Use the following questions to lead you to the identity of M: Determine the mass of barium carbonate produced as above.00 g sample you started with. DO NOT DUMP THIS SOLUTION DOWN THE SINK. Weigh one piece of filter paper and your watch glass. 3. What alkali metal carbonate is your sample most likely to be? Where are errors most likely to enter into the experiment? Which direction do these errors bias the final answer (molar mass of metal carbonate)? 54 . Applying the law of conservation of mass you can determine the mass of CO2 evolved.0 mL of 1 M HCl in a pre-weighed graduated cylinder.5 g of your unknown in a pre-weighed 250 mL beaker. Measure 40. How many moles of CO2 were produced? Use stoichiometry to determine how many moles of M2CO3 were present in the 1. Place 0. 2. Measure the mass of the beaker after the reaction has ceased (no further generation of carbon dioxide gas).Simple weight loss (Method 2) In this experiment you will make use of the principle of conservation of mass to determine the identity of your alkali metal carbonate. The balanced chemical equation will allow you to determine the atomic mass of the alkali metal: M2CO3 + 2 HCl -----> CO2 + 2 MCl + H2O Use the procedure outlined below to help you design your experiment and set up your data table.0 g sample you started with. Answer any questions you encounter along the way. PROCEDURE: 1. K2CO3). The metal carbonate (a base) will react with added acid to produce carbon dioxide (CO2) gas. (Determine the actual mass of HCl added. by difference from the starting mass of the reagents. which will leave the system and go into the gas-phase. Na2CO3. You should perform the procedure three times and should obtain a relative deviation [{(largest result – smallest result)/average result} x 100%] of less than 10%. Determine the molar mass of the unknown metal carbonate (not the metal itself) and compare it to the molar mass of the possible alkali metal carbonates (Li2CO3.) Pour the HCl slowly onto the unknown metal carbonate. Analysis: Use the following questions to lead you to the identity of M: Determine the mass of CO2 produced (use the average from the three trials). 4. Table 1: Characteristic Colors for the Flame Test of Certain Group 1 Salts. E. Touch the wet toothpick into the previously unused half of your unknown solid to pick up a small quantity of the solid. Educ. 55 . Place the portion of the toothpick with the solid on it in the hot part of the flame. 1991. The heat excites the electrons in the metal. J. Record your results. 3. the release a photon of light that has a characteristic energy or color (if the photon is in the visible portion of the electromagnetic spectrum). 2. Experiment adapted from: Dudek. This property is used to identify unknown metals and even quantify their amounts present in a sample. Chem. Group 1 Element Li Na K Rb Flame Color Crimson Golden Yellow Violet Blue-violet PROCEDURE: 1. Repeat steps 1-3 if necessary. emits light whose color is characteristic to the metal ion in the salt (see Table 1).Flame test In this experiment you will make use of the fact that metal salts. 68. Soak a toothpick in water for 1 minute. Observe the color of the flame. When the excited electrons relax to what is called their “ground state”. P.948. when heated in a flame. e. Report your findings of the molar mass of your metal carbonate and the identity of the metal carbonate. Any suggested errors should be accompanied with a discussion as to how the error could have been responsible for the error seen (i. Which Alkali Metal Carbonate? Lab Report: Subm it you r report on tim e and to your TA in the d ropbox on D2L. Discussion: Discuss the experiment and any possible sources of error. Attach hand written sample calculations to the back of your report. a molar mass that was lower or higher than expected). Provide any relevant background (balanced chemical equations and brief explanation of the methods used) and key concepts (how the reactions relate to the laws and how they will allow for the determination of the unknown metal). Include the data from your gravimetric analysis Include a data table for your 3 simple weight loss trials Include flame test information Results: Include a results table that summarizes your results from both methods (be sure this includes all trials of each method). and what your results were (molar mass and which metal carbonate). Na2CO3 and K2CO3). Calculate the percent difference for the molar mass of the unknown metal carbonate with the molar mass of each of the three possibilities (Li2CO3. Introduction: Begin with a statement of the purpose of the experiment (identify the alkali metal in the unknown carbonate using three different methods). Additionally. Data: Report the number of your unknown. Phrase this as a recommendation of a procedure for a single-method determination of the identity of an unknown alkali metal carbonate.Which Alkali Metal Carbonate? Lab Report: Your report for this experiment should include the following sections: Abstract: Keep it to three sentences and be sure to discuss what you did. 56 . you should address which method (gravimetric or weight loss) you felt was the most successful (be sure to support your answer with an explanation). Discuss if the flame test supports your identification. how you did it. Use this to assist in identifying your unknown metal carbonate. Could you use cond uctivity to determ ine the equivalence point of this reaction? Why or w hy not? Part B Prepare your notebook for the lab. 57 . d o you expect the cond uctivity to be high or low ? Why? 2. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. Write the balanced chem ical reaction for the titration of strontium hyd roxid e w ith sulfuric acid . This includes stating the purpose of the experiment. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. w hen you have ad d ed an equal num ber of moles of H 2SO 4 to the m oles of Ba(OH )2 originally present? 4. If you ad d excess H 2SO 4. When the cond uctivity probe is placed in a solution of Ba(OH ) 2.Pre-lab: Using Conductivity to Find an Equivalence Point Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. At the start of your lab. Do you expect the cond uctivity to increase or d ecrease as you ad d H 2SO 4 to the solution of Ba(OH )2? Why? 3. w hat should happen to the cond uctivity of the solu tion in the flask? 5. past the equivalence point. Do you expect the cond uctivity in the flask to be greater or less than the original. an electrical circuit is completed across the electrodes that are located on either side of the hole near the bottom of the probe body. As 0. The balanced chemical equation for the reaction in this experiment is: – – Ba2+(aq) + 2 OH (aq) + 2 H+(aq) + SO42 (aq) BaSO4(s) + 2 H2O(l) Before reacting. Ba(OH)2 and H2SO4 are almost completely dissociated into their respective ions. however. and water on solution conductivity.Using Conductivity to Find an Equivalence Point OBJECTIVES In this experiment. When the moles of H2SO4 added equals the moles of BaSO4 originally present. precipitates. you can find the concentration of the Ba(OH)2 solution. Use a Conductivity Probe to monitor conductivity during the reaction. When you have reached a consensus about what will happen during the experiment. prior to the addition of H2SO4.] From the volume used and known concentration of the sulfuric acid. it is important for you to hypothesize about the conductivity of the solution at various stages during the reaction. Barium sulfate is a solid precipitate and water is predominantly in its neutral molecular form. As H2SO4 is slowly added. When the probe is placed in a solution that contains ions (and thus has the ability to conduct electricity). producing BaSO4 and H2O. you will monitor conductivity during the reaction between sulfuric acid (H2SO4) and barium hydroxide (Ba(OH)2) in order to determine the equivalence point. is significantly dissociated. Observe the effect of ions. You will also directly observed the effect of ions. precipitates. as you were asked to do in the Pre-lab. Prior to doing the experiment. and increasing or decreasing. As excess H2SO4 is added (beyond the equivalence point).02 M H2SO4 is slowly added to Ba(OH)2 of unknown concentration. you will Hypothesize about the conductivity of a solution of sulfuric acid and barium hydroxide at various stages during the reaction. in each of these situations? When the Conductivity Probe is placed in Ba(OH)2. or µS/cm.. 58 . and water on conductivity. [In this reaction. sulfuric acid will function as a diprotic acid and barium hydroxide as a dibasic base. The unit of conductivity used in this experiment is microsiemens per centimeter. Neither of the reaction products. INTRODUCTION In this experiment. proceed with the procedure below. This results in a conductivity value that can be read by the computer interface. changes in the conductivity of the solution will be monitored using a Conductivity Probe. Discuss the following questions with your lab partners: Would you expect the conductivity reading to be high or low. Parafilm can be reused. b. Use two pieces of filter paper per filtration. toxic 59 . Chemical Safety Information Using conductivity to find an equivalence point Chemical Sulfuric acid Barium hydroxide Hazards corrosive corrosive.Equipment Information Each bin should contain: 1 – conductivity probe Notes: a. 0 mL of 0. Obtain approximately 60 mL of ~0. Press ENTER. Continue adding 1. Transfer the solution to a clean 100 mL beaker. Using a 100 mL graduated cylinder.00 mL level of the buret. add one 0. Once filtered. Press ENTER to store the first data pair (volume. Before adding H2SO4 titrant.0 mL of the Ba(OH)2 solution). 2. Then add 15 mL of distilled water to the beaker (this step just adds volume so that the probe can accurately measure the conductivity of the solution). filter the Ba(OH)2 solution. c. When the conductivity has is close to 200 µS/cm.) Use a utility clamp to attach the buret to the ring stand as shown here. Dispose of the waste solution from this step as directed by your instructor. (The stopcock at the bottom is open when the handle is aligned with the tip of the buret and closed when it is at a right angle across the tip. Fill the buret a little above the 0. The Conductivity Probe should extend down into the Ba(OH)2 solution to the bottom of the beaker. 02 M H2SO4 solution into a 250 mL beaker. type the current buret reading in mL. a. Handle it with care. Put on gloves and obtain about 60 mL of the Ba(OH)2 solution in a clean beaker. 3. Obtain a 50 mL buret and rinse the buret with a few mL of the H2SO4 solution. In the edit box. 60 . Drain a small amount of H2SO4 solution so it fills the buret tip and leaves the H2SO4 at a mark that is slightly below the 0. b. until the conductivity has is close to 200 µS/cm.02 M H2SO4 to the beaker. Conductivity Probe and beaker containing Ba(OH)2 as shown in the picture above. When the conductivity stabilizes. each time waiting for the reading to stabilize.0 mL increments of H2SO4. CAUTION: H2SO4 is a strong acid. use a graduated cylinder to aliquot 25 mL to each group. again click . and entering the buret reading. Set the selection switch on the amplifier box of the conductivity probe to the 0-2000 µS/cm range. Arrange the buret.5 mL increment and enter the buret reading as above. measure out the appropriate volume of the barium hydroxide solution for all the groups on your bench (each group will need 25. Read and record this volume. Connect the Conductivity Probe to the computer interface. your group will collaborate with the other groups at your lab bench. This process goes faster if one person manipulates and reads the buret while another person operates the computer and enters volumes. You are now ready to begin the titration. click and monitor the displayed conductivity value (in µS/cm). click . 4. Add about 1. In the edit box. CAUTION: Ba(OH)2 is toxic. 7. 4.00 mL level of the buret. Prepare the computer for data collection by opening the file “Lab 26a: Conductivity to find Eqiv. For the filtration step.MEASURING VOLUME USING A BURET 1. Once the conductivity has stabilized. Record the precise H2SO4 concentration (given) in your data table. conductivity) for this experiment. 6. Obtain and wear goggles. type the current buret reading. asking the TA for help if you are unsure on how to do this. Using a filter flask with two pieces of filter paper. 5. clicking the Keep button. pt” from the Chemistry with computers folder of Logger Pro. You have now saved the second data pair for the experiment. and should be handled with care. Calculate the moles of Ba(OH)2 at the equivalence point. The precise volume of H2SO4 added can be confirmed by examining the data table for the minimum conductivity obtained. Vernier. and the other will be the linear region of data following the equivalence point. Dispose of the beaker contents in the 9. Drag your mouse cursor across the linear region of data that precedes the minimum conductivity reading. 10. M. Next. 2. 8. When you have passed the equivalence point. use 2-drop increments (~0. 1.00 mL. Record the volume of H2SO4. . Print a copy of the graph. in L. PROCESSING THE DATA 1. determine the volume of H2SO4 added at the equivalence point. Finally use 1. using the molarity. but recall that you must subtract the beginning volume if it wasn’t 0.1 mL) until the minimum conductivity has been reached at the equivalence point. . This volume reading will correspond to the equivalence point volume for the titration. Make sure each group member has the data. of the H2SO4 and its volume. Print a copy of the table. Third Edition. Record this volume and comment on whether it matches the one you selected manually in the Discussion section of your report. Click on the Linear Fit button. in molarity (mol/L). Use your answer in the previous step and the ratio of moles of Ba(OH)2 and H2SO4 in the balanced chemical equation (or the 2:2 ratio of moles – of H+ to moles of OH ). This lab was modified from lab 26 “Using Conductivity to find an Equivalence Point” from Chemistry with Computers. Click on the Linear Fit button. click waste jar as directed by your TA. Using Conductivity to Find an Equivalence Point: Name_____________________________ Date______________ Lab Section______________ 61 . Drag your mouse cursor across the linear region of data that follows minimum conductivity reading. EQUIVALENCE POINT DETERMINATION: An Additional Method An alternate way of determining the precise equivalence point of this titration is to perform two linear regressions on the data. Inc. . The equivalence point volume corresponds to the volume at the intersection of these two lines. The graph should give you the approximate volume at this point. 4. calculate the concentration of Ba(OH)2. Read and enter the volume after each 2-drop addition.0 mL increments (read and enter the volume at each increment) until the conductivity reaches about 2000 µS/cm. Choose Interpolate from the Analyze menu. Then move the mouse cursor to the volume reading when both linear fits display the same conductivity reading. 3. continue using 2-drop increments until the conductivity is greater than 100 µS/cm again. Calculate moles of H2SO4 added at the equivalence point. 2. 3. When you have finished collecting data. One of these will be on the linear region of data approaching the equivalence point. From the moles and volume (25 mL) of Ba(OH)2 used. e. From the data table and graph that you printed.d. Be sure to define equivalence point.Provide a brief statement of the purpose of this activity. 62 . Explain conductivity and the idea behind why conductivity can be used in determining the equivalence point in a titration. Show any relevant calculations.Data and Analysis: Report the molarity of sulfuric acid used. Report your determined concentration of Ba(OH)2 from both analysis methods (if they differ). Copy and paste your titration curve in the space below (include a descriptive caption). 63 . Discussion: Briefly discuss which of the two analysis methods you feel was the most accurate (be sure to support your answer with an explanation)? Why does the conductivity not go to zero? Using Conductivity to Find an Equivalence Point Lab Report: There is not a form al lab report for this lab. 64 . subm it the w orksheet on tim e and to your TA in the d ropbox on D2L. Com plete the above pages using the Microsoft version of this file that is available for d ow nload on the lab D2L page. Once the w orksheet is com plete. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. What wavelengths of the electromagnetic spectrum correspond to visible light? 2. What is the equation for the energy levels of the hydrogen atom? Give the units associated with the energy equation you report. 65 . remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. This includes stating the purpose of the experiment.Prelab: Atomic Emission Spectra Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. Part B Prepare your notebook for the lab. When light is emitted from the hydrogen atom. is the atom moving from a higher energy state to a lower energy state or a lower energy state to a higher energy state? 4. At the start of your lab. Why do atoms exhibit line spectra? 3. and calculate the energy of each line To gain an understanding of the quantized nature of the hydrogen atom Supplies Simple transmission grating spectroscope Hydrogen and other elemental emission lamps White. but not in between them. the emission spectrum of hydrogen consists of only four visible lines: red. A spectroscope is an instrument that allows you to analyze light in some way. As an example. 4. appears as one or more sharp narrow lines (colors). since individual’s visual acuity varies). A band spectrum is intermediate in appearance – there will be a brightest color. Each color corresponds to the transition of an electron from an excited state. including heating the atoms or using an electric discharge. red. bluegreen. Background The understanding of the internal structure of the atom was advanced when Niels Bohr explained the cause of the emission spectra of atoms using the concept of quantization. but a range of nearby wavelengths are emitted. when viewed through a spectroscope. and green light sources Definitions 1. violet and deep violet (although your eyes may not be sensitive to the last one. A continuous spectrum is one in which a rainbow of colors is seen when viewed through a spectroscope or prism. 3. After exciting the electrons. Electrons within an atom can be excited to higher energy states through various means. to a lower principal energy level. possibly the 66 . A line spectrum. This emitted light can be passed through a prism or reflected from a diffraction grating to spatially separate it into its individual wavelength components (colors) generating an atomic emission spectrum.Atomic Emission Spectra Activity Goals To view the hydrogen emission spectrum and other atomic line spectra To contrast the line spectra with other broad-band sources To measure the wavelengths of the bright lines in the visible emission spectra of hydrogen and mercury. the atoms emit electromagnetic radiation as the excited electrons relax into a lower energy state. like a spark. a line spectrum characteristic of the particular sample of atoms. a higher principal energy level. The energy of the light emitted is equal to the difference between the energy levels in the atom. In this case. splitting the light into different wavelengths. He stated that electrons in the atom could exist at finite energy levels. 2. your spectroscope acts like a prism. 178 x 10-18 J) (1/n2final .1/n2initial) [Rydberg equation. 67 . Other Equations That Might Be Useful E = (2. Recall that the energy of light is related to the wavelength and frequency of the light. In this experiment you will use a spectroscope and gas discharge lamps to measure the wavelength of each bright line in the visible atomic emission spectra of both hydrogen and mercury.00 x 108 m/sec. As the electron drops.ground state or some other allowed energy level in between. ninitial = 2] E = hc/ [Links wavelength and energy – be careful of the units!] 1 nm = 1 x 10-9 m The final results of this experiment will be the wavelength (in meters) and the photon energy of each bright line measured for hydrogen and mercury. for Balmer series.626 x10-34 J sec and the speed of light in a vacuum c = 3. through Planck’s constant h = 6. it emits energy in the form of a photon which may or may not be in the visible region. You will then use these measurements to calculate the photon energy for each bright line. Each minor graduation represents 10 nm. A very rough correction factor can be made using the difference between your measurement and the accepted value (436 nm). To determine the wavelength of the light you observe. Look through the slit at the small end pointing the widen end toward a fluorescent lamp (see figure 1. is as bright as possible. as seen through the slit. Any additional measurements made with this spectroscope should be corrected by subtracting 19 nm from the measurement. If for instance you measured the violet band at 455 nm. Figure 1 How a Spectroscope Works. You will see the light through a vertical slit on the left and the spectrum of the light source projected onto a wavelength scale on the right. you will make use of the calibrated wavelength scale which is in hundreds of nanometers. note the graduations on the back of the spectroscope. Figure 2: Observing a Spectra Using the Spectroscope.WARNING!! The power supply for the discharge lamps operates at 5000 volts ! DO NOT TOUCH How To Use The Spectroscope Hold the eyepiece of the spectroscope up to your eye. 68 . See Figure 2 for how to read the scale. This band should be at 436 nm. One band of light that you should observe is a violet band the far right of the spectra. First. with a transparent grating in one side and a narrow slit directly opposite the grating. To observe a spectrum you point the slit toward a light source and look through the grating. You will see an image of the spectrum along the back wall of the box. Adjust the position of the spectroscope until the light source. just over the wavelength scale. A spectroscope is a small box. You will be revisiting this spectra later. The large numbers on the scale are hundreds of nanometers. your spectroscope is miscalibrated by about 19 nm. note where that band appears on your spectroscope. above). Be sure to include a sample of your work for one of the calculations. If daylight is available and after giving time for your eyes to adjust. Comment. Qualitatively compare the 3 brightest lines from the emission spectra of the elemental mercury lamp to the spectra of a fluorescent lamp and an incandescent bulb and the red and green colored sources provided. whether or not the solar spectra a continuous spectra. Table 1: Bright Line Spectra for Elemental Hydrogen. observe the solar spectra by looking at the ambient daylight through a window (no worries if it is cloudy.Atomic Emission Spectra Activity Spectrum of a Single Electron Element – Hydrogen • Record the line color and its position for the 3 or 4 brightest lines observed using the hydrogen lamp in table 1 and calculate each wavelength using Equation 1 from your lab manual. what do you think the discontinuities represent? 69 . If not. Line Color Spectral Line Position (nm) Wavelength (nm) Corrected using calibration factor1 Spectrum of Multi-Electron Elements and Other Miscellaneous Spectra Qualitatively observe other atomic spectra and make notes below about the observed differences from the hydrogen spectrum. just aim at the sky). Never aim directly at the sun. Spectral Line Electronic Experimental λ (nm) Accepted λ (nm) from Color Observed Transition from Table 1 Table 2 Percent Error n3 → n2 n4 → n2 n5 → n2 n6 → n2 Below.the predicted wavelength of the emitted photon. ΔE -. clearly show your percent error calculation for the n3 → n2 transition. Electronic Transition ΔE (J) λphoton (nm) Ephoton (J) n3 → n2 n4 → n2 n5 → n2 n6 → n2 Clearly show the following calculations for the n3 → n2 transition. the predicted energy of the emitted photon (Ephoton) and the predicted wavelength of the emitted photon (λphoton).the change in energy of the electron. match the electronic transitions in the Balmer Series to the spectral lines you observed. Table 2: Calculated Values for the Balmer Series of Hydrogen. calculate the change in energy of the electron (ΔE). -. Then calculate the percent error between your experimentally determined and calculated wavelengths. Put the calculated values in Table 2 and be sure to clearly show an example of each calculation in the space provided. λphoton 2) Based on your theoretical calculations. Table 3: Comparison of Experimental and Accepted Wavelengths from the Balmer Series. 70 .Data Analysis 1) For the first four electronic transitions in the Balmer Series. and document your choices in Table 3. 3) It is not possible to observe the n7 → n2 transition in the Balmer Series. Why do you think that is? 4) Emission spectra are sometimes referred to as atomic fingerprints. Is it possible to use them to identify elements in an unknown sample? Explain your reasoning: think about the Hg spectrum and that of a fluorescent bulb. 5) Why do you think sodium vapor lights cast a different color (yellowish) than fluorescent lamps? 6) Calculate the ionization energy of the hydrogen atom. Think about this process as taking an electron from its ground state, n = 1, to a position/energy level far, far away from the nucleus, n = ∞. Atomic Emission Spectra Lab Report: There is not a form al lab report for this lab. Com plete the above pages using the Microsoft version of this file that is available for d ow nload on the lab D2L page. Once the w orksheet is com plete, subm it the w orksheet on tim e and to your TA in the d ropbox on D2L. 71 Pre-Lab: Determining the Concentration of a Solution: Beer’s Law Part A Answer the following questions in your lab notebook (be sure to show your work for any calculations): 1. You are given a colored solution that is labeled 1M . You need to prepare a solution from this that is 0.5 M. Describe your proced ure in d etail. 2. What is the relationship betw een absorbance and transm ittance? 3. Allura Red is a com m only used red food d ye. Does Allura Red transm it or absorb red light? 4. If 5.00 m L of a 0.5 M solution is d iluted to a final volum e of 100.0 m l, w hat is the concentration of the final d ilute solution? Part B Prepare your notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. At the start of your lab, remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. 72 Determining the Concentration of a Solution: Beer’s Law OBJECTIVES In this experiment, you will Prepare Allura Red standard solutions Use a Colorimeter to measure the absorbance value of each standard solution Find the relationship between absorbance and concentration of a solution Use the results of this experiment to determine the concentration of Allura Red in red Gatorade INTRODUCTION The primary objective of this experiment is to determine the concentration of Allura Red in a commercially available beverage. You will be using the Colorimeter shown in Figure 1. In this device, light from the LED light source will pass through the solution and strike a photocell. A higher concentration of the colored solution absorbs more light (and transmits less) than a solution of lower concentration. The Colorimeter monitors the light received by the photocell and reports either an absorbance or a percent transmittance value as compared to a blank, a solution containing no absorber. Figure 1 Figure 2 You are to prepare five Allura Red solutions of known concentration (standard solutions) and conduct a calibration procedure. Each standard solution is transferred to a small, rectangular cuvette that is placed into the Colorimeter. The amount of light that passes through the solution and strikes the photocell is used to compute the absorbance of each solution. When a calibration graph of absorbance vs. concentration is plotted for the standard solutions, a linear relationship should result, as shown in Figure 2. This linear relationship between absorbance and concentration for a solution is known as Beer’s law. The concentration of Allura Red in an unknown solution (Gatorade) is then determined by measuring its absorbance in the same way with the Colorimeter. By locating the absorbance of the unknown solution on the vertical axis of the graph, the corresponding concentration can be found on the horizontal axis (follow the arrows in Figure 2). The concentration of the unknown can also be found using the slope of the Beer’s law line, assuming that the y-intercept of the calibration line is zero. 73 Equipment Information Each bin should contain: 1 – 10mL serological pipet with bulb 5 – cuvettes with lids Do not aspirate liquid into the bulb. Notes: a. Today’s waste can go down the drain. b. Do not leave a cuvette in the colorimeter at the end of class. c. The colorimeter should be left out on the lab bench at the end of class. It should not be put away in a bin. 74 Rinse the cuvette twice with standard solution 2 (next highest concentration). Using standard solution 1 (the lowest concentration sample). Wait until Step 12 to do the unknown. “Blank” the Colorimeter. Connect the Colorimeter to the computer interface. 8. Be sure to record the concentration of the stock solution of Allura Red from the container. click . When the absorbance value stabilizes. type the concentration of the standard solution in the edit box. 5. and press the ENTER key. and fill the cuvette 3/4 full. click . Click “Keep” and then “Done”.000. Empty the water from the cuvette. place it in the Colorimeter. Prepare the computer for data collection by opening the file “Lab 11: Beer’s Law” from the Chemistry 227 folder of Logger Pro. After closing the lid. Click . remember: All cuvettes should be wiped clean and dry on the outside with a tissue. wait for the absorbance value displayed on the computer monitor to stabilize. All solutions should be free of bubbles. To correctly use a Colorimeter cuvette. 7. The data pair you just collected should now be plotted on the graph. 6.] You may need to click on the Autoscale button to rescale the graph as you go along. In Logger Pro click on “Experiment”. Then click and type the concentration of the standard solution into the edit box. Obtain about 30 mL of Allura Red stock solution in a 100 mL beaker. Prepare a blank by filling the cuvette 3/4 full with distilled water. and press the ENTER key. [NOTE: When entering values. 9. You will prepare five solutions of Allura Red varying in concentration from approximately 6 x 10-6 M to 2 x 10-5 M. Add about 30 mL of distilled water to another 100 mL beaker. Discard the cuvette contents in the waste jar as directed by your TA. Set the colorimeter to a wavelength of 470 nm. Place the cuvette with the blank (water) in the colorimeter. Handle cuvettes only by the top edge of the ribbed sides. Thoroughly mix each solution by inverting the stoppered flask ten times. Make the solutions by pipetting the correct quantity of the Allura Red stock solution into the volumetric flask and then filling to the line with distilled water. You should now see an absorbance reading of 0. 75 . Always position the cuvette with its reference mark facing toward the white reference mark at the top of the cuvette slot on the Colorimeter. You may wish to check your concentrations and calculations with your TA before making the solutions. 3. Wipe the outside with a tissue and place it in the Colorimeter. Wipe the outside. 4.PROCEDURE 1. In the popup window check the box next to “one point calibration”. When you have entered all of your standard solutions. then from the drop down menu select “calibrate” and “lab pro colorimeter”. You are now ready to calibrate the Colorimeter. rinse the cuvette twice with ~1 mL amounts and then fill it 3/4 full. Obtain and wear goggles 2. Click “Calibrate now” and enter “100” in the box provided (100% T). Be careful to avoid getting liquid above the fill line. Repeat the Step 8 procedure to save and plot the absorbance and concentration values of standard solutions 3-5. 2 x 10-6 can be entered in standard computer format as 2E-6. and close the lid. You are now ready to collect absorbance data for the five standard solutions. (Important: The reading in the meter is live. Allura Red has the following chemical structure: 76 . Determine the unknown concentration: With the linear regression curve still displayed on your graph. Use the calibration curve equation to determine the concentration of Allura Red in the diluted solution (solve for Concentration. A best-fit linear regression line will be shown for your five data points and your blank. or you may use the following method: 1. 2. move the cursor straight up the vertical cursor line until the tool bar is reached. 11. The Proportional fit has a y-intercept value equal to 0. therefore. Read the absorbance value displayed in the meter. choose Interpolate from the Analyze menu. This line should pass near or through the data points and the origin (0. 13. Calculate the concentration of Allura Red in the undiluted Gatorade. Wipe the outside of the cuvette. 3. . Print a graph of absorbance vs. this regression line will always pass through the origin of the graph. Discard the solutions in the waste jar as directed by your teacher.10. and then select Proportional. To keep the interpolated concentration value displayed. PROCESSING THE DATA You may use Microsoft Excel to plot the data and obtain a linear relationship between the data. click the Linear Fit button. [Note: Another option is to choose Curve Fit from the Analyze menu. Use the pipette to deliver 5 mL of the Gatorade to a clean volumetric flask. so it is not necessary to click to read the absorbance value. Enter your name(s) and the number of copies of the graph you want and print. place it into the Colorimeter. Move the cursor along the regression line until the absorbance value is approximately the same as the absorbance value you recorded in Step 12. The corresponding concentration value is the concentration of the unknown solution. Examine the graph of absorbance vs. Obtain a small amount of Gatorade in a small clean beaker.0) of the graph. A vertical cursor now appears on the graph. in mol/L. and close the lid. given Absorbance). concentration. To see if the curve represents a linear relationship between these two variables. Finish preparing your unknown by diluting the Gatorade to a total volume of 50 mL with distilled water and mix thoroughly. with a regression line and interpolated unknown concentration displayed. record its value.) When the displayed absorbance value stabilizes.] 12. concentration. The cursor’s concentration and absorbance coordinates are displayed in the floating box. Rinse the cuvette twice with the unknown solution and fill it about 3/4 full. Be sure to record the absorbance and concentration data pairs that are displayed in the table. determine the number of molecules of Allura Red you would consume if you drank one 20 ounce bottle of Gatorade (is molar mass necessary for this step?). Vernier. Use the molar mass to determine what mass of Allura Red you would consume if you drank one 20 ounce bottle of Gatorade. Inc.With the help of your TA. calculate the molar mass of Allura Red. Third Edition. 77 . This lab was modified from lab 11 “Determining the Concentration of a Solution: Beer’s Law” from Chemistry with Computers. Finally. Further questions: Answer the following questions in a separate section: 1. How many molecules of Allura Red would you consume if you drank one 20 ounce bottle of Gatorade? 2. What mass of Allura Red you would consume if you drank one 20 ounce bottle of Gatorade? If your value is greater than the mass of one 20 ounce bottle.Determining the Concentration of a Solution: Beer’s Law Lab Report: Your report for this lab should include the following sections: Abstract: Your abstract must be written individually and should include the concentration of Allura Red in Gatorade Introduction: Begin with a statement of the purpose of the experiment Provide any relevant background and key concepts and an explanation of the techniques used (calibration curve. Beer’s Law) Data: Include a table showing the concentrations of your standard solutions and their absorbance values Results: Include a copy of your calibration curve State the concentration of Allura Red in your diluted Gatorade solution and the undiluted Gatorade Attach hand written sample calculations to the back of your report Discussion: Discuss the experiment and any possible sources of error Explain why you set the colorimeter to a wavelength of 470 nm. 78 . Subm it you r report on tim e and to your TA in the d ropbox on D2L. be sure to recheck your calculations. The interactions between water molecules can be described as electrostatic or coulombic. given this electrostatic interaction. Read sections 9. What does the description polar bond refer to? 4. 5.6 and 10. 1. where areas of positive charge are attracted to areas of negative charge.5 in your text book. What are the three types of bonding? 3. Briefly describe the difference between a nonpolar covalent and polar covalent bond. submit the worksheet on time and to your TA in the dropbox on D2L. Complete the below pages using the Microsoft version of this file that is available for download on the lab D2L page. Provide a brief explanation as to why atoms may have different values of electronegativity. 2. What causes a low electronegativity? High? 6.Name_____________________________ Date______________ Lab Section______________ Electron Density Lab: There is not a formal lab report for this lab. 79 . Once the worksheet is complete. Draw a cartoon of how three water molecules might be arranged in space. Adjust the electronegativities of atoms A and B. What does represent? 2. Enter: Phet. Turn on (check) electron density in “Surface” options. you will investigate how an atoms' electronegativity value affects the types of bonds they produce. Investigate how the bond behaves when the atom's electronegativity is changed. 4. 1. What scenarios can bring about a higher electron density around a particular atom? 80 . Turn on (check) all “View” options.Simulation In this atomic-level simulation.compare to the electron density around an atom with a δ+? 6. Can an atom with a high electronegativity form a covalent bond? Describe under what circumstances this can occur.colorado. How does changing the electronegativity of the atoms affect the bond polarity? 5.represent? 3. Describe how a molecule’s polarity is related to it electron density? How does the electron density around an atom with a δ.edu into the browser of your computer select: Play with the Sims Chemistry Molecule Polarity Part A Select the two atom tab in the upper left hand corner of the simulator. What does the symbol δ+ or δ. Explain how the direction of the arrow in the bond dipole symbol ( ) relates to the electron density. Positive electrostatic potential (colored in shades of blue) corresponds to a repulsion of the proton. Describe how electron density is related to the electrostatic potential? In the red shaded regions.Electrostatic potential correlates with electron density. Negative electrostatic potential (colored in shades of red) corresponds to an attraction of the proton. what would cause the repulsion? 8. What if the molecule is nonpolar? 81 . what is responsible for a potential attraction to a proton? In the Blue shaded regions. the partial charges and the electrostatic potential on a molecule. What happens to a polar molecule when the electric field is turned on? Make sure to spin the molecule several times while making observations. 9. Electrostatic potential is basically a measure of how a proton would react when brought to different regions of a molecule. 7. It provides a useful way to quickly predict polarity of a molecule or a region in a larger molecule. Octane (a major component of gasoline) will not dissolve in water because it does not have a molecular dipole and is thus a nonpolar molecule. Provide a scenerio in which a molecule with two strong bond dipoles can have no molecular dipole at all? Explain your answer with a drawing showing individual bond dipoles and the overall molecular dipole. Being able to predict the polarity of a molecule is extremely important since many properties of molecules depend on whether they are polar or non-polar. Explain your answer with a drawing showing individual bond dipoles and the overall molecular dipole. Part C One property between molecules which is explored more in CH222 is the solubility of one substance in another. molecules with similar molecular dipoles will tend to interact favorably and mix. Provide a summary as to how the bond dipoles affect the molecular dipose. Turn on (check) all “view options”. Also. 1. describe how the geometric orientation of the bonded atoms affect the molecular dipole? 2. Investigate how the bond behaves when the electronegativity of the individual atoms is changed.Part B Select the three atoms tab in the upper left hand corner of the simulator. both are polar molecules and possess strong molecular dipoles). There is an adage that describes this ability. For instance. In addition to changing the electonegativities. 3. a polar molecule will mix well (dissolve) other polar molecules (ethanol readily dissolves in water. Predicting a molecule’s polarity is a 82 . Provide a scenerio in which a molecule may have a very large molecular dipole. orient the radial atoms relative to one another by dragging them with the mouse. "Like dissolves like". and then check it in the “Real Molecules” section of the simulation. 2. predict the molecule’s molecular geometry. Using VSEPR. Molecule Lewis Structure Molecular Geometry 3-d Geometry with Bond Polarities Polar or nonpolar? * Tetrahedral CH3F POLAR N2 BF3 CH2F2 C is the central atom HCN C is the central atom *Make a prediction.multi-step process that starts with drawing the Lewis structure. 83 . Using the molecular dipoles/polarity of BF3. 1. For the following molecules complete this step-by-step process. explain why BF3 does not mix with H2O? Subm it you r report on tim e and to your TA in the d ropbox on D2L. Individual bond polarities are finally used to predict the molecular polarity. What is the value of Hreaction (in kJ/mol phosphoric acid) if 50. 84 . remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA.60 M H3PO4? 5. Use 4.0 mL of 0.0 g.6 °C. What is the value of qreaction for the neutralization reaction described in number 2? 4. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. 3. Calculate the amount of heat energy gained by the solution (qsol). A neutralization reaction was carried out in a calorimeter. The change in temperature (∆T) of the solution was 5. How many moles of phosphoric acid are contained in 50. Cs. This includes stating the purpose of the experiment.18 J/(g•°C) as the specific heat. of the solution. The temperature of the solution rose from 20.CH228 LABS Pre-Lab: Enthalpy of Neutralization of Phosphoric Acid Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1.0 °C to 25.60 M H3PO4 was used in the reaction described in number 2? Part B Prepare your notebook for the lab. A neutralization reaction was carried out in a calorimeter.0 mL of 0.6 °C and the mass of the solution was 100. Is this reaction endothermic or exothermic? 2. At the start of your lab. Compare your calculated enthalpy of neutralization with the accepted value. Calculate the enthalpy. In this case. Phosphoric acid will be the limiting reactant in this experiment. If the temperature of the reaction is measured precisely. you may assume that the heat loss to the calorimeter and the surrounding air is negligible. acid-base reactions can be observed and measured thermodynamically. For purposes of this experiment. ΔH. the reaction is carried out in a calorimeter.The Enthalpy of Neutralization of Phosphoric Acid OBJECTIVES In this experiment. of neutralization of phosphoric acid. In this experiment. you will react phosphoric acid with sodium hydroxide. Pages 246-248 and 257-258 in your text will provide background information. and you will accordingly be determining the enthalpy. Figure 1 85 . Selecting a limiting reactant helps ensure that the temperature measurements and subsequent calculations are as precise as possible. of neutralization of the acid. you will Measure the temperature change of the reaction between solutions of sodium hydroxide and phosphoric acid. the enthalpy of neutralization of an acid by a base (or vice versa) can be determined. Calculate the enthalpy. In addition. of neutralization per ionizable hydrogen for phosphoric acid. ΔH. ΔH. You will use a Styrofoam cup nested in a beaker as a calorimeter. INTRODUCTION A great deal can be learned by conducting an acid-base reaction as a titration. as shown in Figure 1. Coffee cups are reusable.Equipment Information Each bin should contain: 1 – temperature probe 2 – coffee cups After use. secure cord as shown Notify TA if damaged Notes: a. Chemical Safety Information Enthalpy of neutralization of phosphoric acid Chemical Phosphoric acid Sodium hydroxide Hazards corrosive corrosive 86 . Do not throw them in the trash. CAUTION: Handle the phosphoric acid with care. Nest a Styrofoam cup in a 250 mL beaker as shown in Figure 1. Use a glass stirring rod to stir the reaction mixture gently and thoroughly. Click the Statistics button. If directed. d. It can cause painful burns if it comes in contact with the skin. After you have recorded three or four readings at the same temperature. and stirring rod. Conduct the experiment. Start the Logger Pro program on your computer. If the minimum temperature is not a suitable initial temperature. if the temperature readings are no longer changing. CAUTION: Sodium hydroxide solution is caustic. 87 . Lower the Temperature Probe into the phosphoric acid solution.0 mL of 0. 5. Data will be collected for 10 minutes. Rinse and dry the Temperature Probe. f. 3. The minimum and maximum temperatures are listed in the statistics box on the graph. examine the graph and determine the initial temperature.85 M NaOH solution in a graduated cylinder and transfer it to a 250 mL beaker. Styrofoam cup. conduct a third trial. Use a utility clamp to suspend the Temperature Probe from a ring stand (see Figure 1). 9.0 mL of 1. e.PROCEDURE 1. Measure out 50. Connect a Temperature Probe to Channel 1 of the Vernier computer interface. You may terminate the trial early by clicking . a.0 mL of NaOH solution to the Styrofoam cup all at once. Open the file “Lab 1 Phosphoric” from the Chemistry 228 folder. Close the Statistics box by clicking the X in the corner of the box. add the 50. Click to begin the data collection and obtain the initial temperature of the H3PO4 solution. 2. It is best to conduct this experiment in a well-ventilated room. Print a copy of the graph of the second trial to include with your data and analysis. Avoid spilling it on your skin or clothing. b. Record the initial and maximum temperatures for Trial 1. Measure out 50. Dispose of the solution as directed. 6. . c. Repeat Steps 4–8 to conduct a second trial. 8.60 M H3PO4 solution into the foam cup. 4. Obtain and wear goggles. 7. Determine the number of moles of phosphoric acid used in the reaction. Use 4.DATA TABLE Trial 1 Trial 2 Trial 3 Maximum temperature (°C) Initial temperature (°C) Temperature change (∆T) DATA ANALYSIS 6. qsol = Cs m ∆T 8.44 kJ/mol. The change in temperature (∆T) is a directional change where ∆T = Tf –Ti. Calculate the percent error in your experimental value. ∆H = qrxn/moles H3PO4 10. of the solution use 1. Use the equation below to calculate the amount of heat energy gained by the solution (qsol). m. If the solution gained heat. Since we are interested in the heat of neutralization of phosphoric acid we need the heat transfer associated with the reaction (qrxn). ∆H. The accepted value for the ∆H of neutralization for phosphoric acid is -156. Cs. 88 . Write the balanced equation for the reaction of phosphoric acid and sodium hydroxide. of the solution. for the reaction in terms of kJ/mol of phosphoric acid. This is your experimental value of ∆H.18 J/(g•°C) as the specific heat.00 g/mL for the density (be sure to use the total volume of the solution after the acid and base are mixed). Use the moles of phosphoric acid along with qrxn to determine the enthalpy change. In determining the mass. 7. This relationship can be expressed by the following equation: qsol = -qrxn 9. the reaction must have given off heat. The heat calculated above represents the heat gained by the solution (the solution being predominantly water). Why can you use the specific heat capacity and density of pure water to determine the enthalpy of reaction? What assumptions must be made in order to do this? .The Enthalpy of Neutralization of Phosphoric Acid Lab Report Name_____________________________ Date______________ Lab Section______________ Describe coffee cup calorimetry and how it is used to find the enthalpy of various reactions that occur in aqueous solutions. Data: Insert your data table (with a caption) below that captures all of the relevant information. Make sure to include the relevant equations. 89 . Results: Report your calculated average value of ∆H of neutralization for phosphoric acid.. Com plete the above pages and su bm it them to your TA. The Enthalpy of Neutralization Lab Report: There is not a form al lab report for this lab. Think of some valid sources of error to account for your difference and explain how they would contribute to the direction of your error. Include hand written sample calculations. Discussion: Is your value for the ∆H of neutralization for phosphoric acid greater than or less than the accepted value. Include any calculations. Report the percent error for the ∆H of neutralization for phosphoric acid. 90 . The enthalpy of the reaction for the reaction of calcium oxide with hydrochloric acid is exothermic. When you measure a temperature rise during a chemical reaction. How will you use the measurement in the Hess’s Law calculation? 5. the second reaction is difficult to measure as written. At the start of your lab. What is the formula that relates the temperature change observed in a substance with the energy released or absorbed? 2. For the reactions described in the lab. 91 . Hess’s Law allows us to combine reactions to determine the heat of reaction for a net reaction that has not been measured. You will measure the heat of reaction for the reverse reaction. Will the reverse reaction have a positive or negative H? 4. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. This includes stating the purpose of the experiment.Pre-Lab: Hess’s Law Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. is the reaction endothermic or exothermic? 3. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA.8 kJ Part B Prepare your notebook for the lab. Calculate the enthalpy for this reaction: 2C(s) + H2(g) ---> C2H2(g) ΔH° = ??? kJ Given the following thermochemical equations: C2H2(g) + (5/2)O2(g) ---> 2CO2(g) + H2O(ℓ) ΔH° = -1299.5 kJ H2(g) + (1/2)O2(g) ---> H2O(ℓ) ΔH° = -285.5 kJ C(s) + O2(g) ---> CO2(g) ΔH° = -393. Consid er the follow ing three reactions: 1) Mg (s) + 2 H +(aq) ---> Mg 2+(aq) + H 2(g) H °1 2) Mg 2+(aq) + H 2O (l) ---> MgO (s) + 2 H +(aq) H °2 3) H 2(g) + 1/ 2 O 2(g) ---> H 2O (l) H °3 4) Mg (s) + 1/ 2 O 2(g) ---> MgO (s) H °4 You w ill d eterm ine the heat of reaction for reactions 1 and 2 experim entally. Be aw are that equation (2) is the reverse of the reaction you actually run and m easure. (N ote: the enthalpy of form ation of MgO cannot easily be m easured . accord ing to H ess's Law (see your textbook for more d etails).9 kJ/ m ol) to calculate H °4 w hich is the enthalpy of form ation of MgO . Rem em ber.Enthalpy of Reaction and Hess's Law Introd uction In this experim ent you w ill be find ing the enthalpy of form ation for MgO (s) using an ind irect m ethod . H ° for the net reaction is sim ply the sum of the H °''s for the reactions w hich are ad d ed . if tw o or m ore reactions can be ad d ed to give a net reaction. then use the know n value of the enthalpy of form ation of w ater (H °3 = -285.) 92 . Do not throw them in the trash.Equipment Information Each bin should contain: 1 – temperature probe After use. Chemical Safety Information Hess's Law Chemical Hydrochloric acid Magnesium oxide Magnesium Hazards corrosive none flammable 93 . secure cord as shown 2 – coffee cups Notify TA if damaged Notes: a. Coffee cups are reusable. Make sure the calorimeter is clean and mostly dry before repeating the experiment. e. a. 94 . 7. Use a utility clamp to suspend the Temperature Probe from a ring stand (see Figure 1). and stirring rod. 5. Start the Logger Pro program on your computer. . Rinse and dry the Temperature Probe. Hydrogen gas is flammable. Conduct the experiment. You may terminate the trial early by clicking .0 M HCl into the calorimeter. Obtain a coffee cup calorimeter from the stockroom. Data will be collected for 10 minutes. f. mL of 1. HCl is a strong acid. Styrofoam cup. Do not use any open flames in the lab. Nest a Styrofoam cup in a 250 mL beaker and put 50. d.0. Connect a Temperature Probe to Channel 1 of the Vernier computer interface.Experimental Procedure Part A 1. Lower the Temperature Probe into the acid solution. 6. weigh out a sample containing between 0. Record the initial and maximum temperatures for Trial 1. if the temperature readings are no longer changing. Close the Statistics box by clicking the X in the corner of the box. Click the Statistics button. Make sure the cup is clean and dry. Use a glass stirring rod to stir the reaction mixture gently and thoroughly.55 grams of magnesium. 4. Conduct another trial as above. examine the graph and determine the initial temperature. Using a weighing boat. Click to begin the data collection and obtain the initial temperature of the acid solution. After you have recorded three or four readings at the same temperature. 3. c. add the magnesium to the styrofoam cup all at once. Dispose of the solution as directed. 2. If the minimum temperature is not a suitable initial temperature. The minimum and maximum temperatures are listed in the statistics box on the graph. Caution: Wear your goggles at all times. Open the file “Lab 1 Phosphoric” from the Chemistry 228 folder.45 . b. because you are using the specific heat of pure water. Report ΔH°rxn for reaction 1 and 2. be certain to use units of kJ/mol. dry calorimeter.) You should use a molar equivalent of MgO (24. not the combined mass of water and solute.184 J/g-°C. cs = 4. this time replacing Mg with MgO.3 g Mg is the molar equivalent of 40. why?. this will require vigorous stirring!! Conduct another trial as above. it is a good approximation to take specific heat of the solution to be the specific heat capacity of water. For mass. your measurement should be within 5%) Be certain all the MgO dissolves. (Use a clean. use the mass of the water only. Calculations To relate heats of reactions (in energy units of Joules) with temperature differences we use: q = m x cs x ΔT For the reactions above.Part B Repeat the above procedure. Calculate ΔH°4 95 . Calculate q for reaction 1 and 2.3 g MgO. you m easure the quantity. For an exothermic reaction.Hess’s Law Lab Report: Your report for this lab should include the following sections: Abstract: Your abstract should be written individually Introduction: Include a statement of purpose for this experiment. answer the following questions: 1. need s to analyzed . relevant conceptual background. q. Is reaction 1 end otherm ic or exotherm ic? reaction 2? 4. a) Using w hat you have learned about enthalpy. For an exothermic reaction. does the temperature observed rise or fall? 2. is H° positive or negative? 3. 96 . (rem em ber the d efinitions and units)? Answ er the follow ing question and attach it to your report: Further Analysis An alloy (a m etal m ixture) containing m agnesium and another m etal. b) If the sam ple w ere 30% m agnesium calculate the heat evolved if a 5 gram sam ple w ere analyzed in that m anner. that d oes not react w ith hyd rochloric acid . d escribe the proced ure you w ould use to d eterm ine the percent m agnesium in the alloy. Subm it you r report on tim e and to your TA in the d ropbox on D2L. and general equations Data: Prepare a data table that includes the initial and final temperatures for each trial Report the mass of Mg and MgO used in each trial Results: Prepare a results table showing the calculated H rxn for each trial and averages for reactions one and tw o and the value of H for reaction four Be sure to attach hand written sample calculations to the back of your report Discussion: Discuss the experiment and any possible sources of error As part of your discussion. In this lab. H ow is this d ifferent from H °rxn . You are asked to determ ine the percent m agnesium in the alloy. Use your value from this experim ent for the H rxn for Mg in ord er to obtain a value. one can derive the mathematical relationships that exist between these variables by holding two of the variables constant. We are constantly being exposed to the behavior of gases. pressure (P). 3. V. we are reminded of how gases behave with changes in temperature (T). In a scientific manner. gases are in constant and random motion with enough kinetic energy such that they rarely interact with one another. When gas particles collide with the walls of a container. you can always hit the reset button at the bottom right of the screen.edu/new/simulations/sims. use a spray can. Go to the Physics Education Technology from the University of Colorado at: http://phet. temperature (T) and numbers of particles of gas (n). These characteristics describe an "Ideal Gas. Together these studies led to the so called "Gas Laws" which relate volume (V). To derive the relationships. 4. click on the MEASURMENT TOOLS button. Each time we pump up a tire.php?sim=Gas_Properties 2. Qualitatively get a feel for the relationships that exist between the four variables that describe gases: P. or experience the cooling of gases as they escape from a gas storage container. n and T.Deriving the Gas Laws Using Computer Simulations Note: If you have a M acintosh computer or are having trouble running the simulators on your personal computer. 97 . blow up a balloon. pressure (P).colorado. on the lower right side of the screen. On the right side of the screen." Experimental evidence suggests that many common gases making up air behave in this manner when studied at temperatures well above their boiling points. Play around with the simulator and see what sorts of tools are available to you to analyze the behaviors of gases. If you have not already done so. PROCEDURE 1: Pressure Volume Relationship 1. followed by Jacques Charles' (1787) and Joseph Gay-Lussac's work (1802). volume (V). changing one and monitoring the effect on the fourth variable. If you ever get to a point that you need to reset the simulator. you will be using an interactive research-based simulation produced by the PhET project at the University of Colorado. they rebound with no apparent loss of energy. click on the RULER option to activate the ruler. click on the RESET button. Next. Introd uction According to the kinetic molecular theory. Click the RUN NOW button under the Gas Properties Simulation window (highlighted in green). please use any of the university computer labs. 5. Chemistry students are welcome to use the computers in the chemistry department sponsored computer lab located in SB1 room 221. or number of particles (n). The behavior of gases has been scientifically investigated starting with Robert Boyle's work in 1662. 0 nm) GRAB AND DRAG 98 . 8. CLICK HERE 7. drag the ruler into a position that will allow you to measure the length of the container.CLICK HERE 6.0 nm and the width of the box will remain 5. click on the TEMPERATURE button under the Constant Parameter heading. grab hold of the man pushing against the container and expand the length of the container so that it measures 9. In the upper right hand corner.0 nm. Record this as your initial length (the height of the box will remain 5. Using the mouse and the right button. Using the mouse and the right button. This will hold temperature constant while allowing you to observe the relationship between pressure and volume. grab hold of the pump handle and inject one cycles worth of gas into the chamber by pulling the handle up then pushing it back down. MOVE UP THEN DOWN 10.0 nm. record your pressure value for the chamber length of 9.0 nm. PRESSURE (ATM) 11. record the new pressure for a length of 8.9. PUSH IN 99 . Once the pressure has stabilized (again. Using the mouse and right button. This will represent your initial pressure in atmospheres.0 nm. this may take a short period of time to happen). Using the mouse and the right button. Once the pressure has somewhat stabilized. grab hold of the man pushing on the container and decrease the length of the container to approximately 8. 0 nm. before adding any gas. you will need to submit neat labeled data tables for each procedure. Add gas (more particles seem to provide less noise in the volume measurement). 0-600 K.0 nm. 3. Graphically represent the Pressure (atm) and Inverse Volume. Along with the graphs and tables for each procedure. 1/V (nm-3) relationship with 1/V on the x axis 3. 4. Be sure to label each axis and include a title for each graph (Please see the information on pages 15 and 16 of this lab manual). answer completely the questions below that correlate with each section. and 2. Record any qualitative observations on the behavior of the gas molecules as the volume decreases. 2. 5. 4. Procedure 3: Temperature Pressure Relationship Devise an experiment using the simulator in which you can elucidate the relationship between Temperature and the Pressure of a gas. 14. record the length value and resulting pressure value in a properly labeled data table. reduce the volume of the container.0 nm (you will probably not get to exactly 2 nm). I suggest that you utilize Microsoft Excel or some other comparable spreadsheet software to produce your tables and graphs. Hint. 13. Collect and record your data over a wide range of number of molecules in a properly labeled table Analysis: For this lab. remove constant volume and change to constant pressure before beginning your measurements.0 nm. Repeat step 9 for approximate lengths of 7. Collect and record your data over a wide range of temperatures. in a properly labeled table Procedure 4: Pressure Quantity Relationship Devise an experiment using the simulator in which you can elucidate the relationship between Quantity and the Pressure of a gas. Finally. Collect and record your data over a wide range of temperatures. as being directly proportional or inversely proportional. Why were you asked to graph pressure and the inverse of volume? 100 .0 nm. Graphically represent the Pressure (atm) Volume (nm3) relationship with volume on the x-axis. not the mathematical equation from the best fit line. when temperature and quantity are held constant. Explain your answer and write an equation that relates pressure and volume to a constant.12. PROCEDURE 2: Volume Temperature Relationship Devise an experiment using the simulator in which you can elucidate the relationship between Temperature and the Volume of a gas. 6. Analysis: Procedure 1: Pressure Volume Relationship 1.0 nm. Click the RESET button to remove all the gas particles from the chamber before moving on to the next section. You must also submit a graphical representation for each relationship. using variables. Reduce the temperature to a good low starting point. For each trial. Identify the mathematical relationship that exists between pressure and volume. Make the volume constant. 10. 8. using variables. Graphically represent the Temperature (K) Volume (nm3) relationship. Make sure the axis that represents temperature includes a range from 0 K to 600 K. What effect does temperature have on molecular motion. as being directly proportional or inversely proportional. Subm it you r report on tim e and to your TA in the d ropbox on D2L. Would you expect this trend to be the same for other gases? Explain your answer. pressure graph. What properties does this number represent? Would you expect it to be the same for other gases? Explain your answer. predict the impact of changing the number of moles of a gas sample on the volume of the gas sample (if pressure and temperature are held constant). explain why both pressure and volume can decrease with decreasing temperature. volume graph. Using this explanation. Describe the impact of increasing the number of molecules (or moles) of a gas on the pressure of a gas sample. Analysis Questions: Procedure 2: Volume Temperature Relationship 6. 16. Graphically represent the Temperature (K) Pressure (atm) relationship. what would you predict to be the pressure and volume of a gas sample whose temperature is decreased to absolute zero? Explain any problems with these expectations using the ideal gas law.5. Graphically represent the Quantity (number of molecules) Pressure (atm) relationship 15. when volume and quantity are held constant. Analysis Questions: Procedure 3: Temperature Pressure Relationship 9. when pressure and quantity are held constant. 11. Explain your answer and write an equation that relates pressure and temperature to a constant. Absolute zero is theorized to be the temperature that all molecular motion stops. not the mathematical equation from the best fit line. Analysis Questions: Procedure 4: Pressure Quantity Relationship 14. What properties does this number represent? Would you expect it to be the same for other gases? Explain your answer. Based on your previous observations. What effect would changing the number of moles of a gas sample have on the temperature of a gas sample (if pressure and volume are held constant)? Explain your answer and state whether these relationships are proportional or inversely proportional. 13. Identify the mathematical relationship that exists between pressure and temperature. Explain your answer and write an equation that relates volume and temperature to a constant. using variables. Identify the mathematical relationship that exists between volume and temperature. Based on this. 1/volume graph. 7. 12. not the mathematical equation from the best fit line. Calculate the slope of the line for your temperature vs. What properties does this number represent? Would you expect it to be the same for other gases? Explain your answer. Calculate the slope of the line for your temperature vs. Calculate the slope of the line for your pressure vs. as being directly proportional or inversely proportional. 101 . 4. If 0.00946 moles of O2 gas is collected from the decomposition of hydrogen peroxide. What is Dalton’s law of partial pressure? 2. a pressure of 743 torr and a temperature of 20 °C. how many moles of hydrogen peroxide were reacted? Part B Prepare your notebook for the lab.Pre-Lab: Decomposition of Hydrogen Peroxide Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. B and C) has a total pressure of 849 torr and the partial pressure of A is 57 torr and the partial pressure of B is 573 torr. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. What is the partial pressure of C? 3. This includes stating the purpose of the experiment. 102 . A mixture of three gasses (A. A gas has a volume of 94 mL. Assume ideal behavior of the gas. Calculate the number of moles of gas present. At the start of your lab. The catalyst (KI) is located in the syringe and can be added to the Erlenmyer flask to initiate the reaction.Decomposition of Hydrogen Peroxide OBJECTIVES Decompose hydrogen peroxide using KI as a catalyst Measure the volume of oxygen gas generated through the decomposition reaction Illustrate Dalton’s Law of partial pressure Determine the number of moles of oxygen gas produced using the ideal gas law Determine the percent hydrogen peroxide in an aqueous solution INTRODUCTION Hydrogen peroxide spontaneously decomposes to form oxygen gas according to the following equation: 2 H2O2 (aq) → 2 H2O (l) + O2 (g) This process usually occurs very slowly. Hydrogen peroxide will be placed in the Erlenmyer flask. Figure 1 103 . The gas will be collected in the graduated cylinder. As the reaction proceeds oxygen gas will be produced in the Erlenmyer flask and travel through the tubing. Here. The graduated cylinder is initially filled with water. The apparatus we will use to collect oxygen gas in this experiment is shown in figure 1. potassium iodide (KI) will be used as a catalyst to make the reaction produce products rapidly enough to study the reaction in the lab. Many different compounds or ions are capable of acting as catalysts increasing the rate of the reaction. As the gas enters the cylinder it displaces water allowing the volume of the gas to be measured. Parafilm is reusable. b. 2x plastic Luer tips. do not put temperature probe in gas capture kit bag. When cleaning up. Chemical Safety Information Decomposition of hydrogen peroxide Chemical Hydrogen peroxide Potassium iodide Hazards corrosive toxic 104 . 1x tubing. and 1x stopcock Notes: a. Water from the water displacement bath is not waste. 1x syringe. secure cord as shown 1 – gas capture kit See below for components Place all components back into bag at end of class Gas capture kit components: 1x rubber stopper.Equipment Information Each bin should contain one 100mL graduated cylinder and: 1 – temperature probe After use. c. The gas collected is therefore a mixture of both oxygen and water. If there is a small amount of air present in the cylinder record the volume. 9. Place the rubber stopper tightly in the flask (this should be air tight). Cover the cylinder with parafilm and invert the cylinder in the 800 mL beaker. Obtain a ring stand and clamp the flask as shown in figure 1. To determine the pressure of oxygen gas we must apply Dalton’s law of partial pressure. Completely fill a 100 mL graduated cylinder with water. you will need to redo the setup. The graduated cylinder should be completely filled with water. Allow the reaction to proceed until no further production of oxygen gas is observed (around 10 to 15 minutes). Record the actual mass used. Repeat the above procedure two more times for a total of three trials. Record the final level of the water in the graduated cylinder.PROCEEDURE 1. Attach the syringe to the adaptor in the rubber stopper. Add approximately 5 g of hydrogen peroxide solution into the Erlenmyer flask. Remove the parafilm and carefully place the end of the tubing just inside the graduated cylinder as shown in figure 1. DATA ANALYSIS 1. Measure and record the temperature of the water. you can find the pressure of oxygen gas by subtracting the partial pressure of water at the temperature of the water (also known as the vapor pressure of water) from the total pressure (or atmospheric pressure). A table of the vapor pressure of water at various temperatures follows. If there is more than 10 mL of air in the cylinder. PTot = pO2 + pH2O In other words. Add your magnetic stir bar. Your TA will provide the current barometric pressure. Place on the magnetic stirrer on and turn on to a low setting 5. The total pressure of the gas is the sum of the pressures exerted by the oxygen gas and water vapor. 3. In a small beaker obtain a small amount (approximately 10 mL) of 0. 7. 6. Place an 125 mL Erlenmyer flask on a balance and tare the scale. Place approximately 600 mL of water in an 800 mL beaker. At least two of your trials should agree well with one another.5 M KI. Draw up 3 mL of the KI solution into the syringe. 2. Be sure to record your measurement to 2 decimal places. Carefully clamp the cylinder in place such that the opening of the cylinder is below the surface of the water in the beaker. 105 . Initiate the reaction by depressing the stopper on the syringe and adding the KI to the hydrogen peroxide. 4. 8. Determine the pressure of the oxygen gas: Because the oxygen gas was collected over water some of the gas collected is water vapor. 5 15. VO2 = Vfinal – Vinitial – 3 mL 3.3 30 31.7 28. calculate the mass percent H2O2 in the initial solution. the moles of gas can be calculated using the ideal gas law. PV = nRT 4. Calculate the molar mass of H2O2 and determine the grams of H2O2 present in the initial solution. 3 mL of KI solution was added. 106 . This volume needs to be subtracted from the volume of gas collected. this must also be taken into consideration. If your initial volume of gas was not zero.4 23.8 25.8 13. Determine the volume of the oxygen gas: When the reaction was initiated.8 21.8 33.2 Temperature o ( C) 27 28 29 30 31 32 Vapor Pressure (torr) 26.6 19. calculate the number of moles of H2O2 present in the initial solution. The temperature of the gas will be considered to be the same temperature as the water temperature measured during the experiment.7 35.6 14.5 Temperature o ( C) 21 22 23 24 25 26 Vapor Pressure (torr) 18. Calculate the mass percent of hydrogen peroxide: Using the mass of H2O2 calculated above and the initial mass of the H2O2 solution.Table 1: Vapor pressure of water at various temperatures Temperature o ( C) 15 16 17 18 19 20 Vapor Pressure (torr) 12. 5. Calculate the amount of H2O2: Using the balanced equation.1 22. Calculate the number of moles of oxygen gas generated: Now that the pressure.5 16.5 17.7 2. volume and temperature of the gas are known. 107 . relevant conceptual background.Hydrogen Peroxide Lab Report: Your report for this lab should include the following sections: Abstract: Your abstract should be written individually Introduction: Include a statement of purpose for this experiment. and general equations Data: Include a data table with data from all 3 trials Results: Include a results table with the mass percent of hydrogen peroxide from each trial Be sure to attach hand written sample calculations to the back of your report Discussion: Discuss the experiment and any possible sources of error Subm it you r report on tim e and to your TA in the d ropbox on D2L. 7 kJ/mol. At 25 °C the pressure is 693 mmHg. what must be held constant? 2. Would you expect most of the components in a perfume to have a low or high vapor pressure? Explain. how much energy is required to vaporize 5. At the start of your lab. A sample of gas is held in a capped flask. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. This includes stating the purpose of the experiment. 108 .0 g of liquid water at 100 °C? 4. When using the equation P1/T1 = P2/T2 to relate temperature and pressure of a gas. Part B Prepare your notebook for the lab. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection.Pre-Lab: Vapor Pressure and Heat of Vaporization Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. If the heat of vaporization of water is 40. What is the pressure of the gas at 37 °C? 3. The pressure at equilibrium is called vapor pressure. ln P H vap 1 C R T where ln P is the natural logarithm of the vapor pressure. you will introduce a specific volume of a volatile liquid into a closed vessel. At this equilibrium. but it relates these factors to the heat of vaporization of a liquid. Calculate the heat of vaporization of the liquid. T is the temperature (in Kelvin) and C is a constant not related to heat capacity. Thus. you will be able to calculate the ΔHvap of the liquid.31 J/mol•K). Figure 1 109 . ΔHvap is the heat of vaporization.Vapor Pressure and Heat of Vaporization When a volatile liquid is placed in a container. R is the universal gas constant (8. The newly formed gas molecules exert pressure in the container. then at some point a physical equilibrium will be reached. a portion of the liquid will evaporate. In this experiment. and will remain constant as long as the temperature in the container does not change. the rate of condensation is equal to the rate of evaporation. In mathematical terms. If the temperature inside the container is held constant. OBJECTIVES In this experiment. you will Measure the pressure inside a sealed vessel containing a volatile liquid over a range of temperatures. the relationship between the vapor pressure of a liquid and temperature is described in the Clausius-Claypeyron equation. Determine the relationship between pressure and temperature of the volatile liquid. and the container is sealed tightly. the Clausius-Clayperon equation not only describes how vapor pressure is affected by temperature. By analyzing your measurements. while some of the gas condenses back into the liquid state. and measure the pressure in the vessel at several different temperatures. ΔHvap is the amount of energy required to cause the vaporization of one mole of liquid at constant pressure. secure cord as shown 1 – gas capture kit See below for components Place all components back into bag at end of class1x syringe. and 1x stopcock Notes: a. 1x tubing. toxic 110 . 2x plastic Luer tips. When cleaning up. do not put temperature probe in gas capture kit bag Chemical Safety Information Vapor pressure and heat of vaporization Chemical Ethanol Hazards flammable.Equipment Information Each bin should contain: 1 – gas pressure sensor 1 – temperature probe Place back in bag at the end of class After use. Gas capture kit components: 1x rubber stopper. Use a hot plate to heat ~200 mL of water in a 400 mL beaker. d. 9. 2. Place the Erlenmeyer flask in the water bath. After 30 seconds. Use the syringe to transfer the hot water. 13. Record these values in your notebook.) Twist the white stopper snugly into the neck of the Erlenmeyer flask to avoid losing any of the gas that will be produced as the liquid evaporates (see Figure 1). Obtain and wear goggles. Stir the water bath slowly with the Temperature Probe. c. Carefully remove the syringe from the stopper. 11. to accelerate the evaporation of the ethanol. Monitor the pressure and temperature readings. After 5 seconds. Thread the syringe onto the valve on the white stopper (see Figure 1).PROCEDURE 1. 12. Notify your teacher immediately if an accident occurs. When the readings stabilize. Hold the flask down into the water bath to the bottom of the white stopper. Prepare a room temperature water bath in an 800 mL beaker. Click to begin data collection. Start the Logger Pro program on your computer. b. When the pressure and temperature readings stabilize. click . using a motion similar to slowly stirring a cup of coffee or tea. add a small amount of hot water. d. Your first measurement will be of the pressure of the air in the flask and the room temperature. 8. b. Condition the Erlenmeyer flask and the sensors to the water bath. 10. Record these values in your notebook. Connect a Temperature Probe to Channel 2 of the interface. 111 . Add ethanol to the flask. to warm the water bath by 3–5°C. While gently rotating the flask in the water bath. Important: Open the valve on the white stopper. Avoid inhaling the vapors. (About one-half turn of the fittings will secure the tubing tightly. Place the Temperature Probe in the room temperature water bath. Draw 3 mL of ethanol into the 20 mL syringe that is part of the Gas Pressure Sensor accessories. from the beaker on the hot plate. Connect a Gas Pressure Sensor to Channel 1 of the Vernier computer interface. Open the valve below the syringe containing the 3 mL of ethanol. Be sure that there are no open flames in the room during the experiment. 4. click . a. Use the clear tubing to connect the white rubber stopper to the Gas Pressure Sensor. f. Avoid contact with your skin or clothing. CAUTION: The alcohol used in this experiment is flammable and poisonous. close the valve on the white stopper. Push down on the plunger of the syringe to inject the ethanol. e. a. The bath should be deep enough to completely cover the gas level in the 125 mL Erlenmeyer flask. Obtain a small amount of ethanol. record these values. Open the file “Lab 4 Vapor Pressure” from the Chemistry 228 folder. 6. When the readings stabilize. Monitor and collect temperature and pressure data. close the valve on the white stopper. 3. Gently rotate the flask in the water bath for a 1 minute. c. 5. Dispose of the ethanol in the liquid waste receptacle and the water from the bath down the sink. as Ptotal. but do not warm the water bath beyond 40°C because the pressure increase may pop the stopper out of the flask. 15. Remove the flask from the water bath and take the stopper off the flask.14. Add enough hot water for each trial so that the temperature of the water bath increases by 3-5°C. open the valve to release the pressure in the flask. and the temperature readings in your data table. 17. Click to end the data collection. do it carefully so as not to disturb the flask. 16. Record the pressure readings. If you must remove some of the water in the bath. Do not exit the Logger Pro program until you have completed 1–4 of the Data Analysis section. After you have recorded the fifth set of readings. DATA TABLE Initial Trial 1 Trial 2 Trial 3 Trial 4 Trial 5 Ptotal (kPa) Pair (kPa) Pvap (kPa) Temperature (°C) 112 . Repeat Step 13 until you have completed five total trials. f. a. Choose New from the File menu. and then select ln vapor pressure to plot on the y-axis. Com plete the below pages and subm it them to your TA. 3. d. Use the gas law relationship shown below to complete the calculations. Prepare and print a second graph. Double-click on the y-axis heading in the table. enter a name and unit. Create a column ln vapor pressure. As you warmed the flask. d. reciprocal of absolute temperature.DATA ANALYSIS 1. b. Create a second column. enter a name and unit.32 kJ/mol. then enter the five values for vapor pressure from your data table above. b. a. then enter the five values for temperature (°C) from your data table above. Calculate the linear regression (best-fit line) equation for this graph. the air in the flask exerted pressure that you must calculate. Double-click on the x-axis heading in the table. Prepare and print a graph of Pvap (y-axis) vs. The Pair for Trials 2-5 must be calculated because the temperatures were increased. An empty graph and table will be created in Logger Pro. the reciprocal of absolute temperature. ΔHvap. Disconnect your Gas Pressure Sensor and Temperature Probe from the interface. and reciprocal of absolute temperature to plot on the x-axis. There is not a form al lab report for this lab. c. if necessary. c. Use the Pair from before the volatile liquid was added to the flask as P1 and the Kelvin temperature of Trial 1 as T1. Calculate and record the Pvap for each trial by subtracting Pair from Ptotal. Does the plot follow the expected trend of the effect of temperature on vapor pressure? Explain. Autoscale the graph. e. 4. Celsius temperature (x-axis). P1 P2 T1 T2 2. you will first need to plot the natural log of Pvap vs. Remember that all gas law calculations require Kelvin temperature. In order to determine the heat of vaporization. click on the respective axes. Calculate ΔHvap from the slope of the linear regression. On the displayed graph. The accepted value of the ΔHvap of ethanol is 42. 113 . Choose New Calculated Column from the Data menu. Compare your experimentally determined value of ΔHvap with the accepted value. 1/(Temperature (°C) + 273). e. 5. Additionally show the Clausius-Clapeyron equation and describe how vapor pressure and temperature data can be manipulated to find the enthalpy of vaporization. Be sure to include the meaning of enthalpy of vaporization and why vapor pressure is temperature dependent.Vapor Pressure and Heat of Vaporization Lab Report: Name_____________________________ Date______________ Lab Section______________ Provide a brief statement of the purpose of this activity. 114 . Include a copy of your graph of ln Pvap vs. . Report the percent error in your calculated value of Hvap for ethanol. Is your value for the Hvap greater than or less than the accepted value. Show the calculation used to arrive at the reported value. show a sample.Data: Initial Trial 1 Trial 2 Trial 3 Trial 4 Trial 5 Ptotal (kPa) Pair (kPa) Pvap (kPa) Temperature (°C) Describe how the Pair was calculated. 115 . Think of some valid sources of error to account for your difference and explain how they would contribute to the direction of your error. 1/T (K). Report your value for Hvap of ethanol. Give the equation for freezing point depression and indicate the units for each term in the expression. The freezing temperature is measured to be -1. What else do you need to measure to determine the molar mass of the solid added to the solvent? 4. This includes stating the purpose of the experiment.0 mL of water.504 g of a solid to 25. What is the molality of the solution? 5. A student adds 1. What is the molar mass of the solid above? Part B Prepare your notebook for the lab. At the start of your lab.Pre-Lab: Using Freezing-Point Depression to Find Molecular Weight Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. A measurement of the freezing temperature of a solution allows you to calculate the concentration of the solution. 3. 116 .20 °C. What is a colligative property? 2. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. an organic non-polar solvent. in g/mol. that is. and the mass of the biphenyl used. is a colligative property. to a known mass of the solvent. cyclohexane. you can use the equation above to find the molar mass of the solute. and m is the molality of the solution (in mol solute/kg solvent). the freezing temperature is lowered in proportion to the number of moles of solute added. you will Determine the freezing temperature of pure cyclohexane.Using Freezing-Point Depression to Find Molecular Weight When a solute is dissolved in a solvent. Figure 1 117 . T. Compare it to the accepted molar mass for biphenyl. Examine the freezing curves for each. Determine the freezing temperature of a solution of the biphenyl and cyclohexane. Kf is the freezing point depression constant for a particular solvent (20. OBJECTIVES In this experiment. known as freezing-point depression. C6H12. an organic solute. You will then add a known mass of biphenyl (C12H10). In this experiment. and determine the lowering of the freezing temperature of the solution. Calculate the experimental molar mass of biphenyl. This property.2°C-kg/mol for cyclohexane used in this experiment). The equation that shows this relationship is: T = Kf • m where T is the freezing point depression. By measuring the freezing point depression. not on the nature of the substance itself. you will first find the freezing temperature of the pure solvent. it depends on the ratio of solute and solvent particles. b. secure cord as shown Use for stirring Notes: a. Use the 100mm test tube. Water from water bath should be poured down the drain (not into the waste container) unless contaminated with cyclohexane. environmental hazard 118 . Do not use the temperature probe to stir. c.Equipment Information Each bin should contain: 1 – temperature probe 1 – copper stirrer After use. toxic. Chemical Safety Information Using freezing point depression to find molecular weight Chemical Cyclohexane Biphenyl Hazards flammable. health and environmental hazard toxic. Use the copper stirrer to stir. Allow the test tube to sit at room temperature to melt the probe out of the solid cyclohexane. Do not attempt to pull the probe out—this might damage it. selecting only the points in the plateau. Measure out approximately 0. Make sure the water level outside the test tube is higher than the solvent level inside the test tube. It can be propped in a plastic 250 mL beaker to facilitate measuring. Repeat steps 5-10 so that you have two trials of the freezing point cyclohexane. 7. Add the biphenyl to the cyclohexane already in the 4” test tube.779 g/ml). 9. If you do. The cyclohexane used in this experiment is flammable. this is useful when the tube is not empty. Part I Freezing Temperature of Pure cyclohexane 3. Prepare the computer for data collection by opening the file “15 Freezing Pt Depression” from the Chemistry with Computers folder. Add 3 mL of cyclohexane to your test tube (C6H12. Do not use Bunsen burners during this lab.PROCEDURE 1. With a very slight stirring motion with the probe. continuously stir the solvent during the cooling. Click . . Do this carefully so that you do not get any solvent on the upper portion of the test tube. Insert the Temperature Probe into the cyclohexne. 4. Move the mouse pointer to the beginning of the graph’s flat part. stop stirring. Lower the test tube into the water bath. MAKE SURE IT LOOKS LIKE 3 mL. 8. About 30 seconds are required for the probe to warm up to the temperature of its surroundings and give correct temperature readings. Click on the Statistics button. Carefully wipe any excess cyclohexane liquid from the probe with a paper towel or tissue and dry the sides of the test tube. Then click to begin data collection. During this time. 5. Part II Freezing Temperature of a Solution of biphenyl and cyclohexane 11. Add tap water until it is about 2/3 full. To determine the freezing temperature of pure cyclohexane. Fill your 250 mL beaker with ice. Obtain and wear goggles. Weigh a CLEAN. It may take several 119 . The mean temperature value for the selected data is listed in the statistics box on the graph. reweigh the test tube and cyclohexane to determine if there is a change in mass of the solvent. 6. 2. Store your data by choosing Store Latest Run from the Experiment menu. 10. Record this value as the freezing temperature of cyclohexane. Continue with the data aquisition until data collection has stopped (10 minute run).06 grams of biphenyl into a weighing boat. Connect the Temperature Probe to the computer interface. Add a small amount of ice to your water bath to bring the temperature down. Determine the mass of the biphenyl by weighing the test tube. Once the solid begins to form. density = 0. you will have to dump the solvent in the waste recepticle. DRY test tube. fasten the utility clamp to the ring stand so the test tube is above the ice water bath. Close the statistics box. you need to determine the mean (or average) temperature in the portion of graph with nearly constant temperature. Weigh the test tube with solvent in it and compare the masses to determine the mass of solvent. Hide the curve from your first run by clicking on the vertical axis label and unchecking the appropriate box. clean and dry your apparatus. Try to prevent the biphenyl from sticking to the inside wall of the test tube where it will be difficult (or impossible) to get into solution. solvent and biphenyl. Press the mouse button and hold it down as you drag across the flat part of the curve. compare to a similar test tube with 3 mL of water in it. and start all over by weighing out a new portion of cyclohexane. When you have completed Step 8. and check the Run 1 and Latest Temperature boxes. Calculate the experimental molar mass of biphenyl. Locate the initial freezing temperature of the solution. Then drag each box to a position on or near its respective curve. and the moles of biphenyl you found in the previous step. 5. Unlike pure cyclohexane. 3. 6. Click .cyclohexane solution. between the pure cyclohexane (t1) and the mixture of cyclohexane and biphenyl (t2). Use the original mass of biphenyl from your data table. the temperature (y) and time (x) data points are displayed in the examine box on the Time graph. Determine the difference in freezing temperatures. Record the freezing point in your data table. c. click More. Click on the vertical-axis label of the graph. using the formula. To determine the freezing point of the biphenyl. time showing all data runs: a. 4. PROCESSING THE DATA (METHOD 1) 1. Dissolution can be encouraged through gentle agitation (be careful not to splash). b. To print a graph of temperature vs. 12. Calculate moles of biphenyl solute. t = Kf • m (Kf = 20. . 2. in mol/kg. As you move the mouse cursor across the graph. using the answer in Step 2 (in mol/kg) and the mass (in kg) of cyclohexane solvent. you need to determine the temperature at which the mixture initially started to Freezing Point freeze. Label both curves by choosing Text Annotation from the Insert menu. Repeat the process with by adding another portion of biphenyl. as shown here. Compare your experimentally determined molar mass of biphenyl with the known value. Calculate molality (m). 14. in g/mol. Repeat Steps 3-8 to determine the freezing point of this mixture. 120 . t. Print the graph.2°C-kg/mol for cyclohexane). t = t1 .minutes for the biphenyl to dissolve. cooling a mixture of biphenyl and cyclohexane results in a gradual linear decrease in temperature during the time period when freezing takes place. Calculate the percent error. and typing “cyclohexane” (or “Biphenyl. Use the formula. 13. To display both temperature runs. click on the Examine button.t2.cyclohexane mixture”) in the edit box. 5. Compare your results from the two methods.PROCESSING THE DATA (METHOD 2) Here is another method that can be used to determine the freezing temperature from your data in Part II. The graph should now have two regression lines displayed. Click again. 2. 121 . Press the mouse button and hold it down as you drag only this linear region of the curve. Click on the Linear Fit button. Choose Interpolate from the Analyze menu. 3. the temperatures shown in either examine box should be equal to the freezing temperature for the biphenyl-cyclohexane mixture. Use the temperature to calculate T and your molar mass for biphenyl. Move the mouse pointer left to the point where the two regression lines intersect. 6. 4. When the small circles on each cursor line overlap each other at the intersection. where the temperature has an initial rapid decrease (before freezing occurred). Move the mouse pointer to the initial part of the cooling curve. Press the mouse button and hold it down as you drag across the linear region of this steep temperature decrease. . With a graph of the Part II data displayed. Now press the mouse button and drag over the next linear region of the curve (the gently sloping section of the curve where freezing took place). use this procedure: 1. 22 g/mol) Calculate the percent error in your determined molar mass of biphenyl Be sure to attach your calculations to the back of your report Discussion: Discuss the results of this activity. 122 . answer the below question. Is your value for the molar mass of biphenyl greater than or less than the accepted value? Think of some valid sources of error to account for your difference and explain if the effect would be cause the measured value to be erroneously high or low as compared to the actual molar mass of the biphenyl. relevant conceptual background. What will be the expected effect on the measured molecular weight? Subm it you r report on tim e and to your TA in the d ropbox on D2L. Additionally. 1.Freezing Point Depression Lab Report: Your report for this lab should include the following sections: Abstract: Your abstract should be written individually Introduction: Include a statement of purpose for this experiment. In the past. many students have listed that they accidentally lost some of the solid biphenyl during transfer to the test tube. and general equations Data: Include a data table with all necessary mass measurements Include graphs for the freezing of cyclohexane and cyclohexane-biphenyl solution Results: Report the freezing point of pure cyclohexane Report your calculated molar mass of biphenyl Determine the percent error in your calculated molar mass (the actual molar mass of biphenyl is 154. the following data were collected. For a reaction where the general form of the rate law is rate = [A]m[B]n. Write the general form of the rate law for the reaction you will be studying this week.Pre-Lab: The Rate and Order of a Chemical Reaction Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. 2.025 M Part B Prepare your notebook for the lab. A first order reaction has a rate constant of 2. What is the overall order of a reaction that has the following rate law? Rate = [A]2[B] 4.075 M [B] 0. This includes stating the purpose of the experiment. What is the order of the reaction with respect to A? What is the order of the reaction with respect to B? Initial Rate 0. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection.050 M 0. 123 .01 M/s 0.01 M/s 0. Calculate the half-life for this reaction. At the start of your lab. 3.025 M 0.025 M 0.90 x 10-4 s-1.025 M 0. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA.09 M/s [A] 0. or absorbance of light. Figure 1 124 . you will determine the effect each reactant has on the rate of the reaction. you must decide how to follow the reaction by measuring some parameter that changes regularly as time passes. By carefully varying the concentrations of the reactants. you will Conduct the reaction of KI and FeCl3 using various concentrations of reactants. In this experiment you will conduct the reaction between solutions of potassium iodide and iron (III) chloride. pH. you can determine the order of the reaction in each species. Once you select a reaction to examine. such as temperature. it undergoes a color change that can be precisely measured by a Colorimeter (see Figure 1).The Rate and Order of a Chemical Reaction OBJECTIVES In this experiment. in ionic form. Determine the order of the reaction in KI and FeCl3. Determine the rate law expression for the reaction. From this information. The reaction equation is shown below. and consequently the order of the reaction. In this way. and determine a rate law expression. you will write a rate law expression for the reaction. pressure. 2 I– (aq) + 2 Fe3+ (aq) → I2 (aq) + 2 Fe2+ (aq) As this reaction proceeds. conductance. INTRODUCTION A basic kinetic study of a chemical reaction often involves conducting the reaction at varying concentrations of reactants. o Work in a team. Chemical Safety Information The rate and order of chemical reaction Chemical Potassium iodide Iron (III) chloride Hazards toxic corrosive. toxic 125 .Equipment Information Each bin should contain: 5 – cuvettes with lids Do not leave cuvettes in colorimeter at the end of class 2 – 10mL graduated cylinders Notes: a. o Measure the KI and the water into the same graduated cylinder. To succeed in this lab: o Ensure good technique and decrease waste by practicing the procedure using tap water first. This lab is very time sensitive. Then mix them with the FeCl3. 5 5. g) Turn the knob of the Colorimeter to the Blue LED position (470 nm).5 5 2.0 2 5. e) Type 0 in the edit box.5 a) Consider opening an online stopwatch to make sure that all measurements are started reproducibly at the same time after mixing. Crystal Violet” from the Chemistry with Computers folder.020 M FeCl3 solution in a separate 100 mL beaker. CAUTION: The FeCl3 solution in this experiment is prepared in 0. h) Type 100 in the edit box.5 6 2. Connect a Colorimeter to Channel 1 of the Vernier computer interface. Start the Logger Pro program on your computer. 4. set the wavelength on the Colorimeter to 470 nm. d) Measure the FeCl3 solution using a graduated cylinder and pour it into the test tube.0 2. c) If adding water. and proceed directly to Step 5. a) Two 10 mL graduated cylinders. then click . Obtain and wear goggles.0 2. d) Turn the wavelength knob on the Colorimeter to the “0% T” position. use the correct volume for each trial based on the table below. This step describes the process for conducting the trials using the Trial 1 volumes.0 2.0 0. d) Approximately 60 mL of distilled water in a third 100 mL beaker.0 0. b) If your Colorimeter has a CAL button. When you repeat this process.020 M KI solution in a 100 mL beaker. If your Colorimeter does not have a CAL button. measure out the water using a graduated cylinder and add to the test tube first. b) Prepare a clean cuvette. then click . continue with this step to calibrate your Colorimeter. Trial FeCl3 (mL) KI (mL) H2O (mL) 1 5. 5. Place the blank in the cuvette slot of the Colorimeter and close the lid.0 5. Set up and calibrate the Colorimeter. 6. 126 .5 5. Obtain the materials you will need to conduct this experiment. 2. c) Choose Calibrate CH1: Colorimeter from the Experiment menu. press the CAL button. f) When the displayed voltage reading for Reading 1 stabilizes. i) When the voltage reading for Reading 2 stabilizes. click . During this experiment you will conduct 6 trials. Open the file “30b. a) Prepare a blank by filling an empty cuvette ¾ full with distilled water. b) Approximately 50 mL of 0.5 4 5.0 5. click .0 2.1 M HCl and should be handled with care.PROCEDURE 1. c) Approximately 50 mL of 0.5 2.0 3 5.5 2. 3. Click to begin collecting absorbance data. and close the lid and begin collecting absorbance data. DATA ANALYSIS 1. carefully remove the cuvette from the Colorimeter. Slide the cursor to the initial time point. When the data collection is complete. practice several times with water before attempting with the KI and FeCl3 solutions. Rinse and clean the beakers and the cuvette for the next trial (this needs to be done as soon as the trial is over). fill the cuvette ¾ full with the mixture. NOTE. Data will be gathered for 2 minutes. 7. select the Slope button. Repeat Steps 6–9 to conduct Trials 2–6. place it in the Colorimeter. 2. followed by I-. Observe the progress of the reaction in the beaker. On the toolbar. 127 . Wipe the outside of the cuvette with a tissue. Examine the graph of the first trial. What is the order of the reaction in FeCl3 and KI? Write the rate law expression for the reaction. 8. 10. Cover the end of the test tube with your thumb and quickly invert to mix. Calculate the initial molar concentration of FeCl3 and KI for each reaction and prepare a data table containing the concentrations of each reaction and the initial reaction rate.e) Measure the KI solution using a graduated cylinder. IT IS IMPORTANT THAT THE SLOPE HAVE MORE THAN ONE SIGNIFICANT FIGURE. Add the solutions to your medium sized test-tube in this order: water followed by Fe3+. Record the slope as the initial rate of the Trial 1 reaction. Dispose of the contents of the beaker and cuvette as directed. f) Within 15 seconds of mixing the two solutions. 3. use the same technique to analyze Trials 2–5 that you used to analyze Trial 1. The timing of this step is imperative to receiving useful data. When you complete Step 9. 9. This tool will determine the initial slope and thus approximate the initial rate of the reaction. 128 . How does absorbance relate to concentration? How can you use rate and concentration data to determine the order in a particular reactant and thus a rate law? .The Rate and Order of a Chemical Reaction Lab Report: Name_____________________________ Date______________ Lab Section______________ There is not a form al lab report for this lab. Com plete the below pages and su bm it them to your TA before leaving lab. Briefly state the purpose of the lab. Why can the absorbance versus time data can be used as a rate. determine a value for the rate constant. Compare the average rate from trials 1 and 2 to the average rate of trials 5 and 6 to determine the order in Fe3+. Report the order in I-. Show your work. Show your work. Report the order in Fe3+. Compare the average rate from trials 1 and 2 to the average rate of trials 3 and 4 to determine the order in I-. Show your work. Write out the rate law. Trial [FeCl3](M) [KI] (M) Initial rate (units of absorbance/s) Value must have >1 sig figure 1 2 3 4 5 6 Incl ude one gra ph sho win g all of the trials overlaid on one another.Results: Complete the following table. 129 . Using the units of rate as “units of absorbance”/s. 130 . remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. what is the value of Kc? 3.0020 M KSCN. Part B Prepare your notebook for the lab. If the equilibrium values of Fe3+. This includes stating the purpose of the experiment.6 x 10-4 M and 5.in the solution. A solution is prepared by adding 18 mL of 0. At the start of your lab.5 x 10-4 M. 3. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection.200 M Fe(NO3)3 and 2 mL of 0. Kc Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. Write the general form of the dilution equation.and FeSCN2+ are 9. SCN. Calculate the initial concentrations of Fe3+ and SCN. 4.Pre-Lab: Chemical Equilibrium: Finding a Constant.7 x 10-5 M respectively. 2. Write the equilibrium constant expression for the experiment you will be studying this week. The initial [Fe3+] in the standard solution is 900 times larger than [SCN-]i. Since the reaction produces the FeSCN2+ ions and this ion transmits the color red. Keq.ions are consumed in order to minimize the disturbance. for the following chemical reaction: FeSCN2+(aq) Fe3+(aq) + SCN–(aq) iron(III) thiocyanate thiocyanoiron(III) When Fe3+ and SCN. This high initial concentration Fe3+ ions on the left side of the equation forces the reaction far to the right. the product’s concentration. one mole of SCN– is used up. the system responds so as to minimize the disturbance and return the system to a state of equilibrium. 131 . containing different concentrations of these three ions (Fe3+. To prepare the standard solution. aA + bB cC + dD This equation gives the equilibrium constant expression of: Keq = [C]c[D]d/[A]a[B]b In order to determine the equilibrium concentrations for the three ions a standard solution needs to be prepared. at equilibrium is assumed to be equal to the [SCN–]i. the blue LED setting on the Colorimeter is used.are combined. a dynamic equilibrium is established between these two ions and the FeSCN2+ ion. it is assumed that for every mole of FeSCN2+ produced. The computer-interfaced Colorimeter measures the amount of blue light absorbed by the colored solutions (absorbance. or trials. Kc. Because the red solutions absorb blue light very well. SCN-. A). In order to calculate the equilibrium constant. [FeSCN2+]std. In this experiment four separate equilibrium systems. Using stoichiometry and the balanced equation.Chemical Equilibrium: Finding a Constant. and FeSCN2+) will be determined experimentally. The Keq. for the reaction. The values for these equilibrium concentrations will be substituted into the equilibrium constant expression to see if Keq is indeed constant despite varied initial concentrations for the reactants. a very large concentration of Fe3+ will be added to a small initial concentration of SCN– (hereafter referred to as [SCN-]i. According to LeChatelier's principle. which states that when a system in dynamic equilibrium is disturbed. using up nearly 100% of the initial SCN– ions. it is necessary to know the concentrations of all the ions at equilibrium. the solution’s absorbance of blue light can be measured through the use of a colorimeter (see Figure 1). Kc The purpose of this lab is to experimentally determine the equilibrium constant. Thus since nearly all of the SCN. is determined by using the Law of Mass Action. The concentration of FeSCN2+ for any of the equilibrium systems. In this case the concentration is the FeSCN2+ ion and the absorbance is blue light (470 nm). to the absorbance of the standard solution. and [FeSCN2+]eq. the equilibrium constant. For each mole of FeSCN2+ ions produced. At equilibrium the [Fe3+] and [SCN-] can be determined according to the following equations: [Fe3+]eq = [Fe3+]i – [FeSCN2+]eq [SCN–]eq = [SCN–]i – [FeSCN2+]eq Knowing the values of [Fe3+]eq. can be found by comparing the absorbance of each equilibrium system. there is a direct relationship between a solution’s concentration and its absorbance. Astd. one less mole of Fe3+ and SCN.ions will be found in the solution (see the 1:1 ratio of coefficients in the equation on the previous page). as the concentration of FeSCN2+ increases so will the absorbance of blue light (see Figure 2). you can now calculate the value of Kc. according to the following equation: [FeSCN2+]std/Astd = [FeSCN2+]eq/ Aeq Since the concentration of [FeSCN2+]std is known and the all of the absorbances for the equilibrium solutions and the standard are measured and recorded all that needs to be done is to solve for the unknown. Aeq [FeSCN2+]eq = A X [FeSCN2+]std std Knowing the [FeSCN2+]eq allows you to determine the concentrations of the other two ions at equilibrium. In other words. trials 1-4. Aeq. 132 .Figure 1 Figure 2 According to Beer’s Law. [SCN–]eq. Chemical Safety Information Chemical equilibrium: Finding a constant. only procure small amounts of the chemicals you need. This lab is done on the micro scale. you will determine the equilibrium constant. To help reduce waste. Kc Chemical Iron (III) nitrate in nitric acid Potassium thiocyanate Hazards corrosive.OBJECTIVE In this experiment. secure cord as shown 1 – 10mL serological pipet with bulb Do not aspirate liquid into the bulb. Kc. for the following chemical reaction: FeSCN2+(aq) Fe3+(aq) + SCN–(aq) iron(III) thiocyanate thiocyanoiron(III) Equipment Information Each bin should contain: 5 – cuvettes with caps 1 – temperature probe After use. oxidizer toxic 133 . Notes: a. Obtain about 25 mL of distilled water in a 100 mL beaker. Pipet 2. 3. Choose Calibrate CH1: Colorimeter (%T) from the Experiment menu and then click . Be sure to clean and dry the stirring rod after each mixing. Prepare the computer for data collection by opening the file “Lab 8 Equilibrium” from the Chemistry 228 folder of Logger Pro 5.0020 M Fe(NO3)3 into a clean. If your Colorimeter does not have a CAL button. 4. b. All solutions should be free of bubbles. to bring the total volume of each test tube to 10 mL. Mix each solution thoroughly with a stirring rod.0020 M KSCN into the same test tube. continue with this step to calibrate your Colorimeter. Close the lid.200 M Fe(NO3)3 into a 20 150 mm test tube labeled “5”. Then pipet 3. Use a pipet pump or bulb to pipet all solutions.PROCEDURE 1. Calibrate the Colorimeter. respectively. Proceed directly to Step 7. Pipet 1 mL of 0. Open the Colorimeter lid. When the LED stops flashing. c. Volumes added to each test tube are summarized below: Test Tube Number Fe(NO3)3 (mL) KSCN (mL) H2O (mL) 1 2 3 4 5 5 5 5 2 3 4 5 3 2 1 0 3. 134 . Prepare a standard solution of FeSCN2+ by pipetting 9 mL of 0. Handle cuvettes only by the top edge of the ribbed sides. respectively. Prepare a blank by filling a cuvette 3/4 full with distilled water.0020 M KSCN into another clean. remember: All cuvettes should be wiped clean and dry on the outside with a tissue. Pour about 25 mL of the 0. the calibration is complete.0 M HNO3 and should be handled with care. place it in the cuvette slot of the Colorimeter. 2. Label four 20 150 mm test tubes 1-4. 4 and 5 mL of this solution into Test Tubes 1-4. Holding the cuvette by the upper edges. Stir thoroughly. Then release the CAL button. 2.0 mL of this solution into each of the four labeled test tubes. Always position the cuvette with its reference mark facing toward the white reference mark at the top of the cuvette slot on the Colorimeter. Pipet 5. dry 100 mL beaker. To correctly use a Colorimeter cuvette. If your Colorimeter has a CAL button. a. Kc. dry 100 mL beaker. Pour about 30 mL of 0. 1 and 0 mL of distilled water into Test Tubes 1-4. First Calibration Point d. 6. CAUTION: Fe(NO3)3 solutions in this experiment are prepared in 1. Measure and record the temperature of one of the above solutions to use as the temperature for the equilibrium constant. Obtain and wear goggles. Press the < or > button on the Colorimeter to select a wavelength of 470 nm (Blue) for this experiment. Press the CAL button until the red LED begins to flash. Connect the Colorimeter to the computer interface. From the table. Wipe the outside of the cuvette with a tissue and then place the cuvette in the Colorimeter. Click to begin data collection. Rinse it twice with ~1 mL portions of the Test Tube 1 solution. and 5 (the standard solution). Rinse the cuvette twice with the Test Tube 2 solution and fill the cuvette 3/4 full. Dispose of all solutions as directed by your instructor. j. Then click . When the displayed voltage reading for Reading 1 stabilizes. Turn the knob of the Colorimeter to the Blue LED position (470 nm). Second Calibration Point h. Empty the water from the cuvette. . type “1” (the trial number) in edit box. click . Follow the Step-c procedure to find the absorbance of this solution. f.e. You are now ready to collect absorbance data for the four equilibrium systems and the standard solution. record the absorbance values for each of the five trials in your data table. After closing the lid. Type “2” in the edit box and press ENTER. wait for the absorbance value displayed in the meter to stabilize. c. 135 . then click 7. a. Repeat the Step-d procedure to find the absorbance of the solutions in Test Tubes 3. Type “0” in the edit box. and press the ENTER key. g. Type “100” in the edit box. f. i. Discard the cuvette contents as directed by your teacher. b. click . Turn the wavelength knob on the Colorimeter to the “0% T” position. 4. g. d. When the displayed voltage reading for Reading 2 stabilizes. e. [Fe3+]eq: Calculate the concentration of Fe3+ at equilibrium for Trials 1-4 using the equation: [Fe3+]eq = [Fe3+]i – [FeSCN2+]eq 6. based on its dilution by Fe(NO3)3 and water: KSCN mL [SCN–]i = total mL (0.0020 M Fe(NO3)3 solution. 8. respectively. Calculate the initial concentration of Fe3+. 3.PROCESSING THE DATA 1. Calculate this for the other three test tubes.00020 M. [SCN–]i = (2 mL / 10 mL)(0. 2. [FeSCN2+]eq is calculated using the formula: Aeq [FeSCN2+]eq = A std [FeSCN2+]std where Aeq and Astd are the absorbance values for the equilibrium and standard test tubes. Be sure to show the Kc expression and the values substituted in for each of these calculations. Calculate Kc for Trials 1-4. based on the dilution that results from adding KSCN solution and water to the original 0.0020 M) = 0.00040 M.0020 M) In Test Tube 1. How constant were your Kc values? 136 . Calculate [Fe3+]i using the equation: Fe(NO3)3 mL [Fe3+]i = (0. Using your four calculated Kc values.0020 M) total mL This should be the same for all four test tubes. [SCN–]eq: Calculate the concentration of SCN. 4. 5. Calculate [FeSCN2+]eq for each of the four trials. and [FeSCN2+]std = (1/10)(0. Write the Kc expression for the reaction in the Data and Calculation table. Calculate the initial concentration of SCN–. See Step 2 of the procedure for the volume of each substance used in Trials 1-4. determine an average value for Kc.at equilibrium for Trials 1-4 using the equation: [SCN–]eq = [SCN–]i – [FeSCN2+]eq 7.0020) = 0. 137 .Equilibrium Lab Report: Your report for this lab should include the following sections: Abstract: Your abstract should be written individually Introduction: Include a statement of purpose for this experiment. and general equations Data: Include your data table with initial concentrations of each reactant for each trial Results: Include a results table with calculated Keq values for each trial and an average value for Keq Be sure to attach hand written sample calculations to the back of your report Discussion: Discuss the experiment and any possible sources of error In addition. answer the following question as part of your report: 1. How are you Keq values to each other? Are they close enough to justify the assertion that an equilibrium constant is constant? 2. relevant conceptual background. What factors could have led to variations in Keq between trials. Subm it you r report on tim e and to your TA in the d ropbox on D2L. Cobalt is highly toxic.Le Chatelier’s Principle in a Cobalt Complex A chemical system will eventually come to a dynamic state of equilibrium. the composition of the system changes in a way to reduce or counteract the disturbance. [CoCl4]2-(aq) + 6H2O(l) [Co(H2O)6] 2+(aq) + 4Cl-(aq) Equipment Information There are no bins this week. Notes: a. b. In this activity. Chemical Safety Information Equilibrium of cobalt chloride Chemical Cobalt chloride Hydrochloric acid Silver nitrate Hazards corrosive. Le Chatlier’s principle states that wen a system in equilibrium is disturbed. corrosive. health and environmental hazard corrosive oxidizer. environmental hazard 138 . Complex ions will be discussed in more detail during the third term but are formed when transition metals are bound to electron pair donors through coordinate covalent bonds. toxic. toxic. you will be monitoring the effect of disturbances of the equilibrium between two complex ions of cobalt. Concentrated hydrochloric acid is very corrosive. such as increasing a reactant or products concentration or by heating or cooling the reaction mixture. This state of equilibrium can be altered by adding some sort of stress to the reaction. 3. exercise caution and be sure to wash your hands after the lab. Take the test-tubes from step 6 and place in your ice bath for two minutes. 6. 9. slowly add 10 mL drop wise to your cobalt chloride solution until vivid color change is observed. something that you can use to compare any color changes to. Note any color changes. filling only half-way. Divide the resulting solution amongst 4 small test-tubes. Note any color changes. 2. To one test-tube add distilled water drop wise. you have just produce the [CoCl4]2(aq) complex ion. Be sure to dispose of all waste properly. 5. Sliver nitrate will stain skin and clothes. 8. Do not overfill the test-tubes (about 5 mL in each). If no color change. Here. Keeping track of which test-tube is which. Obtain 10 mL of a 0. To a test-tube from step 4.1 M cobalt chloride solution.Hazards: HCl is very corrosive! Exercise caution around solution and vapors. Note the color. add 0. Procedure 1. Keep one as a control. Prepare a hot bath by half filling a beaker with water and placing it on a hot plate. Heat to boiling on medium heat. Note the color. 7. Do this step in the fume hood. rather than adding less concentrated solutions to the acid. Always clean up any spills! Disposal: Cobalt chloride and silver nitrate as well as any excess acid must be properly disposed in the labeled waste jars. until the color changes. 139 . prepare an ice/water cold in a second beaker. while gently stirring with your stir rod. Additionally. take a test-tube from step 4 and the test-tube generated in step 5 and place them in the hot bath for 2 minutes.1 M silver nitrate drop wise until a precipitate forms. Always make sure to add acid to less concentrated solutions. then add additional HCl drop wise. Note any color changes. Obtain 15 mL of Concentrated HCl solution and while gently stirring with your stir rod. 4. Step 6 Color change (and observations). Use Le Chatlier’s principle and provide evidence. What was the effect of adding excess chloride ions. Step 5 Color change (and observations). Questions: 1. Step 7 Color change (and observations). Step 8 Color Change (and observations). 140 .Last Name___________________ First Name________________________ TA___________ Color [CoCl4]2-(aq) [Co(H2O)6] 2+(aq) Step 3 Color change (and observations). Use Le Chatlier’s principle and provide evidence (would heat be considered as a reactant or product?). When perturbing the equilibrium with heating and cooling. su bmit the w orksheet on tim e and to you r TA in the d ropbox on D2L .2. Once the w orksheet is com plete. Com plete the above pages u sing the Microsoft version of this file that is available for d ow nload on the lab D2L page. [CoCl4]2-(aq) + 6H2O(l) [Co(H2O)6] 2+(aq) + 4Cl-(aq) 3. how many times do you think the equilibrium can be shifted before it stops working? Why? How about modification of the equilibrium through changes in concentration? There is not a formal lab report for this lab. How did the addition of silver nitrate affect the equilibrium if neither silver ions nor nitrate ions are in the equilibrium expression? Use Le Chatlier’s principle and provide evidence. 4. Based upon the heating and cooling of the two equilibrium mixtures. Additionally. write a net ionic equation to describe the precipitation reaction. 141 . propose if the reaction is endothermic or exothermic. nitrous acid (HNO2) or nitric acid (HNO3)? 4. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. 2.CH229 LABS Pre-lab: Acid Rain Part A Answer the following questions in your lab notebook (be sure to show your work for any calculations): 1. What is the conjugate base of nitrous acid (HNO2)? 3. Carbon dioxide (CO2) reacts with water to produce carbonic acid (H2CO3). At the start of your lab. 142 . summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. Which is a stronger base. Describe the method you will use in this lab to generate the acids found in acid rain. Part B Prepare your notebook for the lab. Write the balanced chemical equations for this reaction and showing what happens when carbonic acid is dissolved in water. nitrite (NO2-) or nitrate (NO3-)? 5. This includes stating the purpose of the experiment. Which is a stronger acid. Nitrogen dioxide gas dissolves in rain drops and forms nitrous and nitric acid: CO 2 (2) 2 NO2(g) + H2O(l) HNO2(aq) + HNO3(aq) Sulfurous acid is produced from another air pollutant. YOU BUY IT. SO2.Lab: Acid Rain In this experiment. H2SO3 Carbonic acid occurs when carbon dioxide gas dissolves in rain droplets of unpolluted air: (1) CO2(g) + H2O(l) H2CO3(aq) Nitrous acid and nitric acid result from a common air pollutant. HNO2 nitric acid. Most of the sulfur dioxide gas in the atmosphere results from burning coal that contains sulfur impurities. Sulfur dioxide dissolves in rain drops and forms sulfurous acid: NO 2 H 2 CO 3 H NO 2 H NO 3 H2SO3 SO2 (3) SO2(g) + H2O(l) H2SO3(aq) In the procedure outlined below. 143 . OBJECTIVES In this experiment. return it to its storage buffer. After the measurement. you will Generate three gaseous oxides. even though protected by a plastic shield. rinse the electrode probe thoroughly with distilled water and dry them gently with a Kim-wipe tissue. sulfur dioxide (SO2). you will first produce these three gases. you will observe the formation of four acids that occur in acid rain: carbonic acid. When finished. The glass electrode is fragile. You will then bubble the gases through water. HNO3 sulfurous acid. again rinse it. nitrogen dioxide (NO2). H2CO3 nitrous acid. and NO2 Simulate the formation of acid rain by bubbling each of the three gases into water and producing three acidic solutions Measure the pH of the three resulting acidic solutions to compare their relative strengths Note about the pH sensor: The pH meter is a device used to measure the electrical potential of a particular process. Handle it with care. producing the acids found in acid rain. Be very cautious about bumping the electrode bulb on the bottom of the container containing a test solution. Before making a pH measurement. The glass bulb is part of an electrode that is responsive to the hydronium ion activity of the test solution. The acidity of the water will be monitored with a pH Sensor. Most of the nitrogen dioxide in our atmosphere is produced in automobile exhaust. IF YOU BREAK IT. CO2. Storage solution available in the stockroom. Refill pH probe with storage solution if needed. environmental hazard 144 . Check the pH probe for breakage. Chemical Safety Information Acid Rain Chemical Hydrochloric acid Sodium bicarbonate Sodium bisulfite Sodium nitrite Hazards corrosive none corrosive oxidizer. b.Equipment Information Each bin should contain: 1 – pH probe 7 – plastic pipets After use. secure cord as shown Notes: a. acutely toxic. as shown in Figure 2. Figure 2 7. Remove the pH Sensor from the pH storage solution. This will be the initial pH value that you will compare your resulting changes to. and “NaHSO3”. enter the pH value on the bottle. When you release the bulb. Use a utility clamp to attach a 20 200 mm test tube to the ring stand. rinse it off with distilled water. CAUTION: HCl is a strong acid. Calibrate the pH probe following this procedure: Use the 2-point calibration option of the Vernier data-collection program. is generated in this pipet and is heavier than air. Squeeze the bulb to expel some of the air. Figure 3 Carbon dioxide. as shown in Figure 1. When you release the bulb. and place the open end of the pipet into the solid NaHCO3. Obtain and wear goggles. 8. When the displayed voltage stabilizes. Connect the pH Sensor to the computer interface. Place the electrode into one of the buffer solutions. 9. “NO2” and “SO2”. CAUTION: Avoid inhaling dust from these solids. and place the open end of the pipet into a beaker containing 1. CO2. Continue to draw solid into the pipet until there is enough to fill the curved end of the bulb. Rinse the electrode with distilled water. enter the pH value on the bottle. Label the long-stem pipets with the formula of the gas they will contain: “CO2”. 8. rinse the electrode and place it into a second buffer solution. Gently swirl the pipet that contains NaHCO3 and HCl. Gently squeeze the HCl pipet to add about 20 drops of HCl solution to the solid NaHCO3. Obtain a beaker containing solid NaHCO3. so that HCl drops do not escape. remove the HCl pipet and place it open side up in the 100 mL beaker. It is now ready to beplaced in the sample to be measured.0 M HCl. Prepare the computer for data collection by opening the file “Exp 23 acid rain” from the Chemistry with Computer folder of Logger Pro. to prevent spillage. with the stem up. Rinse the tip of the electrode in distilled water. You can use a 100 mL beaker to support the pipets. Insert the narrow stem of the HCl pipet into the larger opening of the pipet containing the solid NaHCO3. Wear gloves for this step. 3. For the next calibration point. “NaNO2”. Place the Beral pipet in the 100 mL beaker. Squeeze the bulb of the pipet labeled “NaHCO3” to expel the air. Add about 4 mL of distilled water to the test tube.PROCEDURE 1. solid NaHCO3 will be drawn up into the pipet. Label the short-stem pipets with the formula of the solid they will contain: “NaHCO3”. When the voltage reading displayed on the computer or calculator screen stabilizes. 145 . 6. so it stays in the pipet. When finished. Measure and record the pH of the lab distilled water (it may not be pH =7). Obtain three short-stem and three long-stem Beral pipets as shown in Figure 1. Gently hold the pipet with the stem pointing up. HCl will be drawn up into the pipet. and place it into the distilled water in the test tube. Repeat the Step 3 procedure to add solid NaNO2 and NaHSO3 to the other two Beral pipets. Obtain another Beral pipet and label it HCl. 2. Figure 1 4. gently squeeze the bulb of the pipet so that CO2 slowly bubbles up through the solution. is generated in this pipet. and copy/paste each graph to the digital copy of the worksheet that you will submit to your TA. Nitrogen dioxide. 19. 16. is generated in this pipet. NO2. as shown in Figure 3. Use both hands to squeeze all of the gas from the bulb. but it will be still be displayed while you do your second and third trials. Discard the contents of the test tube as directed by your TA. From the Experiment menu. If not. 146 . rinse the pH Sensor with distilled water and return it to the sensor storage solution. When you are finished.10. 11. 13. Figure 4 17. by subtracting the initial pH from the final pH. choose Store Latest Run. and examine the minimum and maximum values in the pH box displayed on the graph. click . Keep the bulb completely collapsed and insert the long stem of the pipet down into the gas-generating pipet labeled “NaHCO3”. 12. Record the initial and final pH values in your lab notebook. Add 4 mL of tap water to the test tube. To begin collecting data on the computer. To confirm these two values. After 15 seconds have elapsed. Store the gas-generating pipet in the 100 mL beaker and invert the long stem pipet to keep the CO2 in until the next step. examine the data in the table and determine the initial pH value (before CO2 was added) and the final pH value (after CO2 was added and the pH stabilized). 14. When data collection stops after 120 seconds. . Place the pH Sensor in the test tube and check to see that the input display shows a pH that is about the same as the previous initial pH value. Repeat the procedure in Steps 5-13 but this time adding HCl to the pipet containing solid NaHSO3. Label all three curves by choosing Text Annotation from the Insert menu. This stores the data so it can be used later. Squeeze all of the air from the bulb of the long-stem pipet labeled “CO2”. Sulfur dioxide. Remove the pH Sensor from the test tube and rinse its tip thoroughly with distilled water and return it to the sensor storage solution. PROCESSING THE DATA For each of the three gases. calculate the change in pH (pH). alongside the pH Sensor. rinse the test tube again and refill it. Close the Statistics box by clicking in the upper left corner of the box. Record these values in the data table in your lab notebook. Release the pressure on the bulb so that it draws gas up into it. Repeat the procedure in Steps 5-13 finally adding HCl to the pipet containing solid NaNO2. Be sure the tip of the long-stem pipet remains above the liquid in the gas-generating pipet and does not collect any unreacted solid. 15. Clean and return the seven Beral pipets to the stockroom. Leave all three gas-generating pipets in the 100 mL beaker until Step 18. click the Statistics button. SO2. Rinse the test tube thoroughly with tap water. so that its tip extends into the water to the bottom of the test tube (see Figure 4). and typing “carbon dioxide” (or “nitrogen dioxide”. 18. Insert the long-stem pipet labeled “CO2” into the test tube. or “sulfur dioxide”) in the edit box. Make sure to include the final pH and change in pH for all three gases. Data: Neatly draw a data table below that captures all of the relevant information. 147 . describe if it is it possible to control how much of each acid dissolved in the water before measuring the final pH. With the technique used. Say something about the effects of acid rain.Acid Rain Lab Report Lab Report: Name_____________________________ Date______________ Lab Section______________ There is not a form al lab report for this lab. Include information as to how the acids were generated. Briefly describe the purpose of the lab. Com plete the below pages and su bm it them to your TA before leaving lab. Results: Answer the following questions. Use the internet or the appendix in your text to look up the Ka values for HNO2. 4. Ka (the acid ionization constant) is a measure of acid strength. Coal from western states such as Montana and Wyoming is known to have a lower percentage of sulfur impurities than coal found in the eastern United States. Which gas (or gases) caused the largest drop in pH? What was the change in pH? 3. Which acid is the weakest? Which is the strongest? 7. Attach copies of each curve pH curve. H2SO3. Which gas from this experiment would cause rainfall in unpolluted air to have a pH value less than 7 (sometimes as low as 5. which gas caused the smallest drop (change) in pH? What was the change in pH? 2. All three gases from this lab are produced by man but one occurs naturally at relatively high and constant concentrations. Can the nitrogen oxides in automobile exhaust contribute to acid rain? Use specific information from this lab to answer the question. and H2CO3. High temperatures in the automobile engine cause nitrogen and oxygen gases from the air to combine to form nitrogen oxides.6)? 6. 5. the higher the value of the Ka the stronger the acid. 1. In this experiment. How would burning low-sulfur coal lower the level of acidity in rainfall? Use specific information from this lab to answer the question. 148 . 00 M HC2H3O2 required to prepare 50 mL of a 0. Ka Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. for the ionization of acetic acid.10 M solution of HCl? 3. What would be the predicted pH of a 1. Write the equilibrium constant expression.0 M solution of HC2H3O2? 5. What would be the pH of a 0. 4. This includes stating the purpose of the experiment. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. in mL. Determine the volume. At the start of your lab.Pre-lab: Acid Ionization Constant. Ka. 149 . summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. of 2. What would the pH of a 0.30 M HC2H3O2 solution? Part B Prepare your notebook for the lab.10 M solution of NaOH be (in theory)? 2. HC2H3O2. The glass electrode is fragile. starting with solutions of different initial concentrations. you will Gain experience mixing solutions of specified concentration Experimentally determine the dissociation constant. Figure 1 150 . YOU BUY IT. Ka. rinse the electrode probe thoroughly with distilled water and dry them gently with a Kim-wipe tissue. The glass bulb is part of an electrode that is responsive to the hydronium ion activity of the test solution. HC2H3O2. Before making a pH measurement. is a weak acid that ionizies according to the balanced chemical equation: – HC2H3O2(aq) ↔ H+(aq) + C2H3O2 (aq) In this experiment. you will experimentally determine the ionization constant. IF YOU BREAK IT. Handle it with care. of an acid Investigate the effect of initial solution concentration on the extent of dissociation MATERIALS Vernier pH Sensor pipet bulb 10 mL graduated cylinder 50 mL volumetric flasks Note about the pH sensor: The pH meter is a device used to measure the electrical potential of a particular process.Acid Ionization Constant. return it to its storage buffer. Ka Acetic Acid. OBJECTIVES In this experiment. Be very cautious about bumping the electrode bulb on the bottom of the container containing a test solution. After the measurement. again rinse it. for acetic acid. Ka. When finished. even though protected by a plastic shield. Equipment Information Each bin should contain: 1 – 10mL graduated cylinder 1 – pH probe After use. Chemical Safety Information Acid Dissociation Constant Chemical Acetic acid Vinegar Hazards corrosive corrosive 151 . Have your TA check your calculations before you start. Check the pH probe for breakage. secure cord as shown 1 – 50mL volumetric flask Notes: a. b. When the voltage reading displayed on the computer or calculator screen stabilizes. Measure the pH of the vinegar and record your value. When the displayed voltage stabilizes. rinse the electrode and place it into a second buffer solution.0 M. 7. use a wash bottle filled with distilled water or a dropper for the last few mL. 5. 8. When done. enter the pH value on the bottle. Use a utility clamp to secure a pH Sensor to a ring stand as shown in Figure 1. 3. 6. 4. Swirl the solution vigorously. CAUTION: Use care when handling the acetic acid. 9. measure the required volume of ~2 M (write down the actual concentration) acetic acid and pour into the volumetric flask. It is now ready to beplaced in the sample to be measured. enter the pH value on the bottle. Mix thoroughly. To prevent overshooting the mark. Calculate the volume of a 2.20 and 1. (Note: Readings may drift without proper swirling!) Record the measured pH reading in the data table in your lab notebook.0 M HC2H3O2 stock solution necessary to make 50 mL of each solution you were assigned. Rinse the electrode with distilled water. place the pH Sensor in distilled water. Place the electrode into one of the buffer solutions. Determine the pH of your solution as follows: Use about 40 mL of distilled water in a 100 mL beaker to rinse the pH Sensor. Repeat the procedure for your remaining assigned solutions. Put approximately 25 mL of distilled water into a 50 mL volumetric flask. 152 . Using your graduated cylinder. Use the remaining 30 mL portion to determine pH.PROCEDURE 1. Prepare the computer for data collection by opening the file “Exp 27 Acid Dissociation Ka” from the Chemistry w/ Computer folder of Logger Pro. Obtain and wear safety goggles. For the next calibration point. Rinse the tip of the electrode in distilled water. Calibrate the pH probe following this procedure: Use the 2-point calibration option of the Vernier data-collection program. Pour about 20 mL of your first assigned solution into a clean 100 mL beaker and use it to thoroughly rinse the sensor. It can cause painful burns if it comes in contact with your skin or gets into your eyes. 2. Fill the flask with distilled water to the 50 mL mark. Connect the probe to the computer interface. Your TA will assign each group two different concentrations of HC2H3O2 between 0. Obtain 20 mL of vinegar. PROCESSING THE DATA 1. Calculate the [H+]eq from the pH values for each solution. 2. Use the obtained value for [H+]eq and the equation: – HC2H3O2(aq) H+(aq) + C2H3O2 (aq) – to construct an ICE table and determine [C2H3O2 ]eq and [HC2H3O2]eq. 3. Substitute these equilibrium concentrations into the Ka expression for HC2H3O2. 4. Compare your results with those of other students. 5. Using your calculated value of Ka for acetic acid and the pH of the commercial vinegar, determine the concentration of acetic acid in the vinegar solution. 153 Acid Ionization Lab Report: Name_____________________________ Date______________ Lab Section______________ There is not a form al lab report for this lab. Com plete the below pages and su bm it them to your TA before leaving lab. Briefly describe Ka. How does Ka relate to acid strength? Describe how you can use the pH of an aqueous acid solution and its initial concentration to determine the Ka. Data: Insert a data table below that captures all of the relevant information. Make sure to include the concentrations of the HC2H3O2 solutions you made and the volume of the stock solution and water used to make the solutions. Include the pH of vinegar . 154 Results: 2. Report your calculated Ka value for the first concentration used. Include a copy of the ICE table. 3. Report your calculated Ka value for the second concentration used. Include a copy of the ICE table. 4. Report the class average Ka. 5. Compare the class average experimentally determined Ka with the accepted value at 25 °C (1.8 x 10-5) and calculate the percent error. 155 6. Should the initial HC2H3O2 concentration have any effect on Ka? Briefly explain. 7. Calculate the percent ionization of acetic acid for each concentration you used. 8. What effect does initial HC2H3O2 concentration seem to have on the percent ionization? 9. Calculate the concentration of acetic acid in the vinegar solution 156 What would be the expected pH if a Student A accidentally diluted 9. The Ka of acetic acid is 1.18M ? 157 . In the past. Suppose that Student A was supposed to make a 0.8 x 10-5.0 mL of 2. a. what would be the resultant Ka of acetic acid given that they expected the acid to have an initial concentration of 0. many students have listed that the accidental addition of too much acetic acid contributed greatly to the difference between the experimental value and the accepted value.10.0 M acetic acid to 100.18 M solution by diluting 9.0 mL. The expected pH for this solution is 2.0 mL) of the acid to 100 mL? b. If student A measured the above calculated pH.74.1 mL (instead of 9. Pre-lab: Titration of a Diprotic Acid: Identifying an Unknown Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. This includes stating the purpose of the experiment. Give the balanced chemical reaction for the titration of a generic diprotic acid. 3. A student completes a titration of an unknown diprotic acid.48 mL of 1.0 M NaOH to reach the second equivalence point. In this experiment. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA.0 mL of 0. It requires 13. At the start of your lab. with potassium hydroxide.10 M NaOH.10 M H2SO4 with 0. What is a diprotic acid? Give an example not found below in the text for this experiment.79 g of the acid is dissolved in 250. When titrating 50. 0. how many mL of NaOH will you have added to reach the 1st equivalence point? 4. H2X. 158 . What is the molar mass of the acid? Part B Prepare your notebook for the lab.0 mL of water. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. 2. Once grams and moles of the diprotic acid are known. simply divide its NaOH volume by 2 to confirm the first equivalence point. the volume of NaOH added at the second equivalence point is exactly twice that of the first equivalence point (see Equations 3 and 5). use the ratio: 1 mole H2X / 2 mol NaOH 159 . NaOH. The first is somewhat easier. You may use either the first or second equivalence point to calculate molecular weight or both. If the second equivalence point is more clearly defined on the titration curve. all H+ ions from the Figure 1 first dissociation have reacted with NaOH base. A known mass of a diprotic acid is titrated with NaOH solution of known concentration. H2CO3. At + the second equivalence point. and sodium hydroxide base. titration curves of diprotic acids can have two equivalence points. Molecular weight determination is a common way of identifying an unknown substance in chemistry. The primary purpose of this experiment is to identify an unknown diprotic acid by finding its molecular weight. because moles of NaOH are equal to moles of H2X (see Equation 3). Weighing the original sample of acid will tell you its mass in grams. Molecular weight (or molar mass) is found in g/mole of the diprotic acid.Lab: Titration of a Diprotic Acid: Identifying an Unknown A diprotic acid is an acid that yields two H+ ions per acid molecule. H2X. in g/mole. or from Equation 5. however. molecular weight can be calculated. Therefore. The volume and the concentration of NaOH titrant are used to calculate moles of NaOH. H2SO4. all H ions from both reactions have reacted (twice as many as at the first equivalence point). as shown in Figure 1. The equations for the acid-base reactions occurring between a diprotic acid. Moles of unknown acid equal moles of NaOH at the first equivalence point (see Equation 3). and carbonic acid. Examples of diprotic acids are sulfuric acid. are from the beginning to the first equivalence point: NaHX + H2O (3) H2X + NaOH from the first to the second equivalence point: pH Na2X + H2O (4) NaHX + NaOH from the beginning of the reaction through the second equivalence point (net reaction): Na2X + 2 H2O (5) H2X + 2 NaOH 1st Equivalence Point 2nd Equivalence Point V olume NaOH At the first equivalence point. Moles can be determined from the volume of NaOH titrant needed to reach the first equivalence point. A diprotic acid dissociates in water in two stages: H+(aq) + HX–(aq) (1) H2X(aq) – H+(aq) + X2 (aq) (2) HX–(aq) Because of the successive dissociations. b. See the stockroom if needed.Phenol red Hazards toxic corrosive flammable. Chemical Safety Information Diprotic Acid Titration Chemical Unknown acid Sodium hydroxide Indicator Mixture . Check the pH probe for breakage.Ethanol .Equipment Information Each bin should contain: 1 – pH probe After use. Refill pH probe with storage solution as needed. secure cord as shown Notes: a. toxic none toxic 160 .Bromocresol green . Weigh out about 0. Be sure to record the actual buret reading. Use a utility clamp to suspend a pH Sensor on a ring stand as shown here.001 g in the data table in your lab notebook. rinse the electrode probe thoroughly with distilled water and dry them gently with a Kim-wipe tissue. Dispose of the waste solution from this step in the waste jar as directed by your teacher. Use a utility clamp to attach the buret to the ring stand.00 mL level of the buret. Be very cautious about bumping the electrode bulb on the bottom of the container containing a test solution. Record the precise concentration of the NaOH solution in the data table in your lab notebook. Get the stir bar spinning rapidly but smoothly and leave it on. Avoid spilling it on your skin or clothing.1 M NaOH solution in a 250 mL beaker.00 mL level of the buret. Acids can harm your eyes. Drain a small amount of NaOH solution into a waste beaker so it fills the buret tip and leaves the NaOH close to (but below) the 0. After the measurement. and respiratory tract. skin. ALL BURET READINGS NEED TO BE RECORDED TO TWO DECIMAL PLACES. YOU BUY IT. CAUTION: Sodium hydroxide solution is caustic. Obtain and wear goggles. IF YOU BREAK IT. Rinse the tip of the electrode in distilled water. Position the pH Sensor in the diprotic acid solution and adjust its position toward the outside of the beaker so it will be easier to stir the solution with a magnetic stir bar without striking the sensor. Handle it with care. return it to its storage buffer. PROCEDURE 1. Obtain approximately 60 mL of ~0.MATERIALS Vernier pH Sensor Stir plate 50 mL buret Magnetic stir bar Note about the pH sensor: The pH meter is a device used to measure the electrical potential of a particular process. CAUTION: Handle the solid acid and its solution with care. Place the electrode into one of the 161 . Fill the buret a little above the 0. Add 6 drops of the indicator mixture. Before making a pH measurement. When finished. The glass bulb is part of an electrode that is responsive to the hydronium ion activity of the test solution. 6.1 M NaOH solution. Record the mass to the nearest 0. 4.120 g of the unknown diprotic acid on a piece of weighing paper. Calibrate the pH probe following this procedure: Use the 2-point calibration option of the Vernier data-collection program. Connect the pH Sensor to the computer interface. 7. again rinse it. Obtain a 50 mL buret and rinse the buret with a few mL of the ~0. even though protected by a plastic shield. Prepare the computer for data collection by opening the file “Exp 24a Acid Base Titration” from the Chemistry w/ Computer folder of Logger Pro. The indicator mixture contains a mixture of bromocresol green and phenol red. 5. 4. 2. Transfer the unknown acid to a 250 mL beaker and dissolve in 100 mL of distilled water. The glass electrode is fragile. f. buffer solutions. d. Then print a copy of the graph. When the voltage reading displayed on the computer or calculator screen stabilizes. You are now ready to begin the titration. This process goes faster if one person manipulates and reads the buret while another person operates the computer and enters buret readings. e. click .20 units and enter the buret reading after each addition. It is now ready to be placed in the sample to be measured. Additionally. Continue in this manner until a pH of 7. again add larger increments that raise the pH by about 0. Dispose of the beaker contents in the 10. click waste jar as directed by your TA.5 is reached. Proceed in this manner until the pH is 3. When the displayed voltage stabilizes. You have now saved the second data pair for the experiment. Enter the buret reading after each increment. c. Continue in this manner until you reach a pH of 11. When pH 7. change back to 2-drop increments. Rinse the electrode with distilled water.20 units. When you have finished collecting data. 9. 8. Press ENTER. Pay special attention to the color of the solution and be sure to write down the pH when color changes happen. Enter the buret reading after each increment.5 is reached. g. For the next calibration point. When pH 10 is reached.01 mL. Add enough NaOH to raise the pH by about 0. In the edit box.5 is reached. to the nearest 0. rinse the electrode and place it into a second buffer solution. enter the pH value on the bottle. and press ENTER to store the first data pair for this experiment. change to 2-drop increments. Continue adding NaOH solution in increments that raise the pH about 0. a. Once the pH has stabilized. type in the initial buret reading. type the current buret reading. After pH 4. again click . b.20 units and enter the buret reading after each addition. again add larger increments that raise the pH by 0. When the pH stabilizes. 162 . Print a copy of the table. enter the pH value on the bottle.5. . Before adding NaOH titrant. click and monitor the pH for 5-10 seconds.5 is reached. note any change in the color of the solution. Enter the buret reading after each increment. When pH 3. In the edit box.20 units. depending on which equivalence point you selected in Step 1. Determine the percent error for your molecular weight value in Step 6. Using the mass of diprotic acid you measured out in Step 1 of the procedure. 163 . 10. 7. Examine the data to find the largest increase in pH values during the 2-drop additions of NaOH. Then find the NaOH volume after the largest pH jump. Find the NaOH volume just before this jump. From the following list of five diprotic acids. Calculate the number of moles of NaOH used at the equivalence point you selected in Step 1. add the two NaOH volumes determined in Step 2. On your printed graph. Determine the number of moles of the diprotic acid. On your printed graph. Note: Dividing or multiplying the other equivalence point volume by two may help you confirm that you have selected the correct two data pairs in this step. H2X. 2. use your graph and data table to determine the volume of NaOH titrant used.44 = 12. identify your unknown diprotic acid. For example: 12. the two-drop increments near the equivalence points frequently result in larger increases in pH (a steeper slope) at one equivalence point than the other. For the alternate equivalence point (the one you did not use in Step 1). Use your graph and data table to determine the volume of NaOH titrant used for the equivalence point you selected in Step 1. To do so. Indicate the more clearly defined equivalence point (first or second) in your data table. calculate the molecular weight of the diprotic acid.PROCESSING THE DATA 1. Find the NaOH volume just before and after this jump. To do this. Underline both of these data pairs on the printed data table and record them in the Data and Calculations table. 3. 11. Use Equation 3 or Equation 5 to obtain the ratio of moles of H2X to moles of NaOH. Determine the volume of NaOH added at the equivalence point you selected in Step 1. 5. examine the data to find the largest increase in pH values during the 2-drop additions of NaOH. clearly specify the position of the equivalence point volumes you determined in Steps 3 and 10. Identify both of these data pairs and record them. and divide by two. Diprotic Acid Oxalic Acid Malonic Acid Maleic Acid Malic Acid Tartaric Acid Formula H2C2O4 H2C3H2O4 H2C4H2O4 H2C4H4O5 H2C4H4O6 Molecular weight 90 104 116 134 150 8.34 + 12. using the same method you used in Step 3. in g/mol. Specify the NaOH volume of each equivalence point on the horizontal axis of the graph.39 mL 2 4. one of the two equivalence points is usually more clearly defined than the other. using dotted reference lines like those in Figure 1. 6. Determine the volume of NaOH added at the alternate equivalence point. 9. 1 pH unit. is equal to the pKa2 value. The Ka expressions for the first and second dissociations. [HX ] for [X2 ] in the Ka2 expression. logKa2 = log[H+] pK a2 pH pK a1 1 1st EP 2 2nd EP Volume NaOH Figure 2 1. Determine the pH values on the vertical axis that correspond to each of these volumes. is equal to the pKa1 value. Ka1 and Ka2. Determine the precise NaOH volume for the first half-titration point using one-half of the first equivalence point volume (determined in Step 2 or Step 9 of Processing the Data). respectively. The first half-titration point volume can be found by dividing the first equivalence point volume by two. logKa1 = log[H+] Thus. Similarly. 2. use their pH values to confirm your estimates of pKa1 and pKa2 from the graph. are: [H+][HX-] [H+][X2-] Ka1 = [H X] Ka2 = [HX-] 2 The first half-titration point occurs when one-half of the H+ ions in the first dissociation have – been titrated with NaOH. (Note: See if there are volume values in your data table similar to either of the half-titration volumes in Step 1. The second half-titration volume (Point 2 in Figure 2) is midway between the first and second equivalence point volumes (1st EP and 2nd EP). it is possible to determine the acid dissociation constants. draw reference lines similar to those shown in Figure 2. 164 . Similarly. calculate the Ka1 and Ka2 values for the two dissociations of the diprotic acid. the pH value at the second titration point. Start with the first half-titration point volume (Point 1) and the second half-titration point volume (Point 2). On your graph of the titration curve. the pH value at the first half-titration volume.Extension Using a half-titration method. from Equations 1 and 2. so that [H2X] = [HX ].– so that – 2– [HX ] = [X ].) 3. Then determine the precise NaOH volume of the second half-titration point halfway between the first and second equivalence points. the following are obtained: Ka1 = [H+] Ka2 = [H+] Taking the base-ten log of both sides of each equation. Estimate these two pH values to the nearest 0. the second half-titration point occurs when– one-half of the H+ ions in the second–dissociation have been titrated with NaOH. Use the method described below to determine the Ka1 and Ka2 values for the diprotic acid you identified in this experiment. If so. for the two dissociations of the diprotic acid in this experiment. These values are the pKa1 and pKa2 values. Point 1 in Figure 2. From the pKa1 and pKa2 values you obtained in the previous step. Substituting [H2X] for [HX ] in the Ka1 expression and. 165 .EQUIVALENCE POINT DETERMINATION: Another Method An alternate way of determining the precise equivalence point of the titration is to take the first and second derivatives of the pH-volume data. View the first-derivative graph (pH/vol) by clicking the on the vertical-axis label (pH). and choosing Second Derivative. You may need to autoscale the new graph by clicking the Autoscale button. 1. In Method 2. and to the volume where the second derivative equals zero on the second derivative plot. and choose First Derivative. . view the second-derivative on Page 3 by clicking on the Next Page button. 2. . View the second-derivative graph (2pH/vol2) by clicking on the vertical-axis label. The equivalence point volume corresponds to the peak (maximum) value of the first derivative plot. the identity of the acid and the percent error in your calculated molar mass for both equivalence points AND for both methods of determining equivalence point Include the determined values of Ka1 and Ka2 for your acid Attach your calculations to the end of your report Discussion: Discuss the experiment and any possible sources of error. relevant conceptual background. and general equations (titration. diprotic acids and Ka) Explain how these concepts can help to determine the identity of some unknown acid Include equations where necessary Data: Include in your report the recorded concentration of base used for the titration. Question: When the pH of the solution equals the pKa of an indicator. Estimate the pKa of both indicators (bromocresol green is the indicator that made the transition in the acidic region of the titration). Tabulate your results for the calculated moles of acid. Subm it you r report on tim e and to your TA in the d ropbox on D2L. the first derivative plot and the second derivative plot. Which data analysis method (using the indicators or using the Venier pH data) do you feel gave you better (or more easy to interpret) results? Justify your answer. titration curves. 166 . summarize the relevant volume and pH data collected during your titration. the mass of the unknown used. include only the data for the initial reading. it is summarized in the titration curve you will to cite and attach). How sure are you in the identification of your unknown? Use both your Ka values and molecular weight to identify the appropriate acid. at each ½ equivalence point and at each equivalence point (do not turn in your raw data. the solution will have an intermediate color. equivalence point.Titration of a Diprotic Acid Lab Report: You will write a complete lab report for this lab which will contain the following sections: Abstract: State the molar mass and which of the unknown diprotic acids you were given and what you found for the two Ka values Introduction: Include a statement of purpose for this experiment. the molar mass of the acid. You should look up the Ka values of all the possible acids to find which acid best matches your unknown. Additionally. Results: Include in your report a copy of the titration curve. 4. Part B Prepare your notebook for the lab.0 mL of 1. The Ka of acetic acid is 1. The Ka of acetic acid is 1.0 M acetic acid to prepare a pH 4 buffer. This includes stating the purpose of the experiment.1 M acetic acid to prepare a pH 4 buffer. Write a reaction to show how a sodium acetate/acetic acid buffer would respond to a small amount of added strong acid.8 10–5 2. Buffer A: Calculate the mass of solid sodium acetate required to mix with 50. Write a reaction to show how a sodium acetate/acetic acid buffer would respond to a small amount of added strong base.Pre-lab: Buffers Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. At the start of your lab. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. Buffer B: Calculate the mass of solid sodium acetate required to mix with 50.0 mL of 0. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection.8 10–5 3. 167 . pH pK a log [ NO 2 ] [HNO2 ] The pH range in which a buffer solution is effective is generally considered to be ±1 of the pKa corresponding to ten-fold excesses of either the acid over the conjugate base. Determine the buffer capacity of the prepared buffers. A buffer’s function is to absorb small amounts of acids (H+ or H3O+ ions) or bases (OH– ions) so that the pH of the system changes by a smaller amount than it would with most other solutions. a comparable amount of the conjugate base is added. 168 . the Henderson-Hasselbalch equation is shown below. HNO2. OBJECTIVES In this experiment. By design. resulting in a solution that is similar to the original one. A variation of the equilibrium expression above. You will then prepare the buffers and test their buffer capacities by adding solutions of NaOH and HCl while monitoring the pH. [H ][ NO -2 ] Ka [HNO2 ] To prepare a buffer system with nitrous acid. a buffer can be prepared with nitrous acid. The resulting system is a mixture of HNO2 and NO2– ions. Blood plasma. called the Henderson-Hasselbalch equation.6. is a bicarbonate buffer that keeps the pH of blood between 7. such as sodium nitrite (NaNO2). and in each case the product will be one of the original conjugates. For example. you will use the Henderson-Hasselbalch equation to determine the amount of acetic acid and sodium acetate needed to prepare two acidic buffer solutions. For our nitrous acid/sodium nitrate buffer example. The weak acid establishes an aqueous equilibrium as shown below. a natural example in humans. or vice versa. HNO2 (aq) ↔ H+ (aq) + NO2– (aq) The equilibrium constant expression is shown below. The nitrous acid molecule will neutralize hydroxide ions and the nitrite ion (the conjugate) will neutralize hydronium ions.2 and 7. buffers are critical.Lab: Buffers A buffer is a mixture of a weak acid and its conjugate base. or a weak base and its conjugate acid. a buffer is an equilibrium system. In many systems. you will Prepare and test two acid buffer solutions. In this experiment. is a useful reference in preparing a buffer solution. Equipment Information Each bin should contain: 1 – pH probe 1 – 50mL graduated cylinder After use. secure cord as shown Notes: a. Chemical Safety Information Buffers Chemical Acetic acid Hydrochloric acid Sodium hydroxide Sodium acetate Hazards corrosive corrosive corrosive none 169 . b. Check the pH probe for breakage. Parafilm is reusable. 2. b. Once the displayed pH reading has stabilized. CAUTION: Sodium hydroxide solution is caustic.30 mL. Calibrate the pH probe following this procedure: Use the 2-point calibration option of the Vernier data-collection program. Connect a pH Sensor to Channel 1 of the Vernier computer interface. You are now ready to test Buffer A. 7. Once you verify the initial pH of the buffer. rinse the electrode and place it into a second buffer solution. You will stir with a stirring rod during the testing. and ring stand (see Figure 1). click . When the voltage reading displayed on the computer or calculator screen stabilizes. Open the file “19 Buffers” from the Advanced Chemistry with Vernier folder. Figure 1 6. It can cause painful burns if it comes in contact with the skin. 5.5 M HCl solution. Handle the hydrochloric acid with care. Rinse and fill the second buret with 0. enter the pH value on the bottle. buret clamps. Rinse the electrode with distilled water. Use a graduated cylinder to measure out 10. you will slowly and carefully add 0. 170 .5 M NaOH solution. In the edit box.0 mL of 0.1 M acetic acid solution. Take an initial pH reading of the buffer solution. For the next calibration point. Add a small amount of the NaOH solution. 3. When the displayed voltage stabilizes. Enter the current buret reading and press ENTER to store the second data pair. Press the ENTER key to store the first data pair. It is now ready to beplaced in the sample to be measured. Start the Logger Pro program on your computer. Weigh out the precise mass of sodium acetate and dissolve it in 50. enter the pH value on the bottle. Avoid spilling it on your skin or clothing. Record the initial pH value. Place the electrode into one of the buffer solutions. Use your calculations from the Pre-Lab Exercise to prepare 50 mL of Buffer A. NOTE: if the initial pH is not within 0.MATERIALS Vernier pH Sensor two 50 mL burets PROCEDURE Part I Prepare and Test Buffer Solution A 1. Rinse and fill one buret with 0. When the pH stabilizes click .0 you should remake the buffer and begin again. Obtain and wear goggles. type the starting volume on the buret.5 M NaOH solution to the buffer solution. a. Set up two burets. 4. Rinse the tip of the electrode in distilled water. 9.3 pH units of 4. up to 0. Click and monitor pH for 5–10 seconds. Connect the interface to the computer using the proper interface cable.0 mL of the Buffer A solution into a 250 mL beaker and add 15 mL of distilled water. 11.c.0 mL sample of the Buffer A solution. using a fresh 10.0 mL of 1. Repeat the necessary steps to test Buffer B in a manner similar to the Part I trials. (The pH will begin to rise rapidly at the equivalence point. repeat using the sodium hydroxide. and passed. e. refill the burets of NaOH and HCl solution. d. Record the volume of HCl that was used to lower the pH of Buffer B by 2 units or until no significant change continues to occur. continue to add the NaOH solution in small increments until you have reached. Use your calculations from the Pre-Lab Exercise to prepare 50 mL of Buffer B. Your goal is to raise the pH of the buffer by precisely 2 pH units. Rinse the pH sensor with distilled water in preparation for the second titration. Print a copy of the first trial. Part II Prepare and Test Buffer Solution B 12. Print a copy of your graph of the titration using the NaOH solution. but either print the data table or copy the data into your lab notebook. If necessary. Carefully add HCl in small increments until the pH of the solution has been lowered by precisely 2 units or no significant change continues to occur. 10. Repeat Steps 7 and 8.0 mL of the Buffer B solution.) Reaching the equivalence point is important for ensuring consistency in your data. the equivalence point of the titration. 13. Continue adding the NaOH solution in small increments that raise the pH consistently and enter the buret reading after each increment. When the pH of the buffer solution is precisely 2 units greater than the initial reading. Use a graduated cylinder to measure out 10. Dispose of the reaction mixture in the waste jar as directed. For the second trial.5 M HCl solution. For the third trial. There is no need to print a copy of the graph. Record the volume of HCl that was used. Weigh out the precise mass of sodium acetate and dissolve it in 50. Click .0 M acetic acid solution. titrate the buffer with 0. 171 . and does not make for a good buffer system in the lab) 3.0 mL each of Buffer B and Buffer C. a. which buffer’s pH would change less? Explain. balanced chemical reactions. NaC2H3O2. Say. and show how buffers react with a strong acid and with a strong base. Data: Include a summary of the collected data in the form of a data table. Buffer capacity has a rather loose definition. Summarize the data to include those data points that are used to calculate the buffer capacity. Using the moles of acid required to lower the pH by 2 and the moles of base required to increase the pH by 2. Results: Include the graph of pH vs. Why were the buffers asymmetric – that is they seemed to handle added sodium hydroxide with relatively small changes in the pH while small quantities of HCl caused the pH to decrease substantially? Subm it you r report on tim e and to your TA in the d ropbox on D2L. that you had prepared a Buffer C.0 mL of 0. 172 . Use your data to determine the buffer capacity of Buffer A and Buffer B and include this in the results section. Keep in mind this will be different for acid and base. the equation below can be used to express buffer capacity: Buffer capacity = moles of acid or base added / change in pH Discussion: Discuss the experiment and any possible sources of error Answer the following questions as part of your discussion: 1. What would be the initial pH of Buffer C? b. If you wanted to carry out an experiment at ‘physiological pH’ (7. 2. in which you mixed 8.203 g of sodium acetate. This does not mean you should attach your raw data to your report. with 100.4) could you suggest an appropriate buffer system? (Hint: The blood buffer system is incredibly complex. If you add 5. yet it is an important property of buffers. relevant conceptual background.Buffers Lab Report: Your lab report should include the following sections: Abstract: Report the buffer capacity (defined below) for each of your buffers Introduction: Include a statement of purpose for this experiment.0 mL of 1.0 M acetic acid. for example.5 M NaOH solution to 20. volume of NaOH added for both buffer A and B. What is the mole ratio between [OH-] and [H+]. The molar solubility of a slightly soluble ionic compound M2X3 is 2. 2. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. M(OH)2 Ksp = 4. Which of the saturated solutions shown in question 4 would have the lowest pH? Part B Prepare your notebook for the lab.8 x 10-6 M.39 x 10-4 c.64 x 10-3 d. Write the molecular balanced chemical equation and the net ionic equation for the reaction between Ca(OH)2(aq) and HCl(aq). M(OH)2 Ksp = 2. M(OH)2 Ksp = 8. Which of the saturated solutions below would have the highest [OH-]? a. M(OH)2 Ksp = 7.Pre-lab: Determination of the Ksp of Calcium Hydroxide Part A Answer the following questions in your lab notebook (be sure to show your work for any calculations): 1.25 x 10-6 b.52 x 10-8 5. between [Ca2+] and [H+]? 3. Determine the value of Ksp. 173 . At the start of your lab. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. This includes stating the purpose of the experiment. 4. All compounds. even highly soluble ones like sodium chloride. OBJECTIVES In this experiment. In solution it completely dissociates into ions as represented by the following balanced chemical equation: – Ca(OH)2 (s) ↔ Ca2+ (aq) + 2OH (aq) The solubility product expression describes.Lab: Determination of the Ksp of Calcium Hydroxide From Advanced Chemistry with Vernier. you will have the necessary information to calculate the Ksp. The mass of dissolved solid can be used to determine the solubility and thus the Ksp of Ca(OH)2. in mathematical terms. have a Ksp. The third method will use gravimetric analysis to determine the quantity of Ca(OH)2 dissolved in a known volume of solution. By determining the molar concentration of dissolved hydroxide ions in the saturated Ca(OH)2 solution and assuming they all come from calcium hydroxide. Ksp = [Ca2+][OH-]2 (Equation 1) The constant that quantifies a substance’s solubility in water is called the solubility product Ksp. Vernier Software & Technology. 2004 INTRODUCTION Calcium hydroxide is a strong base that is sparingly soluble in water. the Ksp of a compound is usually only considered in cases where the compound is slightly soluble and thus the concentration of solvated ions is small. However. The second method will be to measure the pH of the saturated solution and calculate the concentration of the OH-(aq) ions and use that information to determine the Ksp. the equilibrium that is established between the solid substance and its dissolved ions in an aqueous system. The primary objective in this experiment is to test a saturated solution of calcium hydroxide and use your observations and measurements to calculate the Ksp of the compound. The equilibrium expression for calcium hydroxide is shown below. You will do this by three methods: The first will be to determine the concentration of Ca(OH)2 by titrating with a standardized hydrochloric acid solution. you will use three separate methods to determine the Ksp of calcium hydroxide and will compare the ease of use and reliability of the three 174 . e.Equipment Information Each bin should contain: 1 – pH probe Notes: a. d. Allow the beaker to cool before putting it on a balance. Make sure to write your group number on the beaker for gravimetric analysis. b. Chemical Safety Information Ksp of Calcium Hydroxide Chemical Calcium hydroxide Hydrochloric acid Methyl red indicator Ethanol (indicator mixture) Hazards corrosive corrosive environmental hazard flammable. Only TAs will use the oven. The beakers need to be in the oven for ~1 hour. toxic 175 . After use. so do that part first. secure cord as shown 1 – 10mL serological pipet with bulb Do not aspirate liquid into the bulb Check the pH probe for breakage. c. 1. 3.. Start the Logger Pro program on your computer and allow the default program to open. The beaker must be labeled with your section number and your drawer number. It is now ready to be placed in the sample to be measured. 176 . 2. and filter paper as demonstrated by your instructor. Filter your sample of Ca(OH)2 and transfer to a clean beaker. Obtain about 50 mL of a saturated calcium hydroxide solution.Determination by pH 1.g. Method 1 . Obtain and wear goggles. Follow the acid rinse up with a thorough rise with distilled water. 2. When the voltage reading displayed on the computer or calculator screen stabilizes. After the drying period. 4. pH 7). Rinse the tip of the electrode in distilled water. the mass of the calcium hydroxide. When you are done with the filtering apparatus. Method 2 .g. enter a pH value. Calibrate the pH probe following this procedure: Use the 2-point calibration option of the Vernier data-collection program. be sure to clean the flask out with ~5 mL of acid. Place the beaker in the warming oven until the end of the lab period. Set up a ring stand. The acid rinse should be deposited in the waste receptacle. 3. “10”. 2. Place the electrode into one of the buffer solutions (e. enter a pH value. ring. Rinse the electrode with distilled water. When the displayed voltage stabilizes. pH 10). For the next calibration point. Connect the interface to the computer using the proper interface cable. filter funnel. rinse the electrode and place it into a second buffer solution (e. “7”. Using a 25 mL graduated cylinder transfer exactly 20 mL of the filtered solution into a pre-weighed and labeled 250 mL beaker. Retain this solution. determine the mass of the beaker plus dry calcium hydroxide and by difference. Using the remaining 30 mL of the filtered solution in a 100 mL beaker and measure and record the pH of the filtered solution in your lab notebook. Connect a pH Sensor to Channel 1 of the Vernier computer interface.PROCEDURE CAUTION: Calcium hydroxide solution is caustic.. Avoid spilling it on your skin or clothing. 3.Gravimetric Determination (begin one lab period in advance) 1. The solution will turn yellow. 4. Repeat the steps above to titrate a second sample of the filtered Ca(OH)2 solution. semi-rapidly titrate your Ca(OH)2 sample with the HCl until the indicator starts to turns pink. 8. Dispose of the reaction mixture in the labeled waste bottle in the fume hood 9.0500 M HCl solution.Method 3 . 2. Record the initial volume of HCl in your buret to 0. Obtain about 70 mL of 0. Add 3 drops of the methyl red indicator solution. do a third titration. 177 .050 M HCl solution.01 mL as usual. measure exactly 10 mL of the filtered solution from “Method 2” into a 250 mL beaker. Record the final volume of HCl in your buret and determine the volume of HCl used in the titration (by difference). Using a glass stirring rod to stir the solution. 10. 7. 6. If the volume of HCl used in the two titrations differs by more than 10%. 3. Using a 10 mL volumetric pipette. Rinse the buret with ~10 mL of the acid before filling it with the 0.Determination by Titration 1. 5. Connect a buret to the ring stand. Begin adding the HCl dropwise until the whole solution just turns pink in color that persists for 20 seconds. assuming that both ions came from calcium hydroxide. Determination of Ksp through titration. Again use the balanced chemical equation to link the [OH-] to [Ca2+] and assume calcium hydroxide was the only source of both. and the molar mass of calcium hydroxide. Substitute the values into equation 1 to determine the Ksp. Using the concentration and volume of HCl used. Gravimetric determination of Ksp. With the balanced chemical equation defining the dissolution of Ca(OH)2 use the [OH-] to determine the [Ca2+]. the volume of solution used in the analysis. Determine the mass of dissolved solid.PROCESSING THE DATA Calculations: Determination of Ksp using the pH of the solution Use the pH to determine the [OH-]. determine the [OH-] in the saturated solution. Substitute the molar solubility into a version of equation 1 [Hint: use an ICE table] to determine the Ksp. Substitute the concentrations into equation 1 to determine the Ksp. Calculate the solubility of the solution (M) using the mass of dissolved solid. 178 . Construct a data table summarizing the results. provide an average. say so and average the three results to obtain a “best value”. answer the following question as part of your discussion: 1. relevant conceptual background. Choose the best statement for each scenario. a) The buret was inadvertently left wet with water from cleaning. The measured Ksp would be lower than the true value The measured Ksp would be higher than the true value The measured Ksp would be same as the true value b) Using phenolphthalein as the indicator instead of methyl red. Predict how each of the following “mistakes” would affect the value of the measured Ksp and provide reasoning for your response. The measured Ksp would be lower than the true value The measured Ksp would be higher than the true value The measured Ksp would be same as the true value Subm it you r report on tim e and to your TA in the d ropbox on D2L. and give the equations and a short description of the use of equilibrium constant expressions in calculating solubility Data: Include 3 data tables. Speak about which method provided you the most accurate results and why you think that particular method was more accurate and the others were not.Determination of the Ksp of Calcium Hydroxide Lab Report: Your report for this lab should include the following sections: Abstract Give the Ksp for the three methods and suggest which method seems most reliable. Discussion: Use the data in Appendix II of your test book to determine the accepted value for the Ksp. (Or conclude that all three seemed to work – under what conditions would this seem to be a valid conclusion?) In addition. If all three methods produced similar values for Ksp. Calculate your percent error with respect to each result. provide the percent error. Introduction: Include a statement of purpose for this experiment. or if all seem equally valid. 179 . one for each method used to determine the Ksp Results: Calculate the values of Ksp for each method. In either case. What is ΔG? What does the sign of ΔG tell you about the spontaneity of a reaction? 2. At the start of your lab. Do you expect the dissolution of KNO3 to have a positive or negative ΔS? Use Appendix II in your text to calculate ΔSdissolution. 5. The Ksp for Ag2CrO4 is 9.Pre-lab: Thermodynamics of the Solubility of Potassium Nitrate Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. Do you expect the dissolution of KNO3 to be endothermic or exothermic? Use Appendix II in your text to calculate ΔHdissolution. a. b. Calculate the free energy change associated with making a saturated solution of AgCl in water at 75 °C. A researcher wants to make a solution of AgCl and water at 75 °C.0020 M K2CrO4.0 x 10-12. How many grams of AgCl will dissolve in 1. If 200 mL of 0.0 L of water at 75 °C? 3.0050 M AgNO3 is combined with 300 mL of 0. Part B Prepare your notebook for the lab. For solid AgCl at 75°C. will a precipitate form? 4. Ksp = 1.5 x 10-5. 180 . summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. This includes stating the purpose of the experiment. The data collected will be used to determine the Ksp. The enthalpy change will be negative if the dissolution process releases heat. Ksp. A negative value indicates that the process is spontaneous while a positive value denotes a nonspontaneous process. ΔG = -RTlnKsp (Equation 3) By combining Equations 2 and 3. Thermodynamics may be used to understand the energy changes that occur when a salt dissolves in water. The free energy change (ΔG) for a process will indicate if the process is spontaneous as written (reactants going to products). The Gibbs-Helmholtz equation.Lab: Thermodynamics of the Solubility of Potassium Nitrate In this experiment. ΔG = ΔH . At this point the reaction is said to be at equilibrium. KNO3(s) K+(aq) + NO3-(aq) As the concentration of dissolved K+ and NO3. H 1 S lnK sp (equation 4) R T R 181 . enthalpy. for this process is shown in equation (1). The solution is now considered saturated. is a mathematical expression that relates changes in free energy. In aqueous solution potassium nitrate (KNO3) dissociates according to the following reaction. it is possible to derive an equation that relates Ksp and the Kelvin temperature to the values associated with ΔH and ΔS (equation 4). An equilibrium expression. shown in Equation 2. At one set of ion concentrations the rate of dissolution will equal the rate of precipitation. The energy difference between the solid salt and its dissolved ions is known as the enthalpy change (ΔH). (equation 3). also increases. entropy and free energy of dissolution. enthalpy. Ksp = [K+][NO3-] (Equation 1) The value for Ksp is characteristic of each compound and changes with the temperature. KNO3. When a salt dissolves in water it will dissociate into ions. A positive enthalpy change will occur if heat must be added to dissolve the salt in water. and entropy.TΔS (Equation 2) ΔG can also be expressed in terms of Ksp .increases. and the relative disorder of the dissolved ions is an indication of the entropy change (ΔS). The entropy change for a solid salt dissolving in water will always be positive because the dissolved ions possess more disorder than a solid crystalline salt. you will measure the solubility of KNO3 as a function of temperature. the rate at which the ions will recombine into solid potassium nitrate. Chemical Safety Information Thermodynamics of the solubility of potassium nitrate Chemical Potassium nitrate Hazards oxidizer 182 . ΔH and ΔS for the dissolution process Equipment Information Each bin should contain: 1 – 10mL graduated cylinder 1 – temperature probe After use. It should be poured down the drain unless contaminated. we get a line where the slope of the line is –ΔH/R and the yintercept is ΔS/R (R is the ideal gas constant 8. you will Determine the solubility of KNO3 as a function of temperature Use the solubility data to determine the Ksp for the dissolution of KNO3 Use the data and Equations 4 and 2 to calculate ΔG. secure cord as shown 1 – copper stirrer Use for stirring Notes: a.By plotting lnKsp versus 1/T. Hot water bath is not waste. The enthalpy and entropy of dissolution can be determined by evaluating the temperature dependence of Ksp. OBJECTIVES In this experiment.314 J∙mol-1∙K-1). Be careful not to lose solid as you remove the thermometer probe from the cylinder to add the water. Do not leave it in the water any longer than necessary to get the salt into solution. 13.” For easier removal of the KNO3 reheat the solution until the entire solid has re-dissolved and quickly pour it into the waste jar. heat the cylinder containing the salt and water mixture. The solution will appear to be “snowing. 11. Repeat steps 8 . In this case. Add distilled water to 2 mL in your graduated cylinder and observe if the dissolution of KNO3 is an exothermic or endothermic process (feel the outside of the cylinder). 6. Mass ~2 grams of potassium nitrate and record the actual mass in your lab book. allow the sample to cool while slowly stirring. Prepare a hot water bath by heating a halffilled 400 mL beaker on a hot plate 5. While monitoring the temperature. while stirring gently with the temperature probe. 12.5 mL of distilled water.” The solution is considered to be at equilibrium when the first crystals begin to appear. Remove the probe and add 0.11 four more times until a total of 5 sets of data at different concentrations have been recorded. 7. 4. 9. 8.PROCEDURE 1. Start the Logger Pro program on your computer. At higher concentrations the process happens quite quickly. In your hot water bath. 10. 2. 183 . Remove your sample from the hot water bath and record the total volume of the solution (only do this after all the salt has gone into solution and do not forget to remove the temperature probe from the solution when measuring the volume). 3. Allow Logger Pro to open the default program. It should list columns for collecting time and temperature. the water can be evaporated and the potassium nitrate reused. Connect a Temperature Probe to Channel 1 of the Vernier computer interface. Transfer the salt into your 10 mL graduated cylinder. Obtain and wear goggles. Disposal: Place the KNO3 solution into the waste jar labeled “KNO3 Waste. Record the temperature at the point when the first crystals appear. From your graph. Substitute Ksp into equation (3) to calculate ΔG for each data set. Determine the best linear fit of the data using linear regression and record the slope and intercept in your notebook. calculate the reciprocal of each temperature (1/T). and NO3 –(aq) concentrations.PROCESSING THE DATA For each set of data. Use equation (1) to calculate the Ksp for each data set. calculate the molar concentration of KNO3. Using a spreadsheet program (Logger Pro has this capability or you can use Excel) construct a graph with the y-axis being lnKsp and the x-axis being 1/T. Determine the natural logarithm (ln) of the Ksp at each temperature. and use that value to deduce both K+(aq). determine the values for ΔH and ΔS. while ΔS of the reaction can be determined from the y-intercept. After converting all temperatures into Kelvin. Equation 4 shows that ΔH of the reaction can be determined using the slope of the straight line from the graph. 184 . Compare the signs of your experimental thermodynamic values with your expectations from the prelab. Did you expect the dissolving process to be spontaneous? Do your data confirm your hypothesis? b. Was this process endothermic or exothermic? Does this observation match your calculated enthalpy values? c. Be sure to attach your calculations to the end of your report. relevant conceptual background. and give the equations and a brief description of how they are used to obtain the thermodynamic properties in this experiment Data: Include a copy of the graph with the data indicating the slope and y-intercept Results: Calculate the values of solubility. Discuss the experiment and any possible sources of error Answer the following question as part of your discussion: 1. Did you expect this process to result in an increase or decrease in disorder? How does this compare with your calculated entropy values? Subm it you r report on tim e and to your TA in the d ropbox on D2L.Thermodynamics of the Solubility of Potassium Nitrate Lab Report: Your report for this lab should include the following sections: Abstract: Report the ΔH and ΔS values with percent error Introduction: Include a statement of purpose for this experiment. 185 . Discussion: Use the data in Appendix II of your text book to calculate the ΔHº and ΔSº for the dissolution of KNO3. Calculate your percent error with respect to each result: ΔH and ΔS. Ksp and ΔG for each temperature data set. From the graphical analysis of your data calculate values for ΔH and ΔS. a. 25. For each half-reaction. Part B Prepare your notebook for the lab. 5. What is the concentration of hypochlorite in the original bleach solution? Assume the density of the commercial bleach is 1.+ C2O4 2. 186 .56 mL of 0. 4. remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA.00 mL of commercial bleach was diluted to 100. Calculate the average percent by mass of NaClO in the commercial bleach. Define reduction. This includes stating the purpose of the experiment. Use the equations given below in the introduction of this experiment to determine the mole ratio of thiosulfate (S2O32-) to hypochlorite (ClO-).100 M S2O32-. At the start of your lab.Pre-lab: Redox Titration: Analysis of a Commercial Bleach Part A Answer the following questions in your lab notebook (be sure to show your work for any calculations): 1.08 g/mL. Write balanced oxidation and reduction half-reactions for the following redox reaction equations. [OIL RIG/Leo the lion goes Ger?] 2. identify which substance is oxidized and reduced. 222+ Eqn 1: MnO4 + S2O3 S4O6 + Mn Eqn 2: MnO4 . MnO2 + CO2 3.0 mL. Define oxidation.0 mL of the diluted sample was titrated with 4. summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. the complex that forms may not be reversible. The iodine that forms is then titrated with a standard (known concentration) solution of sodium thiosulfate. 2 H+(aq) + ClO-(aq) + 2 I-(aq) Cl-(aq) + I2(aq) + H2O(l) 2. 187 . and some is reduced to the chloride ion. An acidified iodide ion solution is added to hypochlorite ion solution and the iodide is oxidized to iodine (hypochlorite is quantitatively converted to iodine). ClO-. Triiodide is a combination of a neutral I2 molecule and an I. The most common oxidizing agent in bleach is sodium hypochlorite. The disappearance of this blue colored complex is a much more sensitive method of determining the end point. but this is not a very sensitive procedure (it’s hard to tell when the yellow color just disappears). the starch is not added until shortly before the end point is reached (light yellow solution). in which it forms a complex ion called the triiodide ion. 2. that reduce the iodine back to iodide ions. I2(aq) + I-(aq) I3-(aq) 3.Lab: Redox Titration: Analysis of a commercial Bleach Solution Many commercially available consumer products. hypochlorite ions oxidize iodide ions to form iodine. The balanced chemical equation for the process is: Cl2(g) + 2 OH-(aq) ClO-(aq) + Cl-(aq) + H2O(l) The amount of hypochlorite ion present in a solution of bleach can be determined by an oxidation-reduction titration. Cl-. The iodide ion.ion. I-. contain oxidizing agents. NaClO (sometimes written NaOCl). I3-(aq) + 2 S2O32-(aq) 3 I-(aq) + S4O62-(aq) During this last reaction the red-brown color of the triiodide ion fades to yellow and then the color disappears when you reform the clear iodide ion solution. The volume of thiosulfate used during the titration and its known concentration are converted to moles which are related to moles of hypochlorite through reactions 3. Commercial bleach is made by bubbling chlorine gas through a sodium hydroxide solution. and the hypochlorite concentration is calculated. The triiodide ion is yellow in dilute solution and dark red-brown when concentrated. One of the best methods is the iodine-thiosulfate titration procedure. Iodine is only slightly soluble in water. The analysis takes place in a series of steps: 1. It is possible to use the disappearance of the color of the triiodide ion as the method of determining the end point. Some of the chlorine is oxidized to the hypochlorite ion. if the starch is added to a solution that contains a high concentration of iodine. such as bleaches and hair coloring agents. and 1. but it dissolves very well in an aqueous solution of iodide ion. Addition of starch to a solution that contains iodine or triiodide ion forms a reversible blue complex. is easily oxidized by almost any oxidizing agent. In acidic solution. The solution remains strongly basic. The triiodide is titrated with a standard solution of thiosulfate ions. Therefore. However. I2. Then the solution turns blue and you titrate until the blue color goes away. 188 . Add one dropper-full of starch solution. This is a large excess over what is needed. Obtain 70 mL of the sodium thiosulfate solution and use a few mL to rinse your buret. environmental hazard toxic none none corrosive PROCEDURE 1. Obtain about 100 mL of diluted commercial bleach. Chemical Safety Information Bleach Lab Chemical Dilute bleach Potassium iodide Sodium thiosulfate Starch solution Hydrochloric acid Hazards corrosive.complex ions. Swirl to dissolve the KI. 3. slowly and with swirling. Working in a fume hood. add approximately 2 mL of 3 M HCl.5 mL of commercial bleach into 100 mL. meaning that the sample you are working with is 20 times more dilute than the commercial strength. 5. Note that this bleach solution has been made by diluting 3. Add the solid KI and about 25 mL distilled water. The blue color of the starch-iodine complex should appear. 6. two more times.Equipment Information There are no bins this week. beginning with step 2. You will need to scale back up for the final answer to this experiment. Pipet 25. The solution should be dark yellow to red-brown from the presence of the I3. Continue to titrate until one drop of Na2S2O3 solution causes the blue color to disappear. Repeat the titration. 2. Weigh out approximately 1 g solid KI. Record the concentration of the thiosulfate in your lab notebook. Bring the red-brown triiodide solution back to your bench and titrate the iodine with the standard 0. 4. then fill it as usual.10 M sodium thiosulfate solution until the iodine color becomes light yellow.00 mL of the dilute bleach into an Erlenmeyer flask. PROCESSING THE DATA 1. Use the equations given in the introduction to determine the mole ratio of sodium thiosulfate to sodium hypochlorite. 2. Using the volume of sodium thiosulfate needed for titration of 25.00 mL of diluted bleach, calculate the molarity of the diluted bleach (hypochlorite ion). 3. Calculate the molarity of the hypochlorite ion in commercial bleach (undiluted). 4. Assuming that the density of the commercial bleach is 1.08 g/mL, calculate the percent by mass of NaClO in the commercial bleach. 5. Calculate the average percent by mass of NaClO in commercial bleach from your three trials. 6. Read the label of the commercial bleach to find the percent by mass NaClO that is reported. Calculate the percent error of your value, assuming that the label value is correct. 189 Analysis of a Commercial Bleach Solution Lab Report: Your lab report should include the following sections: Abstract: Be sure to provide a percent by mass of sodium hypochlorite in bleach and the percent error in your measurement Introduction: Include a statement of purpose for this experiment, relevant conceptual background, and include the relevant reactions and information as to how the concentration of sodium hypochlorite is determined Data: Include the recorded concentration of the thiosulfate standard used for the titration, the labeled value of the concentration of commercial bleach and any data tables necessary for the completion of the lab. Results: Tabulate your results for the calculated concentration of bleach, its labeled value and the percent error. Attach your calculations to the end of your report. Discussion: Discuss the experiment and any possible sources of error. Subm it you r report on tim e and to your TA in the d ropbox on D2L. 190 Pre-lab: Synthesis of Acetaminophen Part A Answer the following questions in your lab notebook (be sure to show your work for any calculations): 1. Write the balanced chemical equation for the synthesis of acetaminophen from paminophenol and acetic anhydride. 2. Starting with 1.5 grams of p-aminophenol and an excess of acetic anhydride, calculate the theoretical yield of acetaminophen in grams. It will be necessary to look up the molecular formulas of both the limiting reactant and the product (online?). Part B Prepare your notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted-list format (be sure to include space for observations) and preparing any tables necessary for data collection. At the start of your lab, remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA. 191 Synthesis of Acetaminophen Acetaminophen (N-(4-hydroxyphenyl)ethanamide) is a relatively simple organic compound that is a common over-the-counter pain reliever and fever reducer (Figure 1). Acetaminophen is generally considered a safe medication in the U.S. although overdoses are relatively common and can cause fatal liver damage. Organic compounds such as Acetaminophen are generally classified by their functional groups. Acetaminophen consists of a benzene ring core that has a hydroxyl functional group (-OH) attached to one side and an amide functional group (Figure 2) on the opposite side (referred to as the para position see Figure 1). Figure 1. The molecular structure of Acetaminophen. The apex and juncture of each line represents a carbon atom potentially bound with various numbers of hydrogen to give each carbon a total of four bonds. Figure 2. The structure of an amide. R, R’ and R” represent various other organic groups or carbon chains. Find the amide group in Figure 1. The synthesis of acetaminophen is performed by reacting p-aminophenol with acetic anhydride. A byproduct of this reaction is acetic acid as shown Figure 3. In this lab, you will start with 1.5 grams of p-aminophenol to synthesize the crude acetaminophen product. The crude product will be purified through a common technique called recrystallization. Recrystallization involves dissolving a crude product in a minimal amount of hot solvent. Once the solution cools, a more pure form of your product will recrystallize (precipitate) out of solution. One way of assessing the purity of a substance is through Thin Layer Chromatography (TLC). TLC is a method that can be used to separate nonvolatile components of a mixture. In this process, a sample is applied to a “stationary phase” (filter paper, coated plastic…). A solvent, the “mobile phase” is allowed to pass along the stationary phase through the sample. Depending on the relative solubility of the mixture components, different substances will travel further along the stationary phase than others, resulting in separation of the components of the mixture. Solutes with greater affinity for the mobile phase spend longer in the mobile phase and will thus move faster than solutes that prefer the stationary phase. A formula can be used to calculate the relative movement of the parts of the mixture by the following equation: 192 rinse well with cold water. 193 . move to an area where fresh air is available. move to an area where fresh air is available. symbolized Rf. or variations in the paper or solvent. Safety Precautions: Wear safety glasses or goggles at all times in the laboratory. rinse well with cold water. If the vapors are inhaled. Avoid skin contact and inhalation of the vapors. The distance traveled by one compound is measured from its point of origin to the center of the spot. You will be comparing your final acetaminophen produce with the raw unpurified sample and a given sample of pure unreacted p-aminophenol. The distance traveled by the solvent is measured from the point of origin to the highest point that mobile phase traveled. Phosphoric acid is corrosive. Figure 3. changes in humidity. In the event of skin contact. Their relative values should remain constant. Avoid skin contact. NOTE: Don't use your acetaminophen for a headache! Its purity is not assured. Acetic anhydride is corrosive and its vapor is irritating to the respiratory system. Synthesis reaction of acetaminophen. rinse well with cold water. In the event of skin contact. If the vapors are inhaled. p-aminophenol is harmful by inhalation and by contact with the skin. In the event of skin contact. Rf values may change from day to day due to temperature differences.Rf Distance traveled by thesolute Distance traveled by thesolvent The relative movement of the components is called the retention factor. Chemical Safety Information Synthesis of Acetaminophen Chemical P-aminophenol Ethanol Phosphoric acid ethyl acetate hexanes Acetic Anhydride Acetaminophen Hazards toxic. corrosive. toxic corrosive flammable. The solvent used will degrade the plastic. Phosphoric acid is very corrosive. toxic flammable. health and environmental hazard 194 . toxic toxic. health and environmental hazard flammable. Do not put standards in the well plate. toxic. b.Equipment Information Each bin should contain: 1 – well plate Notes: a. health and environmental hazard flammable. Remove the beaker from the heat and allow the solution to cool. 2. This is done by dipping the toothpick in your sample and quickly tapping the toothpick to the corresponding spot on the TLC sheet. your crude product and your purified product in 1 mL of ethanol. put the solution on ice. Place it in a 125-mL Erlenmeyer flask. When crystals begin to appear. Move to the hood once all of the solid has dissolved. Place the crude acetaminophen crystals in a clean 150 mL beaker. Collect the crystals by vacuum filtration using a Buchner funnel. If the solution starts to boil and undissolved solid still remains. 5. If crystallization does not occur. while swirling the flask add 2 mL of the acetic anhydride. Place the mixture on ice and allow the solid to crystalize for 20 minutes. You should have three marks (one for p-aminophenol. In separate clean test tubes. it may be necessary to scratch the side of the flask with a stir rod. Collect the crystals on a pre-weighed piece of filter paper using the Buchner funnel.Procedures: Synthesis of Acetaminophen 1. Wash the crystals twice on the filter with 5 mL portions of ice cold water. Recrystallization of the Acetaminophen 1. it may be necessary to scratch the side of the flask with a stir rod or add a small seed crystal of pure acetaminophen (ask your TA). add more water.5 g of p-aminophenol. Gently heat to dissolve. Weigh out 2. Wash the product on the filter paper with two 5 mL portions of ice cold water. 4. It may be necessary to dry for a few more minutes. 3. Utilize a clean toothpick to transfer your dissolved samples on to the TLC plate. 2. 3. Weigh the product. TLC of the Acetaminophen Sample 1. Allow the solvent to dry and repeat the transfer procedure 15 times. The solution must not go above the 2 cm mark on the TLC plate or touch the spots when the paper is placed in the beaker. Place about 1 cm of aqueous ethyl acetate mobile phase in a 250 ml beaker. 3. if it still appears wet. dissolve a small spatula full (less than the ½ the size of a small pea) of p-aminophenol. Allow the purified product to dry for 10 minutes under vacuum. 2. 6. If crystallization does not occur. 4. Make marks along this line at approximately 1 cm intervals. carefully draw a line roughly 2 cm from the bottom of the chromatography plate. 7. a few mL at a time until the solid dissolves. Record the final dry weight in your lab notebook for the calculation of % yield. Working in the hood. Add 20 mL of distilled water and heat on a hot plate until all of the solid has dissolved. Using a pencil. 195 . Allow this to react for 5 minutes. 5. cooling for 20 minutes. Add 25 mL of water and 25 drops of 85% H3PO4. Label the spots on the bottom line to indicate the identity of each sample 4. one for the raw sample and one for the purified). When the mobile phase comes to within 1 cm of the top of the TLC plate. 8. The solvent will wick up the plate.6. Illuminate the TLC plate with the handheld UV lamp and circle the spots as shown in the diagram on the right. Carefully prop the chromatography paper in the beaker. 9. remove the TLC plate from the beaker. Immediately mark the position solvent front with the pencil. 7. 196 . Allow the TLC plate to dry in the hood. Synthesis of acetaminophen Worksheet (no lab report) _________________ g Mass of p-aminophenol _________________ g Theoretical yield of acetaminophen (show your work below) _________________ g Mass of pure product _________________ % Percent yield (show your work below) Roughly sketch what the TLC plate looks like when illuminated with UV light. Solvent front Origin 197 . is there a difference between your raw and purified sample? Explain. During the crystallization of acetaminophen. Com plete the above pages u sing the Microsoft version of this file that is available for d ow nload on the lab D2L page. what was the purpose of cooling the sample in an ice bath? Why should you use a minimum amount of hot solvent to dissolve the raw product during the recrystallization step? There is not a formal lab report for this lab. subm it the w orksheet on tim e and to you r TA in the d ropbox on D2L. Once the w orksheet is com plete. 198 .According to your TLC. listing the metals in order of their reactivity. an online resource for learning chemistry. you will identify the unknown. How many electrons are transferred Yes / No ___ 3.php The above link will lead you through a series of activities aimed at improving your understanding of Galvanic Cells. Is there a reaction? (circle the correct response) 2.Pre-lab: Electrochemistry: Galvanic Cells and the Nernst Equation Name ___________________ TA’s Name_____________________ Below is a lab activity that is based on a series of virtual lab exercises and videos that have been put together by The ChemCollective. comparing its reduction potential to a standard list. click on the link in the bottom right of the page to progress to the next page. After watching the video “Zinc strip in copper nitrate solution”. using the Nernst equation. Step 1: .Investigating redox reactions of some metals and solutions of metal salts Activity: Investigating redox reactions. click on the link labeled “start” just below the drawing of the pencil tip. and reading the instructions. To begin.org/chem/electrochem/index. Answer the below questions for the portion of the activity in which Sn(s) is placed in AgNO3(aq) 1. Write the balanced redox reaction for the combination of AgNO3(aq) and Sn(s) Once the 3x3 square is complete. Once finished. you will determine the concentration of an unknown solution. Introduction Please read the introduction page. you will determine the standard reduction potential of an unknown metal. please visit the below website: http://chemcollective. First you will take a series of observations to determine if process is spontaneous. you will construct a series of virtual galvanic cells and use those to power a stopwatch. you will create a situation in which the cells are not in the standard condition and measure the cell potential. 199 . Third. Second. Based upon those observations. click on the link in the bottom right of the page to progress to the next page. you will create an activity series. Finally. Follow the direction to complete the 3x3 grid. Step 1: . 1. Will Sn react with Ni2+? (circle the correct response) Yes / No Yes / No Why? Why? Once finished. which is the most reactive?________ 2.Reduction tendencies of metal ions Read the description and answer the below questions. click on the link in the bottom right of the page to progress to the next page.Practice with redox reactions 1. Will Cu react with Ni2+ ? (circle the correct response) 2. 200 . click on the link in the bottom right of the page to progress to the next page. Step 2: Explaining the electron transfer process Read the description and watch the two videos explaining Galvanic Cells Once finished. List the metal ions from the lowest to highest tendency to undergo reduction Once finished. Of the five metal ions in the list.Step 1: . click on the link in the bottom right of the page to progress to the next page. | || | 1. Step 2: Practice with electrochemical cells In the Mg/Zn Carrou cell that is demonstrated and already set up write down the correct cell diagram. Fill in the blanks on the diagram with the correct terms or phrases. What is the color of the wire that leads to the oxidation half cell (this is the anode) in the above diagram? 3. Is the reduction or oxidation half cell written on the left side of the cell diagram? 2.Step 2: The Electrochemical Cell Read the description and watch the video explaining an electrochemical cell The following is an electrochemical cell diagram for the reaction shown in above video: Zn(s) + Cu2+(aq) --> Zn2+(aq) + Cu(s). Finish this page off by reading the description of Carrou cells. What color of wire leads to the reduction half cell (this is the cathode) in the above diagram? 201 . click on the link in the bottom right of the page to progress to the next page. Once finished. Ag |Ag+ Cu |Cu2+ Sn | Sn2+ Ag+ | Ag Cu2+| Cu Sn2+| Sn Once finished. with units. click on the link in the bottom right of the page to progress to the next page. the voltage will be positive)? 5. Step 3: Measuring cell potentials Please read the description and watch the video showing how the voltage is actually measured using the Carrou cell. Once finished. What is the half reaction that is occurring at the cathode in the above cell? 6. click on the link in the bottom right of the page to progress to the next page. construct your own Carrou cell for the following cell Zn|Zn2+||Ag+|Ag. in the below table.Now. 202 . click on the link in the bottom right of the page to progress to the next page. Construct a series of Carrou cells to determine the standard cell potential for the listed combinations. Is this reduction or oxidation (circle the correct response)? Once finished. Please input the voltages. Please remember that in order for the Carrou cell to function. Step 2: Powering a stopwatch Please watch the video showing that galvanic cells can be used to do useful work Once finished. when correctly set up to be a galvanic cell. Step 3: Calculating cell potentials Please read the description to review how one can calculate the cell potential using a set of standard reduction potentials. click on the link in the bottom right of the page to progress to the next page. What is the voltage of the Zn|Zn2+||Ag+|Ag voltaic Carrou cell (remember. a piece of paper must connect the cells to each other through the well containing KNO3 4. Referencing the linked table of standard reduction potentials. What is the balanced half-cell reaction corresponding to the reduction of metal X? 2. you will measure the cell potential of a Carrou cell containing an a half cell of unknown composition. Once finished. What cell did you use for comparison (write the balanced half reaction)? What was the voltage listed on the volt meter? 3. Using the known half cell standard reduction potential. try to identify the unknown metal. 203 . http://hyperphysics. What value did you obtain for the half-cell reduction potential of metal X? 4.gsu.phy-astr.edu/hbase/tables/electpot. Step 3: Applying standard cell potentials Please read the description to review the table of standard reduction potentials is arranged the way it is. you will determine the standard reduction potential of the unknown and identify it based upon the table of standard reduction potentials that is linked below. click on the link in the bottom right of the page to progress to the next page.Step 3: Practice with standard cell potentials In this activity.html Read the description and create one or two Carrou cells using the unknown as one of the half cells 1. The half cell should now show the word “Diluted” in it. a. what happens to the concentration of the various ions? 1.Step 4: Cells in non-standard conditions Please read the description about non-standard cells. Construct a standard cell. The concentration of Cu2+ will increase / decrease (circle the best answer) Step 4: Practice with cells in non-standard conditions Please read the directions with respect to determining the effect of changing concentrations on cell potential. What is the new voltage? V a. The concentration of Zn2+ will increase / decrease (circle the best answer) 2. Zn|Zn2+||Cu2+ |Cu. In the voltaic cell. as the cell runs. for both reactions. all solutions at 1 M. Reaction 1: Sn| Sn2+(diluted)||Cu2+|Cu ____________ b. 1. Reaction 1: Sn(s) + Cu2+(aq) --> Sn2+(aq) + Cu(s) ____________ b. Reaction 2: Zn| Zn2+ ||Sn2+(diluted)|Sn ____________ V 204 . and determine their voltage. Add water to dilute by dragging the beaker of water from the bottom right of the screen over the cell that you want to dilute. Reaction 2: Zn(s) + Sn2+(aq) --> Zn2+(aq) + Sn(s) ____________ V V 2. Sn| Sn2+||Cu2+|Cu and Zn| Zn2+ ||Sn2+|Sn. in the other it decreases. 4. Show your work. put your name and your TA’s name on a hardcopy of this document. determine the diluted concentration for both cells. Using the Nernst equation. Briefly speculate why. In one reaction the voltage increases. Sn| Sn2+(diluted)||Cu2+|Cu Zn| Zn2+ ||Sn2+(diluted)|Sn Once finished. You must submit your Voltaic Cells results to your TA by your normally scheduled lab period on Week 10. 205 .3.