Chemistry 229 General Chemistry Lab 3
Spring 2012 Lab Coordinators: Dr. Eric Sheagley
[email protected] Dr. Gwen Shusterman
[email protected]
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Table of Content Chemistry 228 Lab Syllabus........................................................................................................... 3
Schedule – Spring 2011 .................................................................Error! Bookmark not defined.
Grading Criteria .............................................................................................................................. 4
Laboratory Safety Rules and Procedures........................................................................................ 5
Keeping a Lab Notebook ................................................................................................................ 9
Structure for your Lab Notebooks: ............................................................................................... 10
An example of a prepared notebook follows. ............................................................................... 11
Report Guidelines ......................................................................................................................... 14
Example Lab Report ..................................................................................................................... 19
General Chemistry Lab Report Checklist ..................................................................................... 22
Acid Rain ...................................................................................................................................... 24
Lab: Acid Rain......................................................................................................................... 25
Acid Rain Lab Report: .............................................................................................................. 28
Acid Dissociation Constant, Ka .................................................................................................... 29
Acid Dissociation Constant, Ka ................................................................................................ 30
Acid Dissociation Lab Report: ................................................................................................. 32
Titration of a Diprotic Acid: Identifying an Unknown................................................................. 33
Lab: Titration of a Diprotic Acid: Identifying an Unknown ................................................. 34
PROCESSING THE DATA ..................................................................................................... 36
EQUIVALENCE POINT DETERMINATION: An Alternate Method .................................. 37
Extension .................................................................................................................................. 38
Titration of a Diprotic Acid Report: ......................................................................................... 39
Buffers .......................................................................................................................................... 40
Lab: Buffers ............................................................................................................................. 41
Buffer Report: ........................................................................................................................... 44
Determination of the Ksp of Calcium Hydroxide.......................................................................... 45
Lab: Determination of the Ksp of Calcium Hydroxide............................................................. 46
Determination of the Ksp of Calcium Hydroxide Lab Report:.................................................. 49
Thermodynamics of the Solubility of Potassium Nitrate.............................................................. 50
Lab: Thermodynamics of the Solubility of Potassium Nitrate ................................................ 51
Thermodynamics of the Solubility of Potassium Nitrate Lab Report: ..................................... 54
Redox Titration: Analysis of a Commercial Bleach ..................................................................... 55
Lab: Redox Titration: Analysis of a commercial Bleach Solution .......................................... 56
Analysis of a Commercial Bleach Solution:............................................................................. 58
Pre Lab: Electrochemistry: Voltaic Cells ..................................................................................... 59
Lab 6: Electrochemistry: Voltaic Cells.................................................................................... 60
Voltaic Cells Lab Report: ......................................................................................................... 63
Synthesis of Acetaminophen ........................................................................................................ 65
Synthesis of Acetaminophen .................................................................................................... 66
Synthesis of acetaminophen worksheet .................................................................................... 69
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Chemistry 228 Lab Syllabus Lab Packet: All printed material for this lab will be available on Blackboard OR may be purchased at Smart Copy (1915 SW 6th Avenue). Prelab Exercises: Prelab instructions are included in the lab packet. You should answer any questions presented and prepare for the weeks lab before your lab meeting. Pre-labs are due 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. Dress for lab: You must wear shoes that cover your entire foot, including the heel. They should fit up near your ankle; leather is preferred but any non-porous material is okay. Short shorts and short skirts are not allowed. Your clothing must cover your torso and legs down to your knees. Grading: The laboratory is graded on a Pass/No Pass basis. An average of 75% of all points available in the lab is required to pass. Late Work: Laboratory reports are due at the beginning of the lab period following completion of the experiment. Lab reports should be typed. Late reports will be docked 5 points per day late. Attendance: Attendance in this lab is mandatory. YOU MUST ATTEND ALL SCHEDULED LABORATORY MEETINGS. If you are not able to attend lab you must notify your laboratory instructor as soon as possible. For a missed lab meeting you must make up the missed lab time during the make-up. The make-up laboratory will take place during week 10 of the quarter, during the regularly scheduled lab period. In addition to completing the make up lab, students are responsible for completing the lab report for the missed lab. Data can be obtained from a lab partner or the lab instructor. The made up work should be clearly labeled and indicate the origin of the data reported. Reports are due the week following the syllabus deadline, unless other arrangements are made with the lab instructor. FAILURE TO DO BOTH WILL RESULT IN A NO PASS GRADE. If you miss two or more labs your grade will be a NO PASS. NOTE: If you are more than 15 minutes late to lab you will be marked late. Two late arrivals during the term will be counted as a missed lab. In addition, late students may be assigned to lab clean up duties at the conclusion of the lab period. If you are chronically late you will be given a NO PASS.
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Plagiarism: Experiments will be done in groups sharing the computer for data analysis and acquisition. You may compare data with other groups, but the content of your lab reports MUST be written individually. It will be considered an act of plagiarism if you borrow tables or graphs from another student (learning how to properly create a table or graph is an important skill, learn how to do it on your own!). You cannot paraphrase the internet, your book or any other source without the proper reference. Additionally, it will be considered an act of plagiarism if you borrow data without prior approval from your TA. There are additional 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, every lab report is worth 90 points, including the prelab, notebook and technique. Each lab report will be graded according to the following point distribution: Prelab: 10 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 20 points assigned on the following basis: Notebook: 10 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: 10 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
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Laboratory Safety Rules and Procedures Safety Rules The guidelines below are established for your and your classmates’ personal safety. Failure to adhere to the guidelines below will result in a loss of Lab Technique points. • Personal Protective Equipment (PPE) is
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Spills and other accidents can occur when least expected. For this reason it is necessary to wear proper PPE. The PPE for student labs consist of goggles, gloves and clothing. 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. 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. These are for sale in the bookstore and stockroom. You should not wear contact lenses in a chemical laboratory. Chemical vapors may become trapped behind the lenses and cause eye damage. Some chemicals may dissolve “soft” contact lenses. The most important aspect of having the goggles fit comfortably is the proper adjustment of the strap length. Adjust the strap length so that the goggles fit comfortably securely and are not too tight. If you find that your goggles tend to fog, you can pick-up anti-fog tissue from the stockroom. • Gloves – Gloves should be worn to protect the hands from chemicals. Gloves are provided through your student fees and are located in the student labs. For health and safety reasons it is important to always remove at least one glove when leaving the student laboratory, this prevents things such as door handles from getting contaminated. • Clothing – Dress appropriately for laboratory work. You must wear shoes that cover your entire foot, including the heel. They should fit up near your ankle; leather is preferred but any non-porous material is okay. Your clothing must cover your torso and legs down to your knees. Short shorts, short skirts, tank tops and halter tops are not allowed. • Eating, drinking and smoking are prohibited in the laboratory at ALL times. Wash your hands after finishing lab work and refrain from quick trips to the hall to drink or eat during lab. If you take a break, be certain to remove gloves and wash hands before ingesting food or drink. • Never work alone in the laboratory or in the absence of the instructor. • Headphones may not be worn in lab.
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Safety Procedures • Know location of safety equipment; fire extinguisher, fire blanket, first aid kit, safety shower, eyewash fountain and all exits. • In case of fire or accident, call the instructor at once. • Small fires may be extinguished by wet towels. • If a person’s clothing catches fire, roll the person in the fire blanket to extinguish the flames. • In case of a chemical spill on the body or clothing, stand under the safety shower and flood the affected area with water. Remove clothing to minimize contamination with the chemical. • If evacuation of the lab is necessary, leave through any door that is safe, or not obstructed; doors that lead to other labs may be the best choice. Leave the building by the nearest exit and meet your TA on the field next to Hoffmann Hall. This would also be the meeting place in the event of an earthquake or other emergency. It is good to know the nearest exits of your lab on the first day of class. • Spilled chemicals must be cleaned up immediately. If the material is corrosive or flammable, ask the instructor for assistance. If acids or bases are spilled on the floor or bench, neutralize with sodium bicarbonate, then dilute with water. Most other chemicals can be sponged off with water. • Avoid contact with blood or bodily fluids. 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. • If a mercury thermometer is broken, do not attempt to clean up yourself. Notify students around you, so that mercury is not spread, then notify your lab instructor or stockroom personnel. The stockroom is equipped for proper clean up and disposal of mercury.
Laboratory Procedures and Protocol General Etiquette: • Leave all equipment and work areas as you would wish to find them. • Keep your lab bench area neat and free of spilled chemicals. Your book bag, coat, etc., should be kept in the designated area at the entrance to the lab, not at your bench. • All chemical waste must be disposed of in proper containers. Proper disposal of chemicals is important for student safety and proper disposal. Putting chemicals into the wrong containers can lead to injury from unexpected chemical reactions. Mixing waste can also 6
make it more difficult or expensive for PSU to dispose of them. Only chemicals should go into waste jars. Waste jars for each experiment will be provided in the lab. They will be labeled specifying which contents should be placed inside. It is important that you replace the lids to the waste containers. When done with the waste jar, make sure it is placed in a secondary container. Do not put anything down the sink unless you are explicitly told to dispose of it this way. Your instructor will provide specific disposal guidelines when needed. Following these guidelines assists us in lowering the environmental impact of the labs. There are several locations for very specific waste. i. Chemical waste – these containers are ONLY for chemical waste generated in the lab. They are each specifically labeled for each lab and waste type. READ THE LABELS. ii. Contaminated paper waste – this is ONLY for paper towels used for clean-up of chemical spills. iii. Broken glass – this is ONLY for broken glassware. iv. Gloves – this is ONLY for used gloves. v. Normal trash – this is for all other trash that is not chemically contaminated, glass, or gloves. • Clean your bench and equipment Clean all your glassware- dirty glassware is harder to clean later. Wash with water and detergent scrubbing with a brush as necessary. Rinse well with water. Do not dry glassware with compressed air, as it is frequently oily. The water and gas should be turned off and your equipment drawer locked. • Clean the common areas before you leave the lab. Point deductions for the entire class will be imposed if the instructor or stockroom is not satisfied. • Return any special equipment to its proper location or the stockroom.
Handling Chemicals: Obtaining reagents: • Read the label CAREFULLY. The Chemicals are organized by experiment in secondary containment bins. Make sure the chemical name and concentration match what is required by the experiment! • Do not take the reagents to your bench. • We recommend always picking up bottles by the label. If all students do this, then any unnoticed spills when pouring will not cause possible problems for the next user. Remember to wear gloves while working with reagents. • Do not put stoppers or lids from reagents down on the lab bench. They may become contaminated. Be sure that the lids or stoppers are replaced. 7
• Do not place your own pipet, dropper, or spatulas into the reagent jar. Pour a small amount into a beaker and measure from that. Please pour on the conservative side to minimize waste and cost of labs. You can always go back for more. • Do not put any excess reagent back in the reagent jar. Treat it as waste and dispose of it properly. • When weighing chemicals on the balances, never weigh directly onto the weighing pan. Weigh into a weighing boat or beaker. Any spills on the balances MUST be cleaned up immediately. If you are unclear how to clean a spill, notify your instructor. The balances you are using are precision pieces of equipment and costs up to $4000. • All chemicals should be treated as potentially hazardous and toxic. Never taste a chemical or solution. When smelling a chemical, gently fan the vapors toward your nose. • Any chemicals that come in contact with your skin should be immediately washed with soap and copious amounts of water.
Laboratory Procedures • Never pipet any liquid directly by mouth! Use a rubber bulb to draw liquid into the pipet. • Never weigh hot chemicals or equipment. • When heating a test tube, always use a test tube holder and be certain never to point the open end of the test tube toward yourself or another person. • Handling glass tubing or thermometers: to insert glass tubing into a rubber stopper, lubricate the glass tubing with a drop of glycerin, hold the tubing in your hand close to the hole, and keep all glass pieces wrapped in a towel while applying gentle pressure with a twisting motion. • To prepare a dilute acid solution from concentrated acid, acid should be added slowly to water with continuous stirring. This process is strongly exothermic, and adding water to acid may result in a dangerous, explosive spattering. • Use the fume hood for all procedures that involve poisonous or objectionable gases or vapors. • Never use an open flame and flammable liquids at the same time.
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Keeping a Lab Notebook In keeping a lab notebook, there are certain principles that should be followed. These boil down to being clear and complete in your entries in your lab notebook. There are also certain conventions for lab notebooks that are universally 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. General Guidelines • Write your name on outside front of notebook • Use black ink, fine-tipped ball-point pen (this will photocopy clearly) • At the front of the notebook, 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, manual calculation, and working notes • Prepare data tables in advance - with columns for calculated results and notes • Working graphs done in lab notebook to monitor progress Usage and Structure You should record all your work in your lab notebook. That is the proper place for all lab planning and observations. Nothing should be recorded on odd scraps of paper, etc. The overriding principle for a lab notebook is to record in it all the pertinent information about your lab work. This boils down to clear descriptions of what you did and what you observed as a result. It is a working tool, and a reference for other researchers who might want to read your notebook and reproduce your work. (This applies to notebooks in learning laboratories: Your lab instructor may want to look at what you did in order to understand your results. This is often the case. So, it needs to be clear.) The word “clear” here is crucial. In order to be clear, data must be recorded in well-thought-out tables, clearly labeled. Descriptions of procedures must be clear and concise; to the point.
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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. Title: With your lab notebook laid open, on the right hand page write down the title of the experiment, and the date. In general, you will use the right-hand page for all your writing. The left-hand page is reserved for recording scratch work. Don’t use this space until you need to. One example of how to use the left-hand page: if your work requires simple calculations using your measurements, use the left-hand page to do the calculations. If unexpected results occur later, sometimes you can look back at your scratch work and discover the error. (“Oh, I subtracted wrong! We put in 10.5 grams of copper sulfate, not 9.5 like we thought!”) Better to discover the error after the fact than never to discover it at all. Purpose: Below the title, write the purpose of the experiment in one or two sentences. This section serves to remind you and notify the reader what the experiment is about. Procedure and Observations: This next section will be labeled Procedure and Observations. As the name suggests, write down what you actually do and what you observe. This section is where you should have preprepared tables for data collection. Set up this section by dividing the page into a right and left column. In the left hand column write your procedure and in the right column next to the procedure, record observations and data or measurements. Results and Discussion: You might want to include a final section that is labeled Results and Discussion. In this section, you would describe what results you got, what conclusions you have reached, ideas for continuing work, etc.
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An example of a prepared notebook follows.
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Writing Style in the Lab Notebook For certain entries in your lab notebook, such as the Introduction before each experiment, you should strive to write as logically and clearly as possible. It is also a good idea to write in the third person passive voice, to get into the habit, and so that in many cases you can copy entries from your lab notebook into your reports without the need for major revisions/rewrite. However, this is a working document. It is not expected that you write perfect prose in your notebook – it is a first draft. Just do the best you can. Also, as a working document, 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. Example Notebook entry: “Added 10 mL of 1M HCl – solution turned red instantly; pcpt.↓ a few secs later→ clr soln.” When written into a lab report or journal article, this would be expanded a bit and made grammatically correct. “10 mL of 1.0 M HCl were added to the clear reaction mixture. This immediately resulted in a crimson solution, and a red precipitate formed a few seconds later, leaving a clear solution.”
Adapted courtesy of Keith James
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Report Guidelines For each experiment performed this term you will turn in a type written report (at the end of each lab you will find a summary of which sections to include in the report for that lab). The reports are due at the beginning of class the week following completion of the experiment. Below is a description of what should be included in each section. The sections are presented here in the order they should appear in your lab report. It is expected that you will complete each experiment and do the necessary calculations and analysis during the scheduled lab period each week. You may discuss the calculations and analysis with your lab mates. Your written lab report should be your own individual work!! The lab report sections should be complete but CONCISE. For most experiments this term, your report should be 1-2 pages long.
Writing Style You will write you reports using a formal scientific writing style. A lab report must be written in the third person, passive voice. Also, it must be in the past tense. It should not contain personal pronouns such as, “I”, “we” or “he” neither should it contain proper names of persons. Good: “50 mL of 1.0 M HCl were poured into a 125 mL Erlenmeyer flask” Bad: “I poured 50 mL of hydrochloric acid into a flask.” Also bad: “Joe Shmoe poured 50 mL of hydrochloric acid into a flask.” This is not the correct form of 3rd person. It includes Joe’s name. Also bad: “We are going to put 50 mL of acid into the flask.” Uses future tense; also, “we”. After you write your report, there is one more thing to do before you print it and hand it in: Proofread it! Read it out loud. If is doesn’t sound right, it isn’t. Fix it. Then do it again until it is right. You will enjoy writing reports more if you take pride in what you hand in.
Abstract: This is like a condensed version of your lab report. It is a stand-alone document. Abstracts are, in fact, often published separately from the articles they describe. A library search of the literature generally involves reading abstracts. This is done with the aim to identify articles that need to be read in full, and eliminate many others whose abstract makes it clear that they are not relevant to the study at hand. So, the abstract needs to be brief, but complete. There are three questions that should be answered in any good abstract 1. What did you do? 2. How did you do it? 3. What did you find? 14
Even though it sequentially appears first, you should consider writing this part of the lab report after you have finished the remaining sections.
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Introduction: Here, you want to address WHY you did this experiment. Your introduction begins with a statement of the purpose of the experiment. You will do this again in the abstract, but remember the Abstract is a stand-alone document. What you say there will not count; you will find that as you write the report that you will be repeating yourself a bit. Next, provide any relevant background, to put the experiment into context. Include any key concepts, mathematical equations or chemical equations needed by the reader to understand your experiment. This means that your Introduction will often include some explanation of the theory behind the experiment. Don’t just write the equations, but provide information as to why they are relevant. You may consider writing your introduction with a central theme, such as density, types of chemical reactions…..
Data: This is section is where your experimental data belong. In this section you would also include observations and descriptions of other pertinent events. This section is not where the calculations, interpretation and discussion of your results belong. (In published papers, a data section is usually not included, but, this is a class so this section will be included.) Tables Whenever possible, data should be presented in the clearest format possible, usually in the form of a table. When you present your data in a table it is necessary to take the following into account. • Number tables sequentially as they appear (Table 1, Table 2….). • Be sure to refer the reader to view the tables in the text. • Construct a descriptive table caption and place it above the table. • Tables should include descriptive column headings, including units. • Tables should not be divided across page boundaries For a simple example, see Table 1. Table 3: Mass of Water as Determined by a Pan Balance (+/- 0.01g): Here the volume of water delivered by a 10 mL volumetric pipet was determined utilizing the mass of water delivered and waters density (0.9980 g/ml).
Run # mass water weighed (g) 1 9.95 2 9.94 3 9.98 Average 9.96 Error +/- 0.02
Volume water (ml) 9.98 +/- 0.02
Graphs When graphical presentation of data is necessary, please prepare graphs using the following guidelines. • Number figures sequentially as they appear (Figure 1, Figure 2….). • In your writing, be sure to cite the tables in the text. • Insert a caption below he graph that indicates what is being plotted on the y-axis vs. what is being plotted on the x-axis (always y vs. x) 16
• • • •
Each axis should be clearly labeled, including units. Figures should not be divided across page boundaries Remove gridlines, titles and equations from the graph. If this information is pertinent, it should be included in the caption. If the slope or intercept is necessary for other parts of the experiment, then place the values in the caption with proper units.
For a simple example, see Figure 1.
Figure 1. A calibration curve for the absorbance at 470 nm of aqueous Allura red solutions as a function of the concentration. A best fit line was rendered resulting in a slope of 5.86 mM-1.
Results: The results section is where you should show sample calculations and report all of your results. For every type of calculation you should show one sample calculation. Each calculation should have a descriptive title, i.e. “Calculating the density of Coca-Cola”. The calculation section should be annotated. The annotation is provided to describe why each calculation is useful and relevant to the lab activity. The description should not be any longer than two or three sentences and should help you describe your results in your discussion section. Sample calculations may be written by hand attached as an appendix to your report. The results of all calculations should be summarized in a table where appropriate. Calculating the density of Coca-Cola The volume (355 mL) and mass (394 g) of the contents of a can of coke had previously been determined above. The density is determined utilizing the relationship d=m/v (equation 1) which was explained in the introduction. d = 394 g / 355 mL = 1.11 g/mL
Discussion: In this section, you will discuss interpretations of the experimental results. This is where you get to present your thinking process. For any labs that have questions to answer, this is also where the answers get written up. 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. This DOES NOT mean to include detailed procedures or that you need to re-explain your 17
calculations in words. It DOES mean that a general description of the experiment can be useful in explaining your results and putting them in context. In this section you should also discuss error analysis. This does not necessarily mean trying to explain what went wrong. (Maybe nothing did go wrong!) It means discussing the limitations of your experiment. For example, if you are doing calorimetry in a coffee cup, and the cup feels warm to your hand, it means that some heat is escaping. Also, if you are reading a 5 degree temperature change with a thermometer that you can only read to the nearest 0.5 degree, there is a significant uncertainty in the exact magnitude of the temperature change. You could easily have a 10% error, or even more, and this needs to be taken into account. It at least needs to be mentioned, to show that you were aware of the issue. This is a limitation of the apparatus, not an error on your part. And, yes, if something did go wrong (your lab partner forgot to write down the exact molarity of your reagent), then that should go here, too, along with an explanation of how you attempted to correct for the error. (In this case, you may have had to re-do the experiment.)
Adapted courtesy of Keith James.
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Example Lab Report Following is an example of a lab report prepared according to the previous report guidelines. Sample calculations can be written on a separate paper and attached to the report.
Calibration of a 10 ml Volumetric Pipette Abstract: A 10 ml volumetric pipette was calibrated by determining the mass of water delivered by a pipette. A pipet was used to precisely deliver 10 mL of water. The mass of water was then converted to volume using the density of water. The volume of the pipette was determined to be 9.98 +/- 0.02 ml when the mass of water was determined on a pan balance and 9.998 +/- 0.002 ml when determined with an analytical balance.
Introduction: A volumetric pipette is designed to deliver a stated volume of liquid; however, the actual amount of liquid any individual pipette delivers may vary slightly from this ideal stated volume. In order to determine the actual volume an individual pipette delivers, it will be calibrated. In this case, calibration refers to the comparison of the actual amount of liquid delivered by the pipette to the standard value of the pipette (10 ml). Because delivered volume is being calculated, another measurable quantity must be used to verify the volume delivered by the pipette. In this case, the relationship between mass and volume (density) will be used. Mass is an easily measurable quantity that can be determined with a high degree of accuracy due to the availability of electronic balances. Mass can then be converted to volume by the use of density. D (density) = m (mass)/ V (volume) The density of water at a variety of temperatures is readily available and will be used here to calibrate the volume of the pipette.
Data: Diameter of beaker: 3.9 cm +/- 0.1 cm Mass of water evaporated in 60 seconds: 0.0016g +/- 0.0002g Temp of water: 20.5 ºC +/- 0.2 ºC Density of water: 0.9980 g/ml Table 1: Mass Determined by Pan Balance (+/- 0.01g)
Run # 1 2 3
mass beaker (g) 27.88 27.88 27.88
mass beaker + water (g) 37.83 37.82 37.86
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Table 2: Mass Determined by Analytical Balance (+/0.0001g)
mass beaker Run # (g) 1 27.2349 2 27.2348 3 27.2335
mass beaker + water (g) 36.5618 36.7813 36.8251
t(transfer) 2:29:00 2:32:00 2:41:20
t(weigh) 2:30:30 2:33:20 2:42:30
Results: Calculation of the volume of water: In this calculation, the average mass of water for the three trials, shown in Table 1, as determined by the pan balance was divided by the know density of water at 20.5 ºC. The data are summarized in Table 3. Volume = 9.96 g / 0.9980 g/mL = 9.98 mL Table 3: Mass of Water as Determined by a Pan Balance (+/0.01g): Here the volume of water delivered by a 10 mL volumetric pipet was determined utilizing the mass of water delivered and waters density (0.9980 g/ml).
Run # mass water weighed (g) 1 9.95 2 9.94 3 9.98 Average 9.96 Error +/- 0.02
Volume water (ml) 9.98 +/- 0.02
Calculation for the mass evaporated: To correct for evaporation of water in the time it takes to measure the mass of the water delivered by the volumetric pipet, the mass of water that evaporated was estimated. The rate of evaporation of water in the 50 mL beaker in 60 seconds: 0.0016g. The data are summarized in Table 4. Mass evaporated = rate of evaporation x time of evaporation = (0.0016 g/ 60 s) x 90 s = 0.0024 g Calculation of the mass transferred: The mass of water initially transferred was the sum of the mass of water evaporated and the mass of water present at the time of weighing (Table 2). The data are summarized in Table 4. Mass transferred = mass water weighed + mass (transferred) = 9.9769 + 0.0024 = 9.9793 g
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Table 4: Mass of Water as Determined by an Analytical Balance (+/-0.0001 g): ): Here the volume of water delivered by a 10 mL volumetric pipet was determined utilizing the mass of water delivered and the density water (0.9980 g/ml). A correction was added to account for the water that evaporated during the measurement.
Run # mass water weighed (g) t(evap) (s) mass (evap) (g) mass (transferred) (g) Volume (ml) 1 9.9769 90 0.0024 9.9793 2 9.9735 80 0.0021 9.9757 3 9.9756 70 0.0019 9.9775 Average 9.978 9.998 Error +/- 0.002 +/- 0.002
Discussion: The mass of water delivered by a 10 ml volumetric pipette was determined on both a pan balance and an analytical balance (Tables 1 and 2 respectively). The mass of water was then converted to volume using the density of water. In the case of the analytical balance, the rate of evaporation of water (which is a systematic error) was taken into consideration. In this case the mass of water that evaporated from the time the water was delivered to the beaker to the time of weighing was added to the weighed mass of water delivered by the pipette. This correction was not necessary when the pan balance was used since the accuracy of the pan balance is +/- 0.01 g and the evaporation rate of water under experimental conditions was found to be 2.7 x 10-5 g/s. The use of the analytical balance increased both the precision and the accuracy of the calculated volume of the pipette (9.98 +/- 0.02 ml with the pan balance and 9.998 +/- 0.002 ml with the analytical balance, see tables 3 and 4). The improvement in the results can easily be seen by the percent error which was calculated to be 0.2 % with the pan balance and 0.02 % with the analytical balance. The largest source of error in this experiment most likely came from the difficulty in accurately filling the pipette to the mark with water which introduced random error into the experiment.
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General Chemistry Lab Report Checklist General _____ Have you listed your name, partner's name, 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, we, they, etc.) _____ Is proper tense is maintained within sections? _____ Have you correctly written your chemical formula and names correctly? _____ Were correct subscripts, superscripts, and symbols are used? _____ Did you separate the numbers from their units (0.25 mL was added…. not 0.25mL was added)? _____ Did you check significant figures? _____ Do your numbers include leading zeros (0.25 mL was added…. not .25 mL was added)? _____ Did you make sure that you did not start a sentence with a number? _____ Are your references cited in one official style? _____ Have you made a citations whenever ideas from outside? _____ 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. Abstracts are often written last. They should be clear, concise, and self-contained and, in the context of this lab, approximately three sentences long. _____What did you do? (Identify the rationale behind the investigation)? _____How did you do it (summarize the procedure, 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. _____Did you include a statement of purpose? _____Is there sufficient background so that the reader can understand what you did? _____Are necessary equations, chemical or mathematical, included?
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Data This section should give only the data and observations from the lab, without results _____Are your data tables properly formatted? _____Are your calculations, either attached as an appendix, or typed neatly into the data section? _____Are your figures and tables numbered sequentially and referred to in the text. Table captions above and figure captions below. 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?
Results:
We will be treating this section as a calculational section. This is where you will be showing all calculations along with a written description as to how the calculations were carried out and what the result of the calculation is and how it relates to the lab. Your readers must easily find your results in order to evaluate and interpret them. _____Are calculations accompanied by text explaining the both the method of calculation and results of the calculation? _____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 discuss interpretations of the experimental results. It will be necessary to describe your results, cite tables or figures. It should include a general description of the experiment to put the results into context. _____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 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?
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Acid Rain Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. Carbon dioxide (CO2) reacts with water to produce carbonic acid (H2CO3). Write the reaction showing what happens when carbonic acid is dissolved in water. 2. What is the conjugate base of nitrous acid (HNO2)? 3. Which is a stronger acid, nitrous acid (HNO2) or nitric acid (HNO3)? 4. Which is a stronger base, nitrite (NO2-) or nitrate (NO3-)? 5. Describe the method you will use in this lab to generate the acids found in acid rain.
Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted 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.
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Lab: Acid Rain In this experiment, you will observe the formation of four acids that occur in acid rain: carbonic acid, H2CO3 • nitrous acid, HNO2 • nitric acid, HNO3 • sulfurous acid, 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, nitrogen dioxide (NO2). Most nitrogen dioxide in our atmosphere is produced from automobile exhaust. 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, sulfur dioxide (SO2). Most sulfur dioxide gas in the atmosphere results from burning coal containing sulfur impurities. Sulfur dioxide dissolves in rain drops and forms sulfurous acid:
NO2
H 2 CO3 H NO2 H NO3 H2SO3 SO2
(3) SO2(g) + H2O(l) → H2SO3(aq) In the procedure outlined below, you will first produce these three gases. You will then bubble the gases through water, producing the acids found in acid rain. The acidity of the water will be monitored with a pH Sensor.
OBJECTIVES In this experiment, you will Generate three gaseous oxides, CO2, SO2, 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. • •
MATERIALS Vernier pH Sensor 3 Beral pipets with a 15 cm stem
1 HCl pipet 3 Beral pipets with a 2 cm stem
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PROCEDURE 1. Obtain and wear goggles. 2. Obtain three short-stem and three long-stem Beral pipets. Label the short-stem pipets with the formula of the solid they will contain: “NaHCO3”, “NaNO2”, and “NaHSO3”. Label the long-stem pipets with the formula of the gas they will contain: “CO2”, “NO2” and “SO2”. You can use a 100 mL beaker to support the pipets. 3. Obtain a beaker containing solid NaHCO3. Squeeze the bulb of the pipet labeled “NaHCO3” to expel the air, and place the open end of the pipet into the solid NaHCO3. When you release the bulb, solid NaHCO3 will be drawn up into the pipet. Continue to draw solid into the pipet until there is enough to fill the curved end of the bulb, as shown in Figure 1.
Figure 1
4. Repeat the Step 3 procedure to add solid NaNO2 and NaHSO3 to the other two Beral pipets. CAUTION: Avoid inhaling dust from these solids. 5. Obtain a Beral pipet and label it HCl. Squeeze the bulb to expel some of the air, and place the open end of the pipet into a beaker containing 1.0 M HCl. When you release the bulb, HCl will be drawn up into the pipet. CAUTION: HCl is a strong acid. Gently hold the pipet with the stem pointing up, so that HCl drops do not escape. Insert the narrow stem of the HCl pipet into the larger opening of the pipet containing the solid NaHCO3, as shown in Figure 2. Gently squeeze the HCl pipet to add about 20 drops of HCl solution to the solid NaHCO3. When finished, remove the HCl pipet. Gently swirl the pipet that contains NaHCO3 and HCl. Carbon dioxide, CO2, is generated in this pipet. Place it in the 100 mL beaker, with the stem up, to prevent spillage. 6. Repeat the procedure in Step 5 by adding HCl to the pipet containing solid NaHSO3. Sulfur dioxide, SO2, is generated in this pipet.
Figure 2
7. Repeat the procedure in Step 5 by adding HCl to the pipet containing solid NaNO2. Nitrogen dioxide, NO2, is generated in this pipe. Leave the three gas-generating pipets in the 100 mL beaker until Step 10. 8. Use a utility clamp to attach a 20 × 200 mm test tube to the ring stand. Add about 4 mL of distilled water to the test tube. Remove the pH Sensor from the pH storage solution, rinse it off with distilled water, and place it into the tap water in the test tube. 9. Connect the pH Sensor to the computer interface. Prepare the computer for data collection by opening the file “Exp 23 acid rain” from the Chemistry with Computer folder. Check to see that the pH is between 6 and 8 for the water. 10. Squeeze all of the air from the bulb of the long-stem pipet labeled “CO2”. Keep the bulb completely collapsed and insert the long stem of the pipet down into the gasgenerating pipet labeled “NaHCO3”, as shown in Figure 3. Be sure the tip of the long-stem pipet remains above the liquid in the gas-generating pipet. Release the pressure on the bulb so that it draws gas up into it. Store the long-stem pipet and the gas-generating pipet in the 100 mL beaker. 11. Repeat the procedure in Step 10 using the pipets labeled “NaNO2” and “NO2”. 12. Repeat the procedure in Step 10 using the pipets labeled “NaHSO3” and “SO2”.
Figure 3
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13. Insert the long-stem pipet labeled “CO2” into the test tube, alongside the pH Sensor, so that its tip extends into the water to the bottom of the test tube (see Figure 4). 14. To begin collecting data, click . After 15 seconds have elapsed, gently squeeze the bulb of the pipet so that bubbles of CO2 slowly bubble up through the solution. Use both hands to squeeze all of the gas from the bulb. When data collection stops after 120 seconds, 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 pH stabilized). To confirm these two values, click the Statistics button, , and examine the minimum and maximum values in the pH box displayed on the graph. Record the initial and final pH values in your data table. Close the Statistics box by clicking in the upper left corner of the box. 15. Rinse the tip of the pH Sensor thoroughly with distilled water and return it to the sensor storage solution. Discard the contents of the test tube as directed by your teacher. Rinse the test tube thoroughly with tap water. Add 4 mL of tap water to the test tube. Place the pH Sensor in the test tube and check to see that the input display shows a pH value that is about the same as the previous initial pH. If not, rinse the test tube again. 16. From the Experiment menu, choose Store Latest Run. This stores the data so it can be used later, but it will be still be displayed while you do your second and third trials. 17. Repeat Steps 13-16 using NO2 gas. 18. Repeat Steps 13-14 using SO2 gas. When you are finished, rinse the pH Sensor with distilled water and return it to the sensor storage solution. Discard the six pipets as directed by your Instructor. Figure 4 19. Label all three curves by choosing Text Annotation from the Insert menu, and typing “carbon dioxide” (or “nitrogen dioxide”, or “sulfur dioxide”) in the edit box. 20. Print copies of the graph, with all three data sets displayed.
PROCESSING THE DATA For each of the three gases, calculate the change in pH (ΔpH), by subtracting the initial pH from the final pH. Record these values in your data table.
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Acid Rain Lab Report: Your report for this lab should include the following sections: Abstract Introduction: Data: Include a table with the data from all three gases Results: Calculate the change in pH for each of the gases and summarize in a table Discussion: Discuss the experiment and any possible sources of error In addition, answer the following questions as part of your discussion: 1. In this experiment, which gas caused the smallest drop in pH? 2. Which gas (or gases) caused the largest drop in pH? 3. 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. How would burning low-sulfur coal lower the level of acidity in rainfall? Use specific information about gases and acids to answer the question. 4. High temperatures in the automobile engine cause nitrogen and oxygen gases from the air to combine to form nitrogen oxides. What two acids in acid rain result from the nitrogen oxides in automobile exhaust? 5. Which gas and resulting acid in this experiment would cause rainfall in unpolluted air to have a pH value less than 7 (sometimes as low as 5.6)? All three gases are produced by man but one occurs naturally at relatively high concentrations. 6. Which acid is the weakest? Which is the strongest?
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Acid Dissociation Constant, Ka Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. What would be the measured pH of a 0.10 M solution of NaOH? 2. What would be the measured pH of a 0.10 M solution of HCl? 3. Write the equilibrium constant expression, Ka, for the dissociation of acetic acid, HC2H3O2. 4. What would be the measured pH of a 1.0 M solution of HC2H3O2? 5. Determine the volume, in mL, of 2.00 M HC2H3O2 required to prepare 100 mL of a 0.30 M HC2H3O2 solution?
Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted 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.
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Acid Dissociation Constant, Ka Acetic Acid, HC2H3O2, is a weak acid that dissociates according to this equation: –
HC2H3O2(aq) → H+(aq) + C2H3O2 (aq) In this experiment, you will experimentally determine the dissociation constant, Ka, for acetic acid, starting with solutions of different initial concentrations.
OBJECTIVES In this experiment, you will • • •
Gain experience mixing solutions of specified concentration. Experimentally determine the dissociation constant, Ka, of an acid. Investigate the effect of initial solution concentration on the equilibrium constant.
MATERIALS Vernier pH Sensor pipet bulb
pipets 2 x 100 mL volumetric flasks
Figure 1
PROCEDURE 1. Obtain and wear safety goggles. 2. Put approximately 25 mL of distilled water into a 100 mL volumetric flask. 3. Your instructor will assign each group two different concentrations of HC2H3O2. Calculate the volume of a 2.0 M HC2H3O2 stock solution necessary to make 50 mL of each solution. Use a pipet bulb (or pipet pump) to pipet the required volume of 2.00 M acetic acid into the volumetric flask. CAUTION: Use care when handling the acetic acid. It can cause painful burns if it comes in contact with your skin or gets into your eyes. Fill the flask with distilled water to the 50 mL mark. To prevent overshooting the mark, use a wash bottle filled with distilled water for the last few mL. Mix thoroughly. 4. Use a utility clamp to secure a pH Sensor to a ring stand as shown in Figure 1.
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5. Connect the probe to the computer interface. Prepare the computer for data collection by opening the file “Exp 27 Acid Dissociation Ka” from the Chemistry w/ Computer folder of Logger Pro. 6. 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. • Pour about 20 mL of your solution into a clean 100 mL beaker and use it to thoroughly rinse the sensor. • Use the remaining 30 mL portion to determine pH. Swirl the solution vigorously. (Note: Readings may drift without proper swirling!) Record the measured pH reading in your data table. • When done, place the pH Sensor in distilled water. •
7. Repeat the procedure for your second assigned solution.
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.
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Acid Dissociation Lab Report: Your lab report should include the following sections: Abstract Introduction: Data: Include the concentrations of the two HC2H3O2 solutions you made and the volume of the stock solution and water used to make the solutions. Include a copy of the ICE table used for each concentration of HC2H3O2. Results: Report your calculated Ka value for each concentration used and an average Ka. Compare your experimentally determined Ka with the accepted value at 25 °C (1.8 x 10-5) and calculate the percent error. Calculate the percent dissociation of acetic acid for each concentration used.
Discussion: Discuss the experiment and any possible sources of error In addition, answer the following questions as part of your discussion: 1. What effect does initial HC2H3O2 concentration seem to have on Ka? 2. What effect does initial HC2H3O2 concentration seem to have on the percent dissociation?
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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. What is a diprotic acid? Give an example not found in the text to the following experiment. 2. Give the balanced chemical reaction for the titration of a generic diprotic acid, H2X, with potassium hydroxide. 3. When titrating 50.0 mL of 0.10 M H2SO4 with 0.10 M NaOH, how many mL of NaOH will you have added to reach the 1st equivalence point? 4. A student completes a titration of an unknown diprotic acid. In this experiment, 0.79 g of the acid is dissolved in 250.0 mL of water. It requires 13.48 mL of 1.0 M NaOH to reach the second equivalence point. What is the molar mass of the acid?
Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted 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.
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Lab: Titration of a Diprotic Acid: Identifying an Unknown A diprotic acid is an acid that yields two H+ ions per acid molecule. Examples of diprotic acids are sulfuric acid, H2SO4, and carbonic acid, H2CO3. A diprotic acid dissociates in water in two stages: (1) H2X(aq)
H+(aq) + HX–(aq)
(2) HX–(aq)
H+(aq) + X2 (aq)
–
Because of the successive dissociations, titration curves of diprotic acids have two equivalence points, as shown in Figure 1. The equations for the acid-base reactions occurring between a diprotic acid, H2X, and sodium hydroxide base, NaOH, are from the beginning to the first equivalence point: (3) H2X + NaOH
NaHX + H2O
from the first to the second equivalence point: (4) NaHX + NaOH
pH
Na2X + H2O
from the beginning of the reaction through the second equivalence point (net reaction): (5) H2X + 2 NaOH
Na2X + 2 H2O
1st Equivalence Point 2nd Equivalence Point Volume NaOH
At the first equivalence point, all H+ ions from the Figure 1 first dissociation have reacted with NaOH base. At + the second equivalence point, all H ions from both reactions have reacted (twice as many as at the first equivalence point). Therefore, the volume of NaOH added at the second equivalence point is exactly twice that of the first equivalence point (see Equations 3 and 5).
The primary purpose of this experiment is to identify an unknown diprotic acid by finding its molecular weight. A diprotic acid is titrated with NaOH solution of known concentration. Molecular weight (or molar mass) is found in g/mole of the diprotic acid. Weighing the original sample of acid will tell you its mass in grams. Moles can be determined from the volume of NaOH titrant needed to reach the first equivalence point. The volume and the concentration of NaOH titrant are used to calculate moles of NaOH. Moles of unknown acid equal moles of NaOH at the first equivalence point (see Equation 3). Once grams and moles of the diprotic acid are known, molecular weight can be calculated, in g/mole. Molecular weight determination is a common way of identifying an unknown substance in chemistry. You may use either the first or second equivalence point to calculate molecular weight. The first is somewhat easier, because moles of NaOH are equal to moles of H2X (see Equation 3). If the second equivalence point is more clearly defined on the titration curve, however, simply divide its NaOH volume by 2 to confirm the first equivalence point; or from Equation 5, use the ratio: 1 mole H2X / 2 mol NaOH
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OBJECTIVE In this experiment, you will identify an unknown diprotic acid by finding its molecular weight.
MATERIALS Vernier pH Sensor
50 mL buret
PROCEDURE 1. Obtain and wear goggles. 2. Weigh out about 0.120 g of the unknown diprotic acid on a piece of weighing paper. Record the mass to the nearest 0.001 g in your data table. Transfer the unknown acid to a 250 mL beaker and dissolve in 100 mL of distilled water. CAUTION: Handle the solid acid and its solution with care. Acids can harm your eyes, skin, and respiratory tract. 4. Add 6 drops of the indicator mixture. The indicator mixture contains a mixture of bromocresol green and phenol red. 4. Use a utility clamp to suspend a pH Sensor on a ring stand as shown here. Position the pH Sensor in the diprotic acid solution and adjust its position toward the outside of the beaker so it will not be easier to stir the solution with a stirring rod without striking the sensor. 5. Obtain approximately 60 mL of ~0.1 M NaOH solution in a 250 mL beaker. Obtain a 50 mL buret and rinse the buret with a few mL of the ~0.1 M NaOH solution. Record the precise concentration of the NaOH solution in your data table. Use a utility clamp to attach the buret to the ring stand. Fill the buret a little above the 0.00 mL level of the buret. Drain a small amount of NaOH solution into a waste beaker so it fills the buret tip and leaves the NaOH close to the 0.00 mL level of the buret. Be sure to record the actual buret reading. ALL BURET READINGS NEED TO BE RECORDED TO TWO DECIMAL PLACES. Dispose of the waste solution from this step as directed by your teacher. CAUTION: Sodium hydroxide solution is caustic. Avoid spilling it on your skin or clothing. 6. Connect the pH Sensor to the computer interface. Prepare the computer for data collection by opening the file “Exp 24a Acid Base Titration” from the Chemistry w/ Computer folder of Logger Pro. 7. 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. Be sure to thoroughly mix the solution with a stirring rod when adding base. Poorly mixed solutions can lead to inaccurate pH readings. a. Before adding NaOH titrant, click and monitor pH for 5-10 seconds. Once the pH has stabilized, click . In the edit box, type in the initial buret reading, and press ENTER to store the first data pair for this experiment. b. Add enough NaOH to raise the pH by about 0.20 units. When the pH stabilizes, again click . In the edit box, type the current buret reading, to the nearest 0.01 mL. Press ENTER. You have now saved the second data pair for the experiment. 35
c. Continue adding NaOH solution in increments that raise the pH about 0.20 units and enter the buret reading after each addition. Proceed in this manner until the pH is 3.5. d. When pH 3.5 is reached, change to 2-drop increments. Enter the buret reading after each increment. Additionally, note any change in the color of the solution e. After pH 4.5 is reached, again add larger increments that raise the pH by about 0.20 units and enter the buret reading after each addition. Continue in this manner until a pH of 7.5 is reached. f. When pH 7.5 is reached, change to 2-drop increments. Enter the buret reading after each increment. g. When pH 10 is reached, again add larger increments that raise the pH by 0.20 units. Enter the buret reading after each increment. Continue in this manner until you reach a pH of 11. 8. When you have finished collecting data, click directed by your teacher.
. Dispose of the beaker contents as
9. Print a copy of the table. Then print a copy of the graph.
PROCESSING THE DATA 1. On your printed graph, one of the two equivalence points is usually more clearly defined than the other; 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. Indicate the more clearly defined equivalence point (first or second) in your data table. 2. 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, examine the data to find the largest increase in pH values during the 2-drop additions of NaOH. Find the NaOH volume just before this jump. Then find the NaOH volume after the largest pH jump. Identify both of these data pairs and record them. 3. Determine the volume of NaOH added at the equivalence point you selected in Step 1. To do this, add the two NaOH volumes determined in Step 2, and divide by two. For example: 12.34 + 12.44 = 12.39 mL 2 4. Calculate the number of moles of NaOH used at the equivalence point you selected in Step 1. 5. Determine the number of moles of the diprotic acid, H2X. Use Equation 3 or Equation 5 to obtain the ratio of moles of H2X to moles of NaOH, depending on which equivalence point you selected in Step 1. 6. Using the mass of diprotic acid you measured out in Step 1 of the procedure, calculate the molecular weight of the diprotic acid, in g/mol. 7. From the following list of five diprotic acids, identify your unknown diprotic acid. 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 36
8. Determine the percent error for your molecular weight value in Step 6. 9. 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. Examine the data to find the largest increase in pH values during the 2-drop additions of NaOH. Find the NaOH volume just before and after this jump. Underline both of these data pairs on the printed data table and record them in the Data and Calculations table. 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. 10. Determine the volume of NaOH added at the alternate equivalence point, using the same method you used in Step 3. 11. On your printed graph, clearly specify the position of the equivalence point volumes you determined in Steps 3 and 10, using dotted reference lines like those in Figure 1. Specify the NaOH volume of each equivalence point on the horizontal axis of the graph.
EQUIVALENCE POINT DETERMINATION: An Alternate 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. The equivalence point volume corresponds to the peak (maximum) value of the first derivative plot, and to the volume where the second derivative equals zero on the second derivative plot. 1. View the first-derivative graph (ΔpH/Δvol) by clicking the on the vertical-axis label (pH), and choose First Derivative. You may need to autoscale the new graph by clicking the Autoscale button, . 2. View the second-derivative graph (Δ2pH/Δvol2) by clicking on the vertical-axis label, and choosing Second Derivative. In Method 2, view the second-derivative on Page 3 by clicking on the Next Page button, .
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Extension Using a half-titration method, it is possible to determine the acid dissociation constants, Ka1 and Ka2, for the two dissociations of the diprotic acid in this experiment. The Ka expressions for the first and second dissociations, from Equations 1 and 2, are: [H+][HX-] [H+][X2-] Ka2 = [H2X] [HX-] The first half-titration point occurs when one-half of the H+ ions in the first dissociation have – been titrated with NaOH, so that [H2X] = [HX ]. Similarly, the second half-titration point occurs when– one-half of the H+ ions in the second–dissociation have been titrated with NaOH,– so that – 2– [HX ] = [X ]. Substituting [H2X] for [HX ] in the Ka1 expression and, [HX ] for [X2 ] in the Ka2 expression, the following are obtained: Ka1 =
Ka1 = [H+]
Ka2 = [H+]
Taking the base-ten log of both sides of each equation, logKa1 = log[H+] Thus, the pH value at the first half-titration volume, Point 1 in Figure 2, is equal to the pKa1 value. The first half-titration point volume can be found by dividing the first equivalence point volume by two. Similarly, the pH value at the second titration point, is equal to the pKa2 value. 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). Use the method described below to determine the Ka1 and Ka2 values for the diprotic acid you identified in this experiment.
logKa2 = log[H+]
pK a2 pH
pK a1
1
1st EP
2
2nd EP
Volume NaOH
Figure 2 1. 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). Then determine the precise NaOH volume of the second half-titration point halfway between the first and second equivalence points. 2. On your graph of the titration curve, draw reference lines similar to those shown in Figure 2. Start with the first half-titration point volume (Point 1) and the second half-titration point volume (Point 2). Determine the pH values on the vertical axis that correspond to each of these volumes. Estimate these two pH values to the nearest 0.1 pH unit. These values are the pKa1 and pKa2 values, respectively. (Note: See if there are volume values in your data table similar to either of the half-titration volumes in Step 1. If so, use their pH values to confirm your estimates of pKa1 and pKa2 from the graph.) 3. From the pKa1 and pKa2 values you obtained in the previous step, calculate the Ka1 and Ka2 values for the two dissociations of the diprotic acid.
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Titration of a Diprotic Acid Report: You will write a complete lab report for this lab which will contain the following sections: Abstract Introduction: Introduce the major concepts in this lab: titration, titration curves, equivalence point, diprotic acids and Ka Explain how you will make use of these concepts to determine the identity of your unknown acid Include equations where necessary Data: Include in your report the recorded concentration of base used for the titration, the mass of the unknown used and a copy of the data table (volume and pH) made during your titration. Results: Include in your report a copy of the titration curve, the first derivative plot and the second derivative plot. Tabulate your results for the calculated moles of acid, the molar mass of the acid, 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 hand written calculations to your report Discussion: Discuss the experiment and any possible sources of error Which method do you feel gave you better (or more easy to interpret) results? Justify your answer. How sure are you in the identification of your unknown? Use both your Ka and molecular weight to identify the appropriate acid Question: When the pH of the solution equals the pKa of an indicator, the solution will have an intermediate color. Estimate the pKa of both indicators (bromocresol green is the indicator that made the transition in the acidic region of the titration).
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Buffers Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. Buffer A: Calculate the mass of solid sodium acetate required to mix with 50.0 mL of 0.1 M acetic acid to prepare a pH 4 buffer. The Ka of acetic acid is 1.8 × 10–5 2. Buffer B: Calculate the mass of solid sodium acetate required to mix with 50.0 mL of 1.0 M acetic acid to prepare a pH 4 buffer. The Ka of acetic acid is 1.8 × 10–5 3. Write a reaction to show how a sodium acetate/acetic acid buffer would respond to a small amount of added strong acid. 4. Write a reaction to show how a sodium acetate/acetic acid buffer would respond to a small amount of added strong base.
Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted 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.
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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’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 very, very little. In many systems, buffers are critical. Blood plasma, a natural example in humans, is a bicarbonate buffer that keeps the pH of blood between 7.2 and 7.6. By design, a buffer is an equilibrium system. For example, a buffer can be prepared with nitrous acid, HNO2. The weak acid establishes an aqueous equilibrium as shown below. HNO2 (aq) ↔ H+ (aq) + NO2– (aq) The equilibrium constant expression is shown below.
To prepare a buffer system with nitrous acid, a conjugate base is added, such as sodium nitrite (NaNO2). The resulting system is a mixture of HNO2 and NO2– ions. The nitrous acid molecule will neutralize hydroxide ions and the nitrite ion from the conjugate will neutralize hydrogen ions. A variation of the equilibrium expression above, called the Henderson-Hasselbalch equation, is the best reference in preparing a buffer solution. For our nitrous acid/sodium nitrate buffer example, the Henderson-Hasselbalch equation is shown below.
The pH range in which a buffer solution is effective is generally considered to be ±1 of the pKa. In this experiment, you will use the Henderson-Hasselbalch equation to determine the amount of acetic acid and sodium acetate needed to prepare two acidic buffer solutions. You will then prepare the buffers and test their buffer capacities by adding solutions of NaOH and HCl.
OBJECTIVES In this experiment, you will • •
Prepare and test two acid buffer solutions. Determine the buffer capacity of the prepared buffers.
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MATERIALS Vernier pH Sensor two 50 mL burets
PROCEDURE Part I Prepare and Test Buffer Solution A
1. Obtain and wear goggles. 2. Use your calculations from the Pre-Lab Exercise to prepare 50 mL of Buffer A. Weigh out the precise mass of sodium acetate and dissolve it in 50.0 mL of 0.1 M acetic acid solution. 3. Set up two burets, buret clamps, and ring stand (see Figure 1). Rinse and fill one buret with 0.5 M NaOH solution. Rinse and fill the second buret with 0.5 M HCl solution. CAUTION: Sodium hydroxide solution is caustic. Avoid spilling it on your skin or clothing. Handle the hydrochloric acid with care. It can cause painful burns if it comes in contact with the skin. 4. Use a graduated cylinder to measure out 10.0 mL of the Buffer A solution into a 250 mL beaker and add 15 mL of distilled water. You will stir with a stirring rod during the testing. 5. Connect a pH Sensor to Channel 1 of the Vernier computer interface. Connect the interface to the computer using the proper interface cable. Suspend the pH Sensor in the pH 4 buffer solution, as shown in Figure 1. 6. Start the Logger Pro program on your computer. Open the file “19 Buffers” from the Advanced Chemistry with Vernier folder.
Figure 1 7. You are now ready to test Buffer A. Once you verify the initial pH of the buffer, you will slowly and carefully add 0.5 M NaOH solution to the buffer solution. a. Take an initial pH reading of the buffer solution. Click and monitor pH for 5–10 seconds. Once the displayed pH reading has stabilized, click . In the edit box, type 0 (for 0 mL added). Press the ENTER key to store the first data pair. Record the initial pH value. NOTE: if the initial pH is not within 0.3 pH units of 4.0 you will need to remake the buffer. b. Add a small amount of the NaOH solution, up to 0.50 mL. When the pH stabilizes click . Enter the current buret reading and press ENTER to store the second data pair. 42
c. Continue adding the NaOH solution in small increments that raise the pH consistently and enter the buret reading after each increment. Your goal is to raise the pH of the buffer by precisely 2 pH units. d. When the pH of the buffer solution is precisely 2 units greater than the initial reading, continue to add the NaOH solution in small increments until you have reached, and passed, the equivalence point of the titration. e. Click . Print a copy of the first trial. 8. Dispose of the reaction mixture as directed. Rinse the pH sensor with distilled water in preparation for the second titration. 9. Repeat Steps 7 and 8, using a fresh 10.0 mL sample of the Buffer A solution. For the second trial, repeat using the sodium hydroxide. For the third trial, titrate the buffer with 0.5 M HCl solution. 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. Record the volume of HCl that was used. There is no need to print a copy of the graph. Part II Prepare and Test Buffer Solution B
10. Use your calculations from the Pre-Lab Exercise to prepare 50 mL of Buffer B. Weigh out the precise mass of sodium acetate and dissolve it in 50.0 mL of 1.0 M acetic acid solution. If necessary, refill the burets of NaOH and HCl solution. 11. Use a graduated cylinder to measure out 10.0 mL of the Buffer B solution and add 15 mL of distilled water. Repeat the necessary steps to test Buffer B in a manner similar to the Part I trials. Print a copy of your graph of the titration using the NaOH solution. Record the volume of HCl that was used to lower the pH of Buffer B by 2 units or no significant change continues to occur.
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Buffer Report: Your lab report should include the following sections: Abstract Introduction: Data: Include all collected data in the form of a data table Results: Include the graph of pH vs. volume of NaOH added for both buffer A and buffer B Include hand written sample calculations for any calculation done in this lab. Discussion: Discuss the experiment and any possible sources of error In addition, answer the following questions as part of your discussion: 1. Write reaction equations to explain how your acetic acid-acetate buffer reacts with an acid and reacts with a base. 2. Buffer capacity has a rather loose definition, yet it is an important property of buffers. A commonly seen definition of buffer capacity is: “The amount of H+ or OH– that can be neutralized before the pH changes by +/- 1 pH unit.” Use your data to determine the buffer capacity of Buffer A and Buffer B. 3. Say, for example, that you had prepared a Buffer C, in which you mixed 8.203 g of sodium acetate, NaC2H3O2, with 100.0 mL of 1.0 M acetic acid. a. What would be the initial pH of Buffer C? b. If you add 5.0 mL of 0.5 M NaOH solution to 20.0 mL each of Buffer B and Buffer C, which buffer’s pH would change less? Explain. 4. If you wanted to carry out an experiment at ‘physiological pH’ (7.4) could you suggest an appropriate buffer system? 5. Why did the buffer B seem to handle added sodium hydroxide with relatively small changes in the pH while small quantities of HCl caused the pH to decrease substantially?
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Determination of the Ksp of Calcium Hydroxide Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. Write the molecular and net ionic equation for the reaction between Ca(OH)2(aq) and HCl(aq). 2. What is the mole ratio between [OH-] and [H+]; between [Ca2+] and [H+]? 3. The molar solubility of a slightly soluble ionic compound M2X3 is 2.8 X 10-6 M. Determine the value of Ksp. 4. Which of the saturated solutions below would have the highest [OH-1]? a. M(OH)2 Ksp = 4.25 x 10-6 b. M(OH)2 Ksp = 7.39 x 10-4 c. M(OH)2 Ksp = 2.64 x 10-3 d. M(OH)2 Ksp = 8.52 x 10-8 5. Which of the saturated solution shown in question 3 would have the lowest pH?
Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted 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.
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Lab: Determination of the Ksp of Calcium Hydroxide From Advanced Chemistry with Vernier, Vernier Software & Technology, 2004
INTRODUCTION Calcium hydroxide is a strong base that is sparingly soluble in water. In solution it completely ionizes as represented by the following equation: Ca(OH)2 (s) ↔ Ca2+ (aq) + 2OH
– (aq)
The solubility product expression describes, in mathematical terms, the equilibrium that is established between the solid substance and its dissolved ions in an aqueous system. The equilibrium expression for calcium hydroxide is shown below. Ksp = [Ca2+][OH-]2 (Equation 1) The constant that illustrates a substance’s solubility in water is called the Ksp. All compounds, even the highly soluble sodium chloride, have a Ksp. However, the Ksp of a compound is commonly considered only in cases where the compound is very slightly soluble and the amount of dissolved ions is not simple to measure. Your 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. You will do this by three methods. The first will be to determine the concentration of Ca(OH)2 by titrating the prepared Ca(OH)2 solution with a standard hydrochloric acid solution. By determining the molar concentration of dissolved hydroxide ions in the saturated Ca(OH)2 solution, you will have the necessary information to calculate the Ksp. The second method will be to utilize the pH of the saturated solution to determine the concentration of the OH-(aq) ions and utilize that information to determine the Ksp. The third method will be to utilize methods of gravimetric analysis to determine the quantitiy of Ca(OH)2 disssolved in a known volume of solution. The mass of dissolved solid can be used to determine the solubility and thus the Ksp of Ca(OH)2.
OBJECTIVES In this experiment, you will utilize three separate methods to determine the Ksp of calcium hydroxide
MATERIALS Vernier pH Probe
Burette
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PROCEDURE CAUTION: Calcium hydroxide solution is caustic. Avoid spilling it on your skin or clothing. 1. Obtain and wear goggles. 2. Obtain about 70 mL of a saturated calcium hydroxide solution. 3. Set up a ring stand, ring, filter funnel, and filter paper as demonstrated by your instructor. Filter your sample of Ca(OH)2 solution into a clean beaker. Method 1 Gravimetric Determination (begin one lab period in advance) 1. Using a 10 mL volumetric pipette transfer exactly 10 mL of the filtered solution into a pre-weighed and labeled 100 mL beaker. The beaker must be labeled with your section number and your drawer number. 2. Place the beaker in the warming oven until the end of the lab period. 3. After the drying period, determine the mass of the beaker with calcium hydroxide. Method 2 Determination by pH 1. Connect a pH Sensor to Channel 1 of the Vernier computer interface. Connect the interface to the computer using the proper interface cable. 2. Start the Logger Pro program on your computer and allow the default program to open. 3. Calibrate the pH probe following this procedure: • Use the 2-point calibration option of the Vernier data-collection program. Rinse the tip of the electrode in distilled water. Place the electrode into one of the buffer solutions (e.g., pH 4). When the voltage reading displayed on the computer or calculator screen stabilizes, enter a pH value, “4”. • For the next calibration point, rinse the electrode and place it into a second buffer solution (e.g., pH 7). When the displayed voltage stabilizes, enter a pH value, “7”. • Rinse the electrode with distilled water and place it in the sample to be measured. 4. Following calibration, measure and record the pH of the filtered solution. Method 3 Determination by Titration 1. Using a 10 mL volumetric pipette, measure out exactly 10 mL of the filtered solution into a 250 mL beaker. 2. Add 3 drops of the methyl red indicator solution. The solution will turn yellow.
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3. Obtain about 70 mL of 0.050 M HCl solution. 4. Connect a buret to the ring stand. Rinse the burette with ~10 mL of the acid before filling it with the 0.0500 M HCl solution. 5. Record the initial volume of HCl in your burette making sure to round to the 0.01 mL in an attempt to achieve the best precision possible with the given instrument). • To read the buret accurately, hold a white card behind the buret. The meniscus will appear black against the white card. Keeping your eye level with the meniscus, read the buret. 6. Using a glass stirring rod to stir the solution, semi-rapidly titrate your Ca(OH)2 sample with the HCl until the yellow color just begins to turn into an orange color. Begin adding the HCl dropwise until the solution just turns red in color. 7. Record the final volume of HCl in your burette making sure to round to the 0.01. 8. Dispose of the reaction mixture in the labeled waste bottle in the fume hood 9. Repeat the necessary steps to titrate a second sample of the filtered Ca(OH)2 solution. 10. If the volume of HCl used in the two titrations varies by more than 10%, do a third titration. PROCESSING THE DATA Calculations: Determination of Ksp using the pH of the solution. Use the pH to determine the [OH-]. With knowledge of the balanced chemical equation defining the dissolution of Ca(OH)2 utilize the [OH-] to determine the [Ca2+]. Substitute the values into equation 1 to determine the Ksp. Gravimetric determination of Ksp. Determine the mass of dissolved solid. Calculate the solubility of the solutions (M) using the mass of dissolved solid, volume of solution used in the analysis and the molar mass of calcium hydroxide. Substitute the molar solubility into equation 1 to determine the Ksp. Determination of Ksp through titration. Using the concentration and volume of HCl used, determine the [OH-] in the saturated solution. With knowledge of the balanced chemical equation defining the dissolution of Ca(OH)2 utilize the [OH-] to determine the [Ca2+]. Substitute the values into equation 1 to determine the Ksp.
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Determination of the Ksp of Calcium Hydroxide Lab Report: Your report for this lab should include the following sections: Abstract Introduction: Data: Include 3 data tables, one for each method used to determine the Ksp Results: Calculate the values of Ksp for each method. Construct a data table summarizing the results. Provide an average of the results. Discussion: Use the data in Appendix II of your test book to determine the accepted value for the Ksp. Calculate your percent error with respect to each result. 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. In addition, answer the following question as part of your discussion: 1. Predict how each of the following “mistakes” would affect the value of the measured Ksp by circling the correct response and providing reasoning for your response. 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. 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 2. Write the solubility product expression for silver chromate. 3. If one has a solution of 0.10 M silver chromate and it is diluted by a factor of 2, what is the new concentration? 4. If the dilution of 0.10 M silver chromate by a factor of 2 is carried out five times, what is the final concentration? 5. The value for the Ksp of silver chromate is 1.1 X 10-12. In a saturated solution of silver chromate, the silver ion concentration is found to be 2.5 X 10-4. What must the chromate ion concentration be? 49
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. What is ΔG? What does the sign of ΔG tell you about the spontaneity of a reaction? 2. A researcher wants to make a solution of AgCl and water at 75 °C. For solid AgCl at 75°C, Ksp = 1.5 x 10-5. a. Calculate the free energy change associated with making a saturated solution of AgCl in water at 75 °C. b. How many grams of AgCl will dissolve in 1.0 L of water at 75 °C? 3. The Ksp for Ag2CrO4 is 9.0 x 10-12. If 200 mL of 0.0050 M AgNO3 is combined with 300 mL of 0.0020 M K2CrO4, will a precipitate form? 4. Do you expect the dissolution of KNO3 to be endothermic or exothermic? Use appendix II in your text to calculate ΔHdissolution. 5. Do you expect the dissolution of KNO3 to have a positive or negative ΔS? Use appendix II in your text to calculate ΔSdissolution.
Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted 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.
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Lab: Thermodynamics of the Solubility of Potassium Nitrate In this experiment, you will measure the solubility of KNO3 as a function of temperature. The data collected will be used to determine the Ksp, enthalpy, entropy and free energy of dissolution. When a salt dissolves in water it will dissociate into ions. In aqueous solution potassium nitrate (KNO3) dissociates according to the following reaction. KNO3(s)
K+(aq) + NO3-(aq)
As the concentration of dissolved K+ and NO3- increases, the rate at which the ions will recombine into solid potassium nitrate, KNO3, also increases. At one set of ion concentrations the rate of dissolution will equal the rate of precipitation. At this point the reaction is said to be at equilibrium. The solution is now considered saturated. An equilibrium expression, Ksp, for this process is shown in equation (1). Ksp = [K+][NO3-] (Equation 1) The value for Ksp is characteristic of each compound and changes with the temperature. Thermodynamics may be used to understand what energy changes occur when a salt dissolves in water. The energy difference between the solid salt and its dissolved ions is known as the enthalpy change (ΔH), and the relative disorder of the dissolved ions is an indication of the entropy change (ΔS). A positive enthalpy change will occur if heat must be added to dissolve the salt in water. The enthalpy change will be negative if the dissolution process releases heat. The entropy change for a solid salt solid dissolving in water should be positive because the dissolved ions possess more disorder than a solid salt molecule. The free energy change (ΔG) for a process will indicate if the process will be spontaneous. A negative value indicates that the process is spontaneous while a positive value denotes a nonspontaneous process. The Gibbs-Helmholtz equation, shown in (equation 2), is a mathematical expression that relates changes in free energy, enthalpy, and entropy. ΔG = ΔH-TΔS (Equation 2) ΔG can also be expressed in terms of Ksp , (equation 3). ΔG = -RTlnKsp (Equation 3) By combining equations 2 and 3, it is possible to derive an equation that relates Ksp and the Kelvin temperature to the values associated with ΔH and ΔS (equation 4). (equation 4)
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By plotting lnKsp versus 1/T, the slope of the line is –ΔH/R and the y-intercept is ΔS/R (R is the ideal gas constant 8.314 J·mol-1·K-1). The enthalpy and entropy of dissolution can be determined by evaluating the temperature dependence of Ksp.
OBJECTIVES In this experiment, 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 to calculate ΔG, ΔH and ΔS for the dissolution process.
MATERIALS Vernier Temperature Probe
PROCEDURE 1. Obtain and wear goggles. 2. Connect a Temperature Probe to Channel 1 of the Vernier computer interface. 3. Start the Logger Pro program on your computer. Allow Logger Pro to open the default program. It should list columns for collecting time and temperature. 4. Prepare a hot water bath by heating half filled 400 mL beaker on a hot plate 5. Mass ~2 grams of potassium nitrate and record the actual mass in your lab book. 6. Transfer the salt into your 10 mL graduated cylinder. 7. Add distilled water to ~2 mL and observe if the dissolution of KNO3 is an exothermic or endothermic process 8. In your hot water bath, while stirring gently with the temperature probe, heat the cylinder containing the salt and water mixture. Do not leave it in the water any longer than necessary to get the salt into solution. 9. 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). 10. Click
to begin the data collection.
11. While monitoring the temperature allow the sample to cool while slowly stirring. Record the temperature at the point when the first crystals appear. The solution will appear to be “snowing.” The solution is considered to be at equilibrium when the first crystals begin to appear. At the higher concentrations the process happens quite quickly. 12. End the data collection by clicking
.
13. Remove the stopper and add between 0.5 and 1 mL of distilled water. Be careful not to lose solid as you remove the thermometer probe from the cylinder to add the water. 14. Repeat steps 8 - 14 four more times until a total of 5 sets of data have been recorded.
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15. Disposal: Place the KNO3 solution into the container labeled “KNO3 Waste.” For easier removal of the KNO3 reheat the solution till the entire solid has re-dissolved. The water can be evaporated and the potassium nitrate reused.
PROCESSING THE DATA For each set of data, calculate the molar concentration of KNO3, and use that value to deduce both K+(aq), and NO3 –(aq) concentrations. Utilize equation (1) to calculate the Ksp for each data set. Substitute Ksp into equation (3) to calculate ΔG for each data set. Determine the natural logarithm (ln) of the Ksp at each temperature. After converting all temperatures into Kelvin, calculate the reciprocal of each temperature (1/T). Using a spreadsheet program (Logger Pro has this capability) construct a graph with the y-axis being lnKsp and the xaxis being 1/T. Determine the best linear fit of the data using linear regression and record the slope in your notebook. From your graph, determine the values for ΔH and ΔS. Equation 4 shows that ΔH of the reaction can be determined using the slope of the straight line from the graph ΔS of the reaction can be determined y-intercept of the straight line from the graph.
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Thermodynamics of the Solubility of Potassium Nitrate Lab Report: Your report for this lab should include the following sections: Abstract Introduction: Data: Include a copy of the graph with the data indicating the slope and y-intercept Results: Calculate the values of solubility, Ksp and ΔG for each data set. From the graphical analysis of your data calculate values for ΔH and ΔS Discussion: Use the data in Appendix II of your test book to calculate the ΔHº and ΔSº for the dissolution of KNO3. Calculate your percent error with respect to each result: ΔH and ΔS. Discuss the experiment and any possible sources of error In addition, answer the following question as part of your discussion: 1. Compare the signs of your experimental thermodynamic values with your expectations from the prelab. a. 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. Did you expect this process to result in an increase or decrease in disorder? How does this compare with your calculated entropy values?
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Redox Titration: Analysis of a Commercial Bleach Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. Define oxidation. Define reduction. 2. Write balanced oxidation and reduction half-reactions for the following equations. For each half-reaction, identify which substance is oxidized or reduced in each Eqn 1: MnO4 - + S2O3 2- S4O6 2- + Mn 2+ Eqn 2: MnO4 - + C2O4 2- MnO2 + CO2 3. Use the equations given in the introduction of this lab to determine the mole ratio of thiosulfate (S2O32-) to hypochlorite (ClO-). 4. 5.00 mL of commercial bleach was diluted to 100.0 mL. 25.0 mL of the diluted sample was titrated with 4.56 mL of 0.100 M S2O32-. What is the concentration of the original bleach solution? Assume the density of the commercial bleach is 1.08 g/mL. Calculate the average percent by mass of NaClO in the commercial bleach.
Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted 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.
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Lab: Redox Titration: Analysis of a commercial Bleach Solution Many commercial products, such as bleaches and hair coloring agents, contain oxidizing agents. The most common oxidizing agent in bleaches is sodium hypochlorite, NaClO (sometimes written NaOCl). Commercial bleaches are made by bubbling chlorine gas into a sodium hydroxide solution. Some of the chlorine is oxidized to the hypochlorite ion, ClO-, and some is reduced to the chloride ion, Cl-. The solution remains strongly basic. The 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. One of the best methods is the iodine-thiosulfate titration procedure. The iodide ion, I-, is easily oxidized by almost any oxidizing agent. In acid solution, hypochlorite ions oxidize iodide ions to form iodine, I2. The iodine that forms is then titrated with a standard solution of sodium thiosulfate. The analysis takes place in a series of steps: 1. Acidified iodide ion is added to hypochlorite ion solution and the iodide is oxidized to iodine. 2 H+(aq) + ClO-(aq) + 2 I-(aq) → Cl-(aq) + I2(aq) + H2O(l) 2. Iodine is only slightly soluble in water, but it dissolves very well in an aqueous solution of iodide ion, in which it forms a complex ion called the triiodide ion. Triiodide is a combination of a neutral I2 molecule with an I- ion. The triiodide ion is yellow in dilute solution and dark red-brown when concentrated. I2(aq) + I-(aq) → I3-(aq) 3. The triiodide is titrated with a standard solution of thiosulfate ions, which reduces the iodine back to iodide ions. 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 to the clear color of the iodide ion. It is possible to use the disappearance of the color of the triiodide ion as the method of determining the end point, but this is not a very sensitive procedure. Addition of starch to a solution that contains iodine or triiodide ion forms a reversible blue complex. The disappearance of this blue colored complex is a much more sensitive method of determining the end point. However, if the starch is added to a solution which contains a great deal of iodine, the complex which forms may not be reversible. Therefore, the starch is not added until shortly before the end point is reached. The quantity of thiosulfate used in step 3 is directly related to the amount of hypochlorite initially present.
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OBJECTIVE In this experiment, determine the amount of hypochlorite ion present in commercial bleach.
MATERIALS 50 mL buret
PROCEDURE 1. Use a 5-mL volumetric pipet to measure 5.00 mL of a commercial bleach solution into a 100-mL volumetric flask. Dilute to the mark with distilled water, stopper and mix well. 2. Weigh out approximately 1 g solid KI. This is a large excess over that which is needed. 3. Pipet 25.00 mL of the dilute bleach into an Erlenmeyer flask. Add the solid KI and about 25 mL distilled water. Swirl to dissolve the KI. 4. Working in a fume hood, slowly and with swirling, add approximately 2 mL of 3 M HCl. The solution should be dark yellow to red-brown from the presence of the I3complex ions. 5. Back at your bench, titrate the iodine. Titrate with a standard 0.100 M sodium thiosulfate solution until the iodine color becomes light yellow. Add one dropper of starch solution. The blue color of the starch-iodine complex should appear. Continue to titrate until one drop of Na2S2O3 solution causes the blue color to disappear. 6. Repeat the titration, beginning with step 2, two more times.
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 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.
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Analysis of a Commercial Bleach Solution: Your lab report should include the following sections: Abstract: Introduction: 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 hand written calculations to your report.
Discussion: Discuss the experiment and any possible sources of error How sure are you in the identification of your unknown?
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Pre Lab: Electrochemistry: Voltaic Cells Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): Use the table of standard reduction potentials in your text, or another approved reference, to complete the following table. An example is provided. Electrodes
Half-reactions
Zn Cu
Zn(s) → Zn + 2e 2+ – Cu + 2e → Cu(s)
Cu
2+
Net Reaction –
Zn (s) + Cu
2+
→ Cu (s) + Zn
E 2+
°
+0.76 V +0.34 V
°
E cell +1.10 V
Pb Pb Ag Pb Mg Pb Zn
Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted 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.
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Lab 6: Electrochemistry: Voltaic Cells In electrochemistry, a voltaic cell is a specially prepared system in which an oxidation-reduction reaction occurs spontaneously. This spontaneous reaction produces an easily measured electrical potential. Voltaic cells have a variety of uses. In this experiment, you will prepare a variety of semi-microscale voltaic cells in a 24-well test plate. A voltaic cell is constructed by using two metal electrodes and solutions of their respective salts (the electrolyte component of the cell) with known molar concentrations. In Parts I and II of this experiment, you will use a Voltage Probe to measure the potential of a voltaic cell with copper and lead electrodes. You will then test two voltaic cells that have unknown metal electrodes and, through careful measurements of the cell potentials, identify the unknown metals. In Part III of the experiment, you will measure the potential of a special type of voltaic cell called a concentration cell. In the first concentration cell, you will observe how a voltaic cell can maintain a spontaneous redox reaction with identical copper metal electrodes, but different electrolyte concentrations. You will then measure the potential of a second concentration cell and use the Nernst equation to calculate the solubility product constant, Ksp, for lead iodide, PbI2.
Figure 1
OBJECTIVES In this experiment, you will • • • • •
Prepare a Cu-Pb voltaic cell and measure its potential. Test two voltaic cells that use unknown metal electrodes and identify the metals. Prepare a copper concentration cell and measure its potential. Prepare a lead concentration cell and measure its potential. Use the Nernst equation to calculate the Ksp of PbI2.
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MATERIALS Voltage Probe 24-well test plate string Cu and Pb electrodes two unknown electrodes, labeled X and Y steel wool
PROCEDURE o
Part I Determine the E for a Cu-Pb Voltaic Cell
1. Obtain and wear goggles. 2. Use a 24-well test plate as your voltaic cell. Transfer small amounts of 0.10 M Cu(NO3)2 and 0.10 M Pb(NO3)2 solution to two neighboring wells in the test plate. CAUTION: Handle these solutions with care. If a spill occurs, ask your instructor how to clean up safely. 3. Obtain one Cu and one Pb metal strip to act as electrodes. Polish each strip with steel wool. Place the Cu strip in the well of Cu(NO3)2 solution and place the Pb strip in the well of Pb(NO)3 solution. These are the half cells of your Cu-Pb voltaic cell. 4. Make a salt bridge by soaking a short length of string in a beaker than contains a small amount of 1 M KNO3 solution. Connect the Cu and Pb half cells with the string. 5. Connect a Voltage Probe to Channel 1 of the Vernier computer interface. Connect the interface to the computer with the proper cable. 6. Start the Logger Pro program on your computer. Open the file “20 Electrochemistry” from the Advanced Chemistry with Vernier folder. 7. Measure the potential of the Cu-Pb voltaic cell. Complete the steps quickly to get the best data. a. Click to start data collection. b. Connect the leads from the Voltage Probe to the Cu and Pb electrodes to get a positive potential reading. Click immediately after making the connection with the Voltage Probe. c. Remove both electrodes from the solutions. Clean and polish each electrode. d. Put the Cu and Pb electrodes back in place to set up the voltaic cell. Connect the Voltage Probe to the electrodes, as before. Click immediately after making the connection with the Voltage Probe. e. Remove the electrodes. Clean and polish each electrode again. f. Set up the voltaic cell a third, and final, time. Click immediately after making the connection with the Voltage Probe. Click to end the data collection. g. Click the Statistics button, . Record the mean in your data table as the average potential. Close the statistics box on the graph screen by clicking the X in the corner of the box.
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o
Part II Determine the E for One Voltaic Cells Using Pb, Cu and an Unknown Metal
8. Obtain a small amount of the unknown electrolyte solution labeled “0.10 M X” and the corresponding metal strip, X. 9. Transfer a small amount of 0.10 M X solution to a well adjacent to the 0.10 M Pb(NO3)2 solution in the test plate. 10. Make a new salt bridge by soaking a short length of string in the beaker of 1 M KNO3 solution. Connect the X and Pb half cells with the string. 11. Measure the potential of the X-Pb voltaic cell. Complete this step quickly. h. Click to start data collection. i. Connect the leads from the Voltage Probe to the X and Pb electrodes to get a positive potential reading. Click immediately after making the connection with the Voltage Probe. j. Remove both electrodes from the solutions. Clean and polish each electrode. k. Set up the voltaic cell again. Connect the Voltage Probe as before. Click immediately after making the connection with the Voltage Probe. l. Remove the electrodes. Clean and polish each electrode again. m. Test the voltaic cell a third time. Click immediately after making the connection with the Voltage Probe. n. Click to end data collection. o. Click the Statistics button, . Record the mean as the average potential and then close the statistics box on the graph screen by clicking the X in the corner of the box. 12. Repeat Steps 8–11 using copper a copper half cell and the unknown metal half cell. Part III Prepare and Test Two Concentration Cells
13. Set up and test a copper concentration cell. a. Using your 0.10 M Cu(NO3)2, prepare 20 mL of 0.050 M Cu(NO3)2 solution by mixing 10 mL of 0.10 M Cu(NO3)2 solution with 10 mL of distilled water. b. Set up a concentration cell in two wells of the 24-well test plate by adding 5 mL of 0.050 M Cu(NO3)2 solution to one well and 5 mL of 1.0 M Cu(NO3)2 solution to a neighboring well. Use Cu metal electrodes in each well. Use a KNO3-soaked string as the salt bridge, as in Parts I and II. c. Click to start data collection. d. Test and record the potential of the concentration cell in the same manner that you tested the voltaic cells in Parts I and II.
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Voltaic Cells Lab Report: DATA TABLE 1: Determine the E for a Cu-Pb voltaic cell Voltaic Cell (Cu / Pb)
Measured Potential (V)
Trial 1 Trial 2 Trial 3 Mean
1. Compare the average cell potential, for your Cu/Pb cell, with the E°cell that you calculated in the pre-lab exercise. Explain why your cell potential is different from the text value.
DATA TABLE 2: Determination of the E for unknown metal Mx using Pb. Voltaic Cell (Pb / Mx)
Measured Potential (V)
Metal of (+) Lead
Metal of (–) Lead
Trial 1 Trial 2 Trial 3 Mean
DATA TABLE 3: Determination of the E for unknown metal Mx using Cu. Voltaic Cell (Pb / My)
Measured Potential (V)
Metal of (+) Lead
Metal of (–) Lead
Trial 1 Trial 2 Trial 3
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Mean
2. The unknown metals X and Y were either magnesium, silver, or zinc. Use the text value for the reduction potential of Pb or Cu and the measured cell potentials for the unknowns to identify X and Y. The cell potential for Mx can be determined using the following equation EMx= Ecell - Eostandard.
DATA TABLE 4: Determination of the E for a copper concentration cell Voltaic Cell (0.10 M / 1.0 M)
Measured Potential (V)
Trial 1 Trial 2 Trial 3 Mean
3. (Part III) Use the Nernst equation to calculate the theoretical value of E of the copperconcentration cell and compare this value with the cell potential that you measured.
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Synthesis of Acetaminophen Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. Write the balanced chemical equation for the synthesis of acetaminophen from paminophenol and acetic anhydride. 2. If 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 reactant and product online. 3.
Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted 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.
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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 as a safe medication though 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 which has a hydroxyl functional group (-OH) attached to one side and an amide functional group on the other (see Figure 2).
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. See if you can 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. After recrystallization, impurities are usually present in much lower concentrations as compared to the crude product.. One measure for assessing the purity of a substance is by determining its melting point. Generally a melting point range is reported. The range is reported from the temperature at which the first drop of the liquid appears to the temperature at which the sample is completely melted and only a clear liquid is present. A pure solid generally melts over a narrow range of temperatures; marking a characteristic melting temperature range (the substances melting point). An impure substance may melt over a wide range of temperatures usually at a temperature lower than that of the pure substance. The greater the amount of impurities, the greater the decrease in the apparent melting point. Here you will measure the melting point of your purified product and compare it to the accepted melting point of acetaminophen which is 169.5-171°C. 66
Figure 3. Synthesis reaction of acetaminophen.
Safety Precautions: • • • • •
Wear safety glasses or goggles at all times in the laboratory. Acetic anhydride is corrosive and its vapor is irritating to the respiratory system. Avoid skin contact and inhalation of the vapors. In the event of skin contact, rinse well with cold water. If the vapors are inhaled, move to an area where fresh air is available. Phosphoric acid is corrosive. Avoid skin contact. In the event of skin contact, rinse well with cold water. p-aminophenol is harmful by inhalation and by contact with the skin. In the event of skin contact, rinse well with cold water. If the vapors are inhaled, move to an area where fresh air is available. NOTE: Don't use your aspirin for a headache! Its purity is not assured.
Procedures: Synthesis of Acetamenophen 1. Weigh out 1.5 g of salicylic acid. Place it in a 125-mL Erlenmeyer flask. 2. Add 25 mL of water and 25 drops of 85% H3PO4. 3. Gently heat to dissolve. 4. Move to the hood once all the solid is dissolved. Working in the hood, while swirling the flask add 2 ml of the acetic anhydride. Allow to react for 5 minutes. 5. Place the mixture on ice and allow to crystalize for 20 minutes. If crystallization does not occur, it may be necessary to scratch the side of the flask with a stir rod or add a small seed crystal of pure acetaminophen. 6. Collect the crystals by vacuum filtration using a Buchner funnel. 7. Wash the crystals twice with 5 mL portions of ice cold water. Recrystallization of the Acetaminophen 1. Place the crude acetaminophen crystals in a clean 150 mL beaker. Add 20 mL of distilled water and heat on a hot plate until all of the solid has dissolved. If the solution starts to boil and undissolved solid still remains, add more water, a few milliliters at a time. 2. Remove the beaker from the heat and allow the solution to cool. When crystals begin to appear, put the solution on ice, cooling for 20 minutes. If crystallization does not occur, it may be necessary to scratch the side of the flask with a stir rod. 3. Collect the crystals on a preweighed piece of filter paper using the Buchner funnel. Wash the product with two 5 ml portions of ice cold water. Allow to dry for 10 minutes under vacuum. 4. Weigh the product. It may be necessary to dry for a few more minutes. 67
Determine the Melting Point of the Aspirin Sample 1. Fill a capillary melting point tube to a depth of 0.2 cm with the recrystallized product as demonstrated by your TA. 2. Place the capillary tube in the melting point apparatus. Determine its melting point (Your TA will demonstrate the use of this apparatus). Pure acetaminophen melts at 169.5171°C. 3. The aspirin sample should be labeled with your name, the mass of the aspirin, the percent yield, and its melting point.
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Synthesis of acetaminophen worksheet _________________ 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)
_________________ °C
Accepted melting range of acetaminophen
_________________ °C
Measured melting point of acetaminophen
What does the comparison of your measured melting point to the know melting point tell you about the purity of your product?
During the crystallization of acetaminophen, 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?
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