clewett- Web viewChem 201L is an inquiry-based lab course. In Chem 201L this semester, the students...

Laboratory Manual for Chem 201Used as a Companion to Chemistry, 7th Edition, by Zumdahl & Zumdahl

Department of ChemistryUniversity of Nevada, Reno

Contributors: Heather S. Clewett and Jason Shearer

With thanks to Thomas Brown, Danil Kaliakin, Kelly Chen, & Kristy Peck, UNR Chemistry Department TA’s, for significant assistance and feedback.

Contents

Introduction 3About this Lab Course 3

Course Objectives & Philosophy 4

Necessary Supplies 4

Lab Management - Things You Need to Know to be Successful 5

Lab Requirements and Grading 6

Pre-Lab Requirements 7

In-Lab Requirements 9

Post-Lab Requirements 10

Notes on Grading 12

Experiment 1: Density Investigations 13

Skill-Building Inter-session 22

Experiment 2: Reactions with Copper (2 week lab) 25

Experiment 3: Qualitative Analysis (2 week lab) 39

Experiment 4: Titrations of Acids 49

Experiment 5: Calorimetry and Thermodynamics 58

Experiment 6: Redox Titration 67

Experiment 7: Atomic Spectroscopy 76

Experiment 8: Particle in a Box Simulations 87

Appendix: Sample Lab 98

2

About this Lab Course

Welcome to Chem 201L!

Chem 201L is an inquiry-based lab course. In Chem 201L this semester, the students fundamentally assume ownership of the investigation; your TA will be a valuable resource but will not tell you what to do or how to do it. Instead, your TA will support you with guiding questions and may also recommend that you discuss your question(s) further with your group members. You will not be abandoned, but you will come to rely more on yourself and your peers, spending time and effort to think through the problem(s) you have encountered, rather than passively accepting an expert’s answer.

As a student in an inquiry lab, you will have input in deciding the question to be investigated and the hypothesis statement. write your own procedures. create your own data tables. analyze your data yourself, including looking up necessary formulas in your text. compare your data and analysis with the class as a whole and actively exchange ideas

with your fellow lab members to make sense of the data and any discrepancies. write conclusions that provide answers to the beginning questions (these answers are

also called claims) and explain how the evidence supports the claim, demonstrating your comprehension of both theory and how the data relates to theory.

be graded more on the quality of your thinking and writing than on whether or not you arrive at the “right” answer.

You will most likely also make some mistakes (because you aren’t following a scripted “right” way of doing

things).

Once again, inquiry-based labs are not focused exclusively on “getting right answers.” Answers in the science community commonly undergo modifications as more evidence comes to light, and a three-hour student lab is not expected to achieve the same degree of accuracy as years of research using high-precision instruments.

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Course Objectives & Philosophy

One goal of this course is to prepare you to start undergraduate research. Expect to be challenged. Plan to be highly organized. Manage your time efficiently and complete lab write-ups promptly, while information is still fresh. Practice your best active reading and note-taking strategies as you read/study this lab manual. (These are the most common suggestions made by previous years’ Chem 201 students, in fact.)

Interaction with peers is encouraged and even necessary for completion of lab work. This is representative of professional chemical practices in industry and academia. Discussions are always encouraged; actual word choice should never be copied directly, of course.

Students will: demonstrate safe laboratory practices and correct handling of equipment, supplies, and

waste. follow mainstream science conventions when reporting and interpreting data as well as

when performing calculations. exercise scientific ethics, including thorough individual preparation for lab experiments,

use of bound lab notebooks and recording all data in ink. demonstrate conceptual and practical understanding of chemical theories, practices and

calculations pertinent to lecture topics.

Necessary Supplies

**Find the Lab Schedule for the day of the week that your lab meets. Print it out and refer to it throughout the semester. It will be your main guide for due dates of all lab assignments.

Use the lab schedule as a guide, but know that pre-lab and post-lab requirements will vary slightly with each lab (and some labs span 2 weeks). Read the instructions thoroughly for every lab and ask your TA questions any time you need clarification.

Visit WebCampus and print out the lab and any additional materials that will be useful to you, for each lab.

Additionally, you will need to purchase a carbon copy lab manual (these are available at many locations in town and online; 100 carbonless, graph-paper marked pages are recommended.) Your lab manual must be numbered and include a table of contents (you may add these by hand, if needed).

Bring a pen and calculator to every lab class. You may also wish to bring a ruler. (Note that pre-lab and post-lab work may be completed in pencil, but all original data must be recorded in pen.)

Wear appropriate personal protective gear as described in the Lab Safety Assignment.

4

Lab Management - Things You Need to Know to Be Successful

Be prepared to spend an average of 3 hours per week outside of lab (in addition to the 3 hours per week spent in lab) working on lab assignments.

This lab complements the lecture; consistent attendance at lecture (and reference to both lecture notes and your textbook) will increase your performance in lab.

You will occasionally be dependent on other students’ results; you will not always be able to collect all your own data; cooperation and collaboration are part of the real world. Get to know your lab mates, and support one another.

Plan to consult with your lab mates and TA frequently. Ask questions! Share your ideas!

Failure to come prepared to lab in appropriate clothing, with safety goggles and completed pre-lab may prevent you from participating in lab, resulting in an automatic zero for that lab. This policy extends to late arrivals to lab, as well.

There are no make-up labs.

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Lab Requirements & Grading“It doesn’t count--and can’t be graded--until it’s on paper.”

Overview

All Lab Reports are to have all of the following Headings:

(Pre-lab)

Pre-lab Exercises

Original Beginning Question

Selected Question (Note: This will be determined as a result of whole class discussion.)

Hypothesis

Procedure

Safety Notes

(In-lab)

{possible adjustments to Procedure and Safety Notes}

Data, Observations & Calculations

(Post-lab)

Results

Claim(s)

Post-lab Question(s)

Reflection

The requirements for each of these headings are discussed more thoroughly on the following pages.

Also, please view the SAMPLE LAB provided in the Appendix (and also available on WebCampus) for a template/exemplar.

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Pre-Lab Requirements

READ the laboratory description and any supplementary materials. Lab descriptions are dense; you need to spend quality time and practice active, thorough, close reading.

Complete the Pre-lab Exercises, as indicated in the laboratory description.

Construct a productive, Original Beginning Question (details for this process are provided below).

Consider the goals of the laboratory experiment and the suggested “Beginning Question Seeds,” provided on all lab experiments except Experiment 1, which instead provides the Beginning Question. Note how the “Seeds” address the Goals. You may refine one of the provided “Beginning Question Seeds” or synthesize a Beginning Question of your own.

Requirements for your own, Original Beginning Question:

It is your original creation or original spin on a “Seed” question.

It is a Productive question and lends itself to the creation of a testable hypothesis.

Examples of Productive Beginning Questions for Experiment #1: What factors (phase, mass, instrument precision, etc.) hinder or

improve accuracy of measurements? How does the test sample volume of an unknown liquid affect the

accuracy and/or uncertainty? How does the accuracy of the density of air measurement vary with

the number of “syringe pumps” used to evacuate the bell jar?

Examples of Non-productive Beginning Questions for Experiment #1:Simplistic, “single answer” questions: What is the density of air? What is my unknown liquid? What is the percent error in my calculation?Procedural questions: How do I use a pipet? “Why” questions (can not be answered by performing the experiment): Why are some materials more dense than others?

Hopefully, it is a question that you are personally interested in investigating.

Note that the ultimate choice of experiment will be driven by the class discussion of these Beginning Questions.

Contribute to your section’s discussion of the Beginning Question; your TA will give you specific directions on when/how to discuss the Beginning Questions and finalize the Selected Question.

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Pre-Lab Requirements, continued

It is each student’s responsibility to ascertain the final Selected Beginning Question and record it in his/her lab notebook.

Write a thoughtful Hypothesis that addresses the Selected Beginning Question and can be tested during the lab.

Provide a section for a brief (practice) Procedure in your lab notebook describing how you will collect data to test your hypothesis. Provide enough detail so that another student could pick up your lab notebook and be able to conduct your experiment.

Provide a section for (practice) Safety Notes specific to this particular lab and your particular procedure.

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In-Lab Requirements

Students must come to lab with their individual pre-lab work stapled for submission, ready to engage in class discussion about the pre-lab, and prepared for productive work on the current topic.

In lab, table groups will discuss their procedures and safety notes. Students will then engage in a class-wide discussion to determine a final procedure and set of safety notes. Note that the final procedure may require different groups to use different quantities of material or systematically alter other factors to test the results. Students will assign their own work groups and/or division of labor. Individual students may write changes to their original procedure or add specific notes that apply to their group’s data, as needed for their assigned tasks.

Students will develop data tables on the board and/or on the lab computer and fill them in as results are obtained. A benefit of electronic data tables is that they may be emailed or copied onto a zip drive at the end of the lab. Students may also use Excel or similar programs to generate a graph while in lab.

During the lab, each student or group of students will record their own Data and Observations in their lab notebook in pen. Every student is also responsible for his or her own Calculations, equations, chemical information, etc. Copying calculations is plagiarism.

Students examine the collected data to assess patterns and relationships. If an inconsistent value is found (and time permits), students may elect to repeat an experiment.

All students stay in lab until all data is collected and engage in a class-wide discussion about the outcome(s) of the experiment. Students compare observations and insights; they engage in scientific argumentation over their proposed claims and explanations. The TA facilitates the discussion but does not tell students what they should have learned from the experiment.

Post-Lab Requirements

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If not already completed during lab, students will complete the “Calculations portion” of their Data, Observations & Calculations section by writing out in full one representative calculation of each type needed in the lab. This full representation includes writing the formula to be used and showing one sample calculation using data from the lab. For example, Experiment #1 requires 5 calculations: calculation of Density, calculation of Density uncertainty, calculation of average Density, calculation of standard deviation, and calculation of percent error. All calculations must show correct units throughout. All calculations for Experiment #2 forward must show attention to significant figures and correct significant figures in the final answer. Note that it is traditional to carry an extra digit or two throughout intermediate calculations to avoid excessive rounding error.

Also note that students are to select representative calculations independently; exact duplication of this portion of the lab write-up will be construed as plagiarism.

It is anticipated that students will increase their efficiency by using spreadsheet programs to automate their calculations; furthermore, these spreadsheets may be shared among students in a given lab section via Google docs or similar. Each student is required to provide a Results section (which may simply state, “See attached data tables.”) to summarize the results (data and/or calculations of interest) of the lab. The scope of these Results must be sufficient to answer the Beginning Question and substantiate the Claim (described next) that will be crafted in response to the student’s Hypothesis.

Students will provide at least one Claim, or synthesis statement, addressing the Beginning Question and the student’s Hypothesis.

An appropriate format for each Claim is: a statement relating to the Beginning Question and/or Hypothesise.g., “In contrast to my hypothesis, the accuracies are not equivalent for all phases.” or “In contrast to my hypothesis, the liquid unknowns did not have the highest precision.”

a couple of sentences citing the specific Results and any applicable theory to support the Claim

e.g., “Accuracy is assessed by percent error, with low percent error corresponding to high accuracy. Inspection of the Results table shows that air (gas phase) had by far the highest percent error, 23%, more than twice as large as the second highest (solid unknown #4). The accuracy of air density measurements may have been compromised by low quality equipment, as we also note that all experimental densities fall below the reported true density for air in Reno, suggesting a systematic error.”

Claims are to be the result of independent student thought and interpretation of the data, though they may be guided by the whole class discussion. Copying another student’s wording (whether exactly or with minor paraphrasing) will result in zero points for both students.

Claims must be written concisely; points will be lost for writing exceeding 1 page.

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Post-Lab Question(s) should be answered completely.

The Reflection explains what you learned about chemistry as a result of each lab. You will need to consider thoughtfully the specific questions posed for each lab experiment; these tie into the original goals of the lab. It is essential that you completely, thoughtfully, yet concisely answer ALL questions asked. As with the Claims section, points will be lost for writing exceeding 1 page.

Notes on Grading

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A note about grading inquiry labs: Inquiry labs are fundamentally different from traditional labs. Students must actively engage and think for themselves as they discover scientific principles based on data they have collected, in experiments they have designed. It is OK to make mistakes...it is part of the learning process and consistent with real scientific research to do so. There is little-if-any emphasis on “getting the right answer.”

Each lab report will be worth 30 points. Lab TAs will use the rubric shown below and will provide comments (to give you a “path to improvement”) as needed.

Rubric Categories 0 1 2 3 4 5

1. Quality of pre-lab preparationThis includes completion/effort/correctness of the Pre-lab Exercises, thoughtful construction of student’s Original Beginning Question and Hypothesis, correct use of personal protective equipment, and overall readiness for the subject material of the lab experiment.Comments:

2. Completion of entries for selected question, procedure, and safety notes sectionsThis is a simple completion grade.Comments:

3. Quality of the individual data, observations and calculationsThis includes use of pen to record data while in lab, overall performance while in lab, use of data tables to organize individually collected data, thoroughness of qualitative observations (when required), and demonstration of all calculations.Comments:

4. Quality and thoroughness of resultsAre your Results complete? Did they address all requirements stated in the lab? Are the calculations correct? Do they include units and appropriate attention to significant figures? Are they presented attractively/logically?

Comments:

5. Quality of claims and use of evidence to support the claimsDo your Claims correlate directly to the Beginning Question? Is sufficient evidence presented to support the claims? Is evidence organized and presented logically? Are results analyzed and interpreted correctly?

Comments:

6. Quality and accuracy of responses to post-lab questionsAre all questions answered correctly and completely?Comments:

7. Quality of the reflectionDoes the Reflection completely, thoughtfully, and logically answer all of the questions asked?Comments:

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Experiment 1: Density Investigations


13

Overview: In Experiment 1, students will collaborate to investigate density. Density is expected to be a familiar topic; however, students will gain experience with the format and expectations for inquiry laboratory experiments in Chem 201. Equally important, students will gain insight into precision and accuracy (concepts that are frequently misunderstood) and will relate uncertainty to significant figure conventions.

Materials to investigate include air (using a small bell jar apparatus), two unknown liquids, and six unknown solids (where all except Metal Unknown #4 are pure elemental metals). Students are encouraged to perform measurements with a range of instruments, which may provide a range of precision values. Students may also consider other factors to vary systematically, (for example, the amount of substance measured). At the completion of the lab, the TA will provide the density of the unknowns, (except for honors classes, where the identity will be provided and students will research the density), and students will qualitatively and quantitatively compare their calculated density values to the known values.

In a traditional science lab, you would be given a procedure and theory to read, and you would follow the steps, more or less “cookbook style” until you got a result, which you would report in your lab notebook or perhaps compare with the expected result.

This course is going to be structured significantly differently. Not only will you be “doing” science, you’ll be doing it the way that scientists do, as a primary investigator, not always sure what the right answers or right procedures are. The goal of an inquiry-based experiment is that you learn how and why science experiments are conducted, so that you take ownership of the process of learning and experimenting rather than just following a routine to prove an established principle.

Background: It is anticipated that all students in Chem 201L are well acquainted with the formula and calculations for density. (If you have a “hole” in your academic background, please speak with your TA as soon as possible.)

Background - Accuracy, Precision, and Uncertainty: The concepts of accuracy, precision, and uncertainty are commonly misunderstood; each carries implications about the quality of the measurement performed.

Accuracy: How close you are to the true value. When multiple measurements are performed, the average value is used to determine accuracy; accuracy is expected to increase with increased number of measurements. In this lab course, we will report accuracy as a percent error; see page 20 for the formula and further information.

Precision (First definition): How close your values are to one another (internal consistency). Precision often refers to a set of measurements; high precision sets of measurements will have low standard deviations; see page 17 for the formula for standard deviation.

Precision (Second definition): How detailed and/or sensitive your measurements are (reflects quality/sensitivity of instrument used to perform measurement). Precision may also refer to a single measurement. For example, a scale that measures to the thousandth of a gram yields more precise measurements than one that measures to the tenth of a gram; 2.267 g is thus a


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more precise value than 2.3 g because it gives more detailed information.Uncertainty: The potential difference between the measured value and the true value, as judged by the individual performing the measurement, at the time of the measurement. Consider how you might measure the thickness of a quarter using a metric ruler; also imagine how you would report your measurement. Would you be confident about reporting, say, 1.5 mm? You might prefer to estimate 1.5 mm ± 0.3 mm; this provides you with a range in which you may be reasonably confident the true value may be found. If you are measuring the same quarter with calipers, you might estimate 1.74 mm ± 0.01 mm. In both cases, the format best estimate ± uncertainty is used, where the choice of uncertainty is made at the discretion of the individual performing the measurement. If you have never used calipers before, and you’re not sure if you’re using them correctly, you may have a bigger range than ±0.01 mm. The concept of uncertainty also extends to include the possibility of systematic error (e.g., the calipers may have been damaged, or have “electronic drift” and consistently yield values 0.13 mm too large). By convention, uncertainty from a digital reading is taken to be ± (1 times the last digit

place value reported); for a mass of 145.086 g, we would report the uncertainty to be ± 0.001 g. (This uncertainty may be increased is systematic error is suspected, such as if the instrument has not been calibrated.)

Otherwise, uncertainty is gauged by the individual performing the measurement and is simply an estimate intended to convey the range of values judged necessary to span between the reported measurement value and the true value.

Goals: Accurately measure and record mass and volume; reasonably assess and record the

uncertainty of your measurements. (This latter is perhaps the more important of the two.) Calculate density and estimate the uncertainty of your density measurement based on the uncertainty of your mass and volume measurements.

Assess the precision of a body of data: compare calculated density values to one another; evaluate trends in class data.

Assess the accuracy of a body of data: compare calculated density values to literature values. (Honors scholars will research literature values themselves with TA guidance.)

Explain how significant figure conventions relate to reported uncertainties of measured value.

Extra Equipment List (Above and Beyond your Lab Drawer Contents):

Bell jar apparatus (see figure at right) Air in room

Solid metal cube unknowns(about 1 inch long per side)

Liquid unknowns

Metric rulers

Graduated pipets (1 mL, 2 mL, 5 mL, 10 mL, and a limited number of 15 mL,


15

25 mL, or 50 mL) may be checked out from the Stockroom, if desiredPre-Lab Requirements

Pre-lab Exercises: (Expected length in lab notebook: 2-3 pages)Record the heading “Pre-lab Exercises” into your lab notebook (this heading applies to questions 1-8 below).

1. Record the text below into your lab notebook. Paraphrase if you wish.

“When reading values from lab equipment, it is customary to use all digits indicated by the tick marks on the measurement device AND estimate one more digit as an indication of where the measured quantity falls within the tick marks.”

“All data measurements must include units (in a table heading or on each entry), and must be recorded in ink.”

“When measuring liquid volumes, it is customary to record the volume as read from the bottom of the meniscus.”

2. Copy (or print and tape, etc.) each figure shown below into your lab notebook. Next to the figure, record the value as you would enter it into your lab notebook.

a. b.

25

(25 mL graduated cylinder) (graduated 1 mL pipet)

3. It is an understood convention that the final digit of a recorded, measured value is an estimate and contains an inherent, implied uncertainty, typically ±1 of the final digit’s place value. Thus, 2.24 cm is understood to have an uncertainty of ±0.01 cm, 0.502 g is understood to have an uncertainty of ±0.001 g, and 13 mL is understood to have an uncertainty of ±1 mL.

State the implied uncertainty for each of the following values:Note that uncertainties should include units.a) 1.3 g/mLb) 1.30 g/mLc) 1.300 g/mL

20

25

0.3

0.5

0.6

0.4


16

4. State a definition for precision and evaluate which of the values listed in question 3 implies the greatest precision. Explain your reasoning.

5. Explain what is inconsistent about the following calculated value and estimated uncertainty:4.58372 g/mL ± 0.1 g/mL

6. Calculate the density of a liquid sample with a volume of 18.8 mL and a mass of 21.256 g. Show your work.

7. a) Calculate the density of the liquid sample in question 6 if the volume was actually 18.9 mL (and the mass was still 21.256 g). Show your work.

b) Calculate the density of the liquid sample in question 6 if the mass was actually 21.257 g (and the volume was still 18.8 mL). Show your work.

8. Suppose students are told to find the mass of a coin. The following values are recorded:coin # mass (g)

1 3.0132 4.9763 5.6414 2.5455 4.9666 2.3707 5.624

Calculate the average and the standard deviation for this data set. What might you say about precision, if anything? What might you say about accuracy, if anything?

ALSO: Record the following section label and text into your lab notebook: (Italics not required.)

Selected Question: “Do the precision and/or accuracy of density measurements depend on the phase of material being measured?”(In other words, will solids, liquids, or gases have more consistent--thus, more precise--measurements than the other phases; alternately, will one or more of the phases produce average densities that are significantly closer to the accepted literature values--thus, more accurate values?)

AND: Record the section label “Hypothesis:” into your lab notebook. Write your educated guess for what your findings will be. State a reasonable explanation and/or give some background in support of your hypothesis.

Standard Deviation=√∑ (x−xi)2

N−1

where x is the average of all values¿N isthe number of recorded values


[1]

17

FOR THIS EXPERIMENT ONLY: Students will NOT construct an Original Beginning Question. The Procedure and Safety Notes sections will be completed in class. They are not required for the pre- lab assignment.

In-Lab Requirements Be prepared to work together as a class and stay for the full 2 hours, 50 minutes.

Pre-Lab Discussion:You will turn in a stapled copy of your pre-lab work at the start of class. For the first 10 or so minutes of class, discuss your pre-lab work with the class. You may be called on to take a leadership role in this discussion (stand at the board, share your values, help brainstorm ideas, etc.). It is your responsibility to ask for help on anything that has been problematic for you; if you don’t understand what you will need to do (and why), you put the whole class at risk of having faulty data. The pre-lab work is foundational for the work in the lab; thus, it is critical that you understand it thoroughly, as well. You are always welcome to make additional notes and/or corrections to your own copy of the pre-lab work as the discussion progresses.

Spend about 10 more minutes talking as a whole class. Consider and agree on:

What data do we need to address the beginning question?

What are the proposed hypotheses that have been developed in response to this question? (Students’ hypotheses do NOT need to agree throughout the class; however, this is an opportunity to make some changes, if desired. Verify that your hypothesis will be testable based on the data being collected and will result in potential claims.)

What procedure(s) will we use to collect the necessary data? (Include number of trials; also agree on how uncertainty will be reported--each measured value that is recorded for this lab MUST include an explicit estimate of uncertainty; finally, be sure to discuss how volume of air will be measured in the bell jar.)

How will we divide labor? (Make efficient use of time and human resources; if you divide work effectively, you will probably get to leave early.)

What are the safety concerns?

What form and final units will be used to report the densities? As a class, write the agreed upon procedure and safety notes on the board.

Also, as a class, prepare data tables (on the board and/or in an Excel spreadsheet) before you begin collecting data. Give space to record uncertainties (and units, of course) for all original measurements.

A note about neatness:

You are expected to write neatly and legibly in your lab notebook at all times.

A failure to submit legible copies of your work may result in significant reductions in your grade.

Your TA has the authority to require you to type all or some of your lab report, at any time and for any reason, in addition to submitting your original carbon copy pages.


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In-Lab Work: (Expected length in lab notebook: 2-3 pages through the end of the calculations)

Each student in the class will need to provide a notebook heading and entry for the Procedure and Safety Notes and complete the entry with the information developed in class.

Each student will also need to write a heading for Data, Observations & Calculations and draw his or her own data collection table, modeled after the one on the board, but almost certainly smaller.

Each student will record his or her own original data into his or her own table, in ink. Perform any required calculations, remembering to show formulas being used (even simple ones such as the density formula). Because it is unfamiliar at this stage, a “propagation of error” formula is provided for calculating the uncertainty of density based on the uncertainty of mass and of volume. Even though the formula is provided, do show the equation and calculations clearly.

The formula for “propagation of uncertainty” for density (where uncertainty is denoted by ) Δis:

ΔDD

=√( Δmm

)2

+( ΔVV

)2

It should be clear that students will need to solve the above equation for ΔD and then plug in the experimental values for m, m, V, V, and the calculated value for D to obtain D.Δ Δ Δ

Each student will then contribute his or her original data and uncertainty estimates, as well as all calculated densities and propagated uncertainties, to the whole class data tables. The class data as a whole may be printed from the Excel spreadsheet and attached to the lab notebook.

Once students are able to evaluate the whole class data, the following calculations are to be performed for each substance (these calculations may be performed in lab or after lab; calculations may be performed using Excel but at least one representative calculation must be shown in full in the lab notebook, and all values are to be recorded in the lab notebook):

Calculate the average density and standard deviation of density using all available data, where standard deviation simply follows equation [1] on page 17.

Given the true values from literature sources and/or suppliers (or as researched for Honors sections), calculate the percent error of the average experimental value for each substance.

percent error (% )=¿ (true value−experimental value )∨ ¿true value

×100%¿


[2]

[3]

19

Clean up; wipe down your lab bench area, etc. If you finish your work early, find other groups to help; you can’t perform the above calculations without them.

Post-Lab Discussion (takes place in lab):When all class data has been recorded in all data tables, the TA will provide the true density values and identities of the unknowns measured. Be sure that you make note of these values and identities.

Note:

Your lab safety goggles must stay on until everyone in the lab has put away all glassware and locked their lab drawers.

Students who wish may choose to stay for a class discussion and/or lab completion round table. This will be student-directed and may address any trends observed in the data or calculations, the methods for determining uncertainties, the precision and accuracy of measurements …or anything else that is helpful to you. You are encouraged to stay and be an active participant in this discussion; if you are able to complete your lab report during this time, so much the better!

Checklist before you leave lab: You have all the class data and the information from the TA about the unknowns

recorded in the lab notebook and/or on electronic file. You understand how to assess precision of class measurements. You understand how to assess accuracy of class measurements. You perceive and understand the connection between significant figure conventions and

measured uncertainties. (ASK your TA before you leave, if this is not yet clear.) Your lab drawer is securely locked.


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Post-Lab Requirements (MAXIMUM length in lab notebook: 2 pages; failure to be concise will result in loss of points.)

Results:Provide data tables (class-wide tables are okay) showing densities, density uncertainties calculated per Equation [2], average densities and standard deviations, and percent error for each sample. You may also include additional information, including raw data values, other calculations that support your claims (below), etc. Honors students must provide sources (references) for literature values of unknowns, as described by your TA.

Claims:Craft at least two Claims directly addressing the Selected Question and your Hypothesis. For this lab, at least one claim must relate to precision, and one claim must relate to accuracy. Claims are to be the result of independent student thought and interpretation of the data, though they may be guided by the whole class discussion. Copying another student’s wording (whether exactly or with minor paraphrasing) will result in zero points for both students.

Be sure to cite evidence in support of your claims!

Post-lab question:

1. Choose: Follow significant figure conventions to perform the following calculations, OR determine the implied uncertainty for each value and determine the uncertainty using Eq. [2].

(Correct answers will include units.)

a) 11.645 g / 10.0 mL

b) 0.032 g / 1.7 mL

c) 2 g / 5.32 mL

d) 17 g / (2.45 cm x 2.51 cm x 2.55 cm)

Reflection:The Reflection explains what you learned about chemistry as a result of this lab. Looking back on this lab, how would you define a “good measurement”? Did this lab alter how you think about measurement, uncertainty, or significant figures? How do you think significant figure conventions might relate to explicit uncertainty estimates? What are the meanings of the calculated uncertainty of density (from Eq. [2]) and the standard deviation: do they measure the same thing; how are they different; was one consistently larger than the other in your data; did one do a better job predicting agreement between the experimental and actual values?


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Skill-Building Inter-session

Skill-Building Inter-session #1

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Overview: Students will be trained on lab equipment and procedures that will be important for multiple future laboratory experiments. Students will be responsible for recording detailed information on available tools and correct operation and cleaning of each item into their lab notebooks. This process is very similar to what students may expect to encounter when they join a research lab; a more experienced mentor will provide guidance, and they will be expected to take notes so that they may perform independently in the future.

Topics that will be covered: Checking and cleaning glassware before use Dispensing reagents Reagent overages Avoiding cross-contamination Use of the hood Waste management Separation procedures: decantation, filtration, and evaporation Drying to constant mass Rinsing procedures: when, why, and how Pipets Burets solubility testing for Exp. 3 pH testing with red & blue litmus paper for Exp. 3 conductivity tests for Exp. 3 chemical reactivity testing using microwell plates for Exp. 3 flame tests for Exp. 3 Using Excel effectively and efficiently

o Entering formulas and automating calculationso Adjusting significant figures in Excelo Showing units when using Excelo Flow charts on Excel (useful but not required for Exp. 3)

** If you have access to a laptop, please bring it to lab. It will be helpful in learning Excel. (Remember that @One, in the basement of the Knowledge Center, offers a free 2-day check-out of laptop computers.)

Skill Building Inter-session

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In-Lab Requirements

At a minimum:Show headings in your lab notebook for each topic covered by your TA. Write notes on anything you are not certain you will remember in future weeks. Any handouts that are useful may be stapled or taped into your lab notebook.

Also complete the following:Chemical handling protocols are essential to any chemistry lab. Write a sentence suggesting a reason for each of the following UNR chemical handling protocols; why do you think this procedure has been adopted; what purpose does it serve?

1) If you accidentally pour out too much of a chemical from the stockroom supply bottle, you may NEVER return the excess back to the bottle (though you may certainly ask around to see if any other groups could use your excess, or even have a community “excess storage beaker,” which would need to be transferred to the appropriate waste container at the end of lab if not used).

2) At the end of lab, liquid waste should be dumped into the appropriate waste container. The glassware that had been holding the waste should then be rinsed with ~2 trials of small volumes of water (filling < 10% of the glassware, NOT completely filling the glassware each time), with “swirling” to rinse and wet the sides of the glassware. Each rinse should be dumped into the waste container. Further washing may take place, if needed.

Inter-Session GradingStudents will receive a grade out of 10 points for participation, productivity, and professionalism during the lab.

Skill Building Inter-session

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Experiment 2: Reactions with Copper


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Overview: **Note that Experiment 2 is a 2-week lab.** In Experiment 2, students will perform reactions on copper and copper-containing products. Reactions proceed in order (see Pre-lab question #2 for the sequence of each step), where the copper-containing product of each step is one of the reactants for the next step. Reactions include reacting bare copper wire with concentrated nitric acid (in the hood!) to form a copper (II) nitrate salt, a precipitation reaction with a reagent of your choice--selected from the chemical list on page 29--to create a copper (II) hydroxide precipitate, a dehydration reaction of the copper (II) hydroxide, an acid-base reaction of copper (II) oxide with hydrochloric acid, and a redox (or oxidation-reduction) reaction of a copper (II) salt with zinc.

Students will consider and calculate what amounts of reactants to use for each step. This experiment will require the use of stoichiometric calculations, an in-depth understanding of reactions, including use of the solubility table to predict products, and will provide students with practical understanding of limiting reactants.

Background - Reaction Types and Generic Formulas: Several types of chemical reactions will be explored in this experiment. The following generic chemical equations represent four out of the five reactions that will be explored (the reaction of bare copper wire with nitric acid may be considered a complex redox reaction):

Dehydration reaction: A B + H2O (Note that alternate generic equations exist, e.g., A + B C + H2O)Essentially, water is being removed from the reactant(s). Thus, the name “Dehydration reaction” is particularly logical. (This reaction type is in addition to those listed in your text.)

Oxidation-reduction (or redox) reaction: AX + B BX + A (Note that numerous alternate generic equations exist, e.g., AX + Y AY + X)Note that one “giveaway” characteristic denoting many redox reactions is when an elemental species converts into a member of an ionic compound (such as B going from elemental B as a reactant to a member of the compound BX as a product in the first redox generic equation) or vice versa (A going from a member of compound AX as a reactant into pure elemental form, A, as a product). Not all redox equations have this feature; however, all equations that do have this feature are redox; it is thus a sufficient but not necessary characteristic of redox reactions.

Precipitation reaction: AX (aq) + BY (aq) AY (s) + BX (aq) Note that A and B are listed first in the all compounds shown and are thus metals or polyatomic cations, while X and Y must be nonmetal or polyatomic anions; there would never be a product AB or XY. Finally, one of the products is required to be a solid which precipitates out of solution.


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Acid-base (or neutralization) reaction: HX + AOH AX + H2ONote that the species HX is an acid, and the species AOH is a base. The products of the neutralization reaction are the salt, AX, and water.

For this experiment, we are using a less traditional base with the chemical formula AO, where O is oxygen. This base would necessitate a slightly different generic chemical equation:

2 HX + AO AX2 + H2O

Important information about all generic chemical equations:

While no subscripts or coefficients are shown in the generic equations, they are frequently necessary in real equations. Two examples are provided below.

The generic reactants AX + BY may be Ca(NO3)2 + (NH4)2CO3, in which case the products AY + BX would necessarily be CaCO3 + 2 NH4NO3.

Neutralization may occur between sulfuric acid, H2SO4 , and sodium hydroxide, NaOH. The generic equation HX + AOH AX + H2O would then become H2SO4 + NaOH Na2SO4 + 2 H2O.

Subscripts and coefficients are necessary to provide accurate chemical formulas for each compound and to balance the chemical equation, respectively.

The benefit of the generic equation is its simplicity for grouping types of reactions and for allowing a clear representation of the essential process occurring in the reaction.

Background - Percent Yield: Stoichiometry allows the use of balanced chemical equations to calculate the theoretical yield of a reaction. This calculated, theoretical yield is an ideal value; chemical reactions in the laboratory rarely produce the exact amount of product predicted by the stoichiometric calculations. In real life, some fraction of the chemical is “lost” for a wide range of possible reasons, including: the product adheres to the surface of the glassware and can not be collected with 100% efficiency; the limiting reactant engaged in a side reaction with a contaminant that was present on the glassware; or the reactant(s) simply did not react (perhaps due to the forward reaction being incomplete or the presence of a reverse/equilibrium-style reaction).

The amount of product that is predicted by calculations is the theoretical yield of that product and is typically reported in grams. The amount of product that is actually collected in a laboratory (again, typically in grams) is the actual yield. The percent yield is the ratio of these two yields, and is reported as a percentage by convention.

Percent yield=( Actual yieldTheoretical yield )×100% [ 1 ]

Chemists often work long and hard to determine the experimental conditions (temperature, pressure, ratio of reactants, effective catalysts, etc.) that will produce optimal percent yields. Note also that all reactions do not simply “proceed to completion” instantaneously, just because


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the correct ratios of reactants are mixed; rates of reaction may cause significant slowdowns. Percent yield is a topic that is very highly used in current research and publications.

Background - Limiting Reactants: Imagine that you are a chemist working to develop a new material with one reactant that is very costly. You would not want to waste any of that reactant. Let’s look at an example.

In the chemical equation A + 2 B + 3 C D, we need twice as many moles of substance B as we do of substance A, and we need three times as many moles of substance C as we do of substance A. It is critical that all substances be compared in units of moles.

Suppose substance B is our expensive reactant, and we are starting with 0.15 moles of substance B. We must choose how much of substances A and C to react with substance B, and we must predict how much of substance D we hope to get.

If we count on ideal reaction conditions, we can use stoichiometry to calculate each of these values.

A + 2 B + 3 C D

0.15moles B× 1mole A2molesB

=0.075moles A

0.15moles B× 3molesC2moles B

=0.23molesC

0.15moles B× 1mo≤D2molesB

=0.075moles D

We can see that if we performed the reaction with 0.15 moles of B, 0.075 moles of A, and only 0.30 moles of C, we would limit our ability to produce D and would expect a significant reduction of product, effectively wasting our most expensive reactant, substance B. (Pre-lab question 1 explores this further.)

In practice, we would finish each calculation above by converting moles of each substance into grams of that substance, by multiplying each value above by the corresponding molar mass of each substance.

e.g., If substance C is sodium hydroxide, we would perform the calculation:

0.45moles NaOH × 40.00 g NaOH1mole NaOH

=18grams NaOH


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Be alert to the following common mistake! Molar mass conversions always use exactly 1 mole of substance, no matter what the coefficient in the balanced equation may be. (Also, as an exact number, 1 mole has infinite significant figures.) That is because each conversion is “legal” only if the numerator is equivalent to the denominator--that is an important concept and extends to all conversions, not just molar mass. When the numerator is equivalent to the

denominator, the conversion is essentially

the same as multiplying by 1, always a legal operation.

Limiting reactant problems often present students with the question of determining which reactant will most limit the amount of product, given specific starting masses of each reactant. Students balance a chemical equation for the reaction (this is always the first, necessary step), convert grams of reactant to moles of reactant, convert moles of reactant to moles of product, and compare to see which calculation results in less product. Detailed examples of this process are shown in the course text.

In this experiment, the limiting reactant situation follows more along the lines of the scenario described previously, with one relatively expensive reactant: copper. Thus, the problem changes from one in which the limiting reactant identity is determined from known starting masses of each reactant to a problem in which the starting masses are determined from knowing the identity of the limiting reactant. This is a closer representation to the thinking process used in industry and research work. The lab description will not state how much of each substance to react; rather, students will perform stoichiometric calculations based on a selected starting mass of copper and a student-selected paradigm that guarantees that the copper (or copper-containing reactant species) will be the limiting reactant.

Goals:

Attain proficiency performing stoichiometric calculations.

Determine the percent yield of the final product. Evaluate class data for trends that maximize the percent yield.

“Along-the-way” Goals:

Convert between chemical names and chemical formulas; write and balance chemical equations.

Utilize solubility rules to determine phase (aqueous or solid) of various compounds.

Recognize classes or types of reactions.

Chemical List:Note: These chemicals are available to you, but not all are necessary for this lab. Maximum amounts may apply.

barium hydroxide


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calcium carbonatecopper wire (maximum of 8.0 grams total, per class)hydrochloric acidisopropyl alcoholconcentrated nitric acidpotassium hydroxidesodium hydroxidesulfuric acidgranulated zinc

Pre-Lab Requirements Due the FIRST Week of this Experiment

Pre-lab Exercises:

1. In order to determine the actual yield of a solid product, the solid product must be collected. This often requires a separation procedure, and it may also be helpful to perform rinses on your product to be sure that you do not have any other substances adhering to the surfaces of your glassware or attached to your product. Review your notes from the Skill Building Inter-session, as needed.

Suppose that a reaction of interest produces a lacy precipitate that is suspended throughout a ~100 mL aqueous reaction system. Which method (decantation, filtration, or evaporation) is best suited to the task of separating out the lacy precipitate product? Explain your answer with a couple of solid, logical sentences.

2. Complete and balance each equation shown below, including phases where indicated; write each Step # and the completed equation into your lab notebook.

These steps represent the chemical reactions you will perform in the laboratory; note that the copper-containing product of each step is a reactant for the following step.

Step # Chemical equation

(1) ¿¿+26HNO3⟶¿Cu ¿¿

Hint: Begin by considering the element with only one unknown coefficient (H). This allows you to solve for the coefficient for nitrous acid, HNO2. Then, you may assign variable x as the coefficient for copper (II) nitrate and variable y as the coefficient for nitric oxide, NO. Write an algebraic equation for oxygen ( Σ Oxygenreactants = Σ Oxygenproducts) in terms of known coefficients and x and y. Write another algebraic equation for nitrogen. This yields two equations and two unknowns; solve via your favorite method…or, just guess and check. Finally, solve for copper.

(2) ¿¿


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Note that you must determine and show phases for Steps 2, 3, 4, and 5. Also, Step 2 must produce copper (II) hydroxide, and the second reactant must be an aqueous solution. It may be helpful to refer to Table 4.1 from the text, replicated below.

(3)

Cu(OH )2( )∆→

¿+¿H 2O(g)

Note that this dehydration reaction involves the application of heat, typically shown as a Delta symbol (∆ ¿ above the arrow in the chemical equation.

(4) ¿¿ ( s )+¿HCl(aq)⟶¿+¿( )

Recall from “Background - Reaction Types and Generic Formulas” that metal oxides are bases. Copper (II) oxide will react with a strong acid such as hydrochloric acid to form a copper (II) salt and water.

(5) ¿¿ ()+¿Zn(s)⟶¿+¿( )

Your copper salt formed in Step 4 will react with zinc to produce pure copper metal and a zinc salt.

3. This lab has been called the “Copper Cycle” lab. In what way do Steps 1 through 5 comprise a complete cycle? The percent yield for the full cycle or series of reactions is the actual, experimental mass of copper product resulting from Step 5 divided by the mass of copper product expected or calculated on paper. How does the theoretical yield or mass of copper product expected from Step 5 compare with the mass of copper wire present at the beginning of Step 1, given ideal reactant ratios throughout?

4. For each of the following reaction types, write the generic equation for that reaction type and identify which chemical equation from question 2 represents an example of that reaction type.

example) Acid-base reaction: HX + AOH AX + H2O or 2 HX + AO AX2 + H2O; Step 4.

a) Dehydration reaction


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b) Oxidation reduction (or redox) reaction

c) Precipitation reaction

Honors. View the MSDS document, “Concentrated nitric acid MSDS,” available on Webcampus, and report at least two hazards as well as what treatment is required (e.g., the appropriate first aid measures) if it comes into contact with your skin.

ALSO: Since Experiment #2 is a two-week lab, only the Original Beginning Question section is required (to be written into your lab notebook) for the First Week Pre-Lab assignment; the Selected Question, Procedure, and Safety Notes will be developed jointly in lab.

Original Beginning Question:Consider the goals of the laboratory experiment and the suggested Beginning Question Seeds below (next page). Write an Original Beginning Question of your own that you would be interested in investigating. (Beginning Question recommendations for all labs are provided on page 7.)

Beginning Question Seeds:

Does the starting amount of copper wire (ranging from 0.1 g - 1.0 g) affect the percent yield?

Does the selection of sodium hydroxide vs. potassium hydroxide as the second reactant in Step 2 affect the percent yield?

Does the amount of sodium hydroxide added during Step 2 affect the percent yield?

How consistent are percent yield values across the class, when everyone follows the same procedure exactly?


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In-Lab Requirements for the FIRST Week of this Experiment

NOTE: Be prepared to work together as a class and stay as long as it takes to develop a skeleton procedure for the class to use on the second week of lab. No experimentation will be performed today. An important and relatively time-consuming task that must be completed by the end of lab today is that each individual student must calculate the values (masses and/or volumes) of substances to use for his or her group’s assigned work. Once all group members have completed their independent calculations, they may compare values and verify that values are consistent and correct with one another and/or the class as a whole.

Your TA will write the molarity of stock solutions being provided for next week’s lab work on the board.

Pre-Lab Discussion:For the first ~5 minutes of class, discuss your pre-lab work with your table group. Share your most brilliant, innovative ideas, ask for help on anything that has been problematic for you, and reach a consensus on one or two Original Beginning Questions to write on the board as a submission for the whole class discussion.

Spend about 10-20 minutes talking as a whole class. Consider and agree on:

Which Original Beginning Question will the class investigate? How much copper wire will each group start with? How will we isolate the product of Step 5 from all other substances? What might aid us

to dry it efficiently? Additionally, are there any risks in over-drying or overheating it? (Is this an appropriate time to use the “drying to constant mass” technique?)


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For the excess reactant of each step, how much excess should be provided? (Please consider using ratios to describe the amount of excess reactant. Which excess reactant, if any, will have its ratios vary across the class? --if so, what will the ratios be?)

What are some potential hypotheses we might make about this question, and what data must be collected to prove or disprove our hypotheses?

Walk your way through the sequence of reactions. What details should your procedure contain to describe how each of the reactions will be conducted? What glassware will be used? Will you simply keep adding more and more chemicals to the same beaker or flask you used for Step 1? (What side reactions might you expect to occur if you did so?)

What else do you need to specify for a complete “skeleton procedure” with only the amounts left to be filled in, pending the completion of each set of calculations?

How will we divide labor? What are the safety concerns?

Each student in the class will need to provide a notebook heading and entry for the Selected Question chosen by the class and will individually develop and write a Hypothesis that addresses the Selected Question. Students will also record the Procedure and Safety Notes developed by the class; each student will work independently to complete the Calculation Guide shown on the next page to “flesh out” the class Procedure with specific values.First Week Calculation Guide (This is a subsection of your Procedure.)

Historically, many chemistry students struggle with stoichiometric calculations. This in-lab opportunity to perform a sequence of calculations is intended to provide meaningful practice that connects to concrete lab activities next week. Calculations must be performed individually; comparisons and collaboration are permitted only after completion of your calculations.

1. State your starting amount of copper wire, to 2 significant figures. (This will be give you a bit of tolerance next week when attempting to cut the wire.)

2. Show (and label) a calculation for the stoichiometrically equivalent amount of concentrated nitric acid--the amount that should theoretically, exactly react with the starting mass of copper wire. (Be sure that you specify the molarity of concentrated nitric acid being used and/or show it clearly in your calculation.)

Show (and label) a second calculation (using the results of the first, if desired) for the actual amount of {specified molarity} nitric acid to be used.

3. Calculate the theoretical yield (in grams) of the aqueous copper (II) nitrate product of Step 1.

4. Show (and label) a calculation for the stoichiometrically equivalent amount of {specify identity and molarity of reactant being used for Step 2} that will exactly react with the theoretical yield amount of copper (II) nitrate calculated previously.

Show (and label) a second calculation (using the results of the first, if desired) for the actual amount of {specified identity and molarity of reactant being used for Step 2} to be used.

5. Calculate the theoretical yield (in grams) of solid copper (II) hydroxide product of Step 2.


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6. Calculate the theoretical yield (in grams) of solid copper (II) oxide product of Step 3.

7. Show (and label) a calculation for the stoichiometrically equivalent amount of hydrochloric acid that will exactly react with the theoretical yield amount of copper (II) oxide calculated previously. (Be sure that you specify the molarity of hydrochloric acid being used and/or show it clearly in your calculation.)

Show (and label) a second calculation (using the results of the first, if desired) for the actual amount of {specified molarity} hydrochloric acid to be used.

8. Calculate the theoretical yield (in grams) of the aqueous copper (II) chloride product of Step 4.

9. Show (and label) a calculation for the stoichiometrically equivalent amount of zinc that will exactly react with the theoretical yield amount of copper (II) chloride calculated previously.

Show (and label) a second calculation (using the results of the first, if desired) for the actual amount of zinc to be used.

10. Compare your values with other group members and check for consistency; write the actual values to be used into your Procedure.

If possible, finish this class by planning your data tables for next week. (As always, give space for units and remember to use pen for all original measurements.) The lab portion of next week will be very lengthy, and anything you can do now to save time next week will pay off.

Next week, as a class, you will have a brief refresh chat, make sure everyone is ready to go, and jump into your experimental procedure.


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Pre-Lab Requirements Due the SECOND Week of this Experiment

None. You get this week off. If you struggled with the stoichiometric calculations, however, it is strongly recommended that you put some time into practice, stop by the Chem Help Center, etc.


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In-Lab Requirements for the SECOND Week of this Experiment

Pre-Lab Whole Class Mini-Discussion:This is your brief refresh chat. Keep it brief because the lab takes a long time; however, be sure to review the safety notes for this experiment. Take enough time to be sure everyone is fully prepared and confident about what to do. If you’re the one who is not sure, then you owe it to the class to ask questions so you don’t foul up everyone’s data. Write the model data table on the board and get the official version into a computer spreadsheet. Then, as always, each individual student will need to model his or her own data collection table after the one on the board (under a section titled Data, Observations & Calculations).

Note that for this experiment, your Data, Observations & Calculations section must have an area for recording qualitative observations about the experiments. Describe how each copper-containing species looks; also include solution color change, description of gas evolution, and notes where applicable such as “slight powdery black residue remains on the beaker surface,” etc.

Perform Your Experiment and record your original data and observations into your data table, in ink. Perform any required calculations, remembering to show formulas being used. Carefully and systematically follow significant figure conventions.

IF PART OF YOUR PROCEDURE ISN’T WORKING AND YOU WISH TO MODIFY IT, the entire class must stop working and agree how to handle the situation. (e.g., If your copper wire doesn’t “dissolve” completely in the concentrated nitric acid, is it kosher to set up a hot plate in the hood and apply gentle heat to coax the reaction to completion? …or do you want to establish a protocol for finding the mass of the rinsed, unreacted sliver of copper, perhaps including a trip to the stockroom to obtain tweezers? …or devise something else entirely?)


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Calculate percent yield while in lab. Show a complete calculation in your lab notebook. Contribute your original data and calculated percent yield values to the whole class data tables. The class data as a whole may be copied onto zip drives and/or emailed to students, and they may be printed from the Excel spreadsheet and attached to the lab notebook.

Clean up; wipe down your lab bench area, etc. If you finish your work early, find other groups to help.

Post-Lab Discussion (takes place in lab):Time may or may not permit a whole class discussion this week. If time allows, students who wish may choose to stay for a class discussion and/or lab completion round table. As usual, discussions will be student-directed and may address anything that is relevant or helpful.

As always:Your lab safety goggles must stay on until everyone in the lab has put away all glassware. Remember to lock your lab drawer before leaving.Post-Lab Requirements (MAXIMUM length in lab notebook: 2 pages; failure to be concise will result in loss of points.)

Results:Provide data tables showing initial and final mass of copper and percent yield (class-wide tables are okay and do not count against the 2 page limit). Check that tables show units and the correct number of significant figures. Include any additional information pertinent for your claim and reflection discussion (e.g., value of other reactant of interest, graph showing trends in data or lack thereof, etc.).

Claim:Craft at least one Claim directly addressing the Selected Question and your Hypothesis. As always, claims are to be the result of independent student thought and interpretation of the data, though they may be guided by the whole class discussion. Copying another student’s wording (whether exactly or with minor paraphrasing) will result in zero points for both students.


Post-lab questions:

1. State the average and standard deviation of class data for percent yield. (Show work. Use correct units and by convention, state the standard deviation to 1 significant figure and the average to the place value of the standard deviation. Good examples include: 92.3% ± 0.5% and 87% ± 4%; note the consistency in the place value of the final significant figure of the average and the lone place value of the standard deviation. Incorrect examples include: 89.1% ± 6% and 77% ± 0.08%.)


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2. What is the theoretical yield (in grams) of copper in Step 5, given a hypothetical experiment in which the copper-containing reactant is the limiting reactant for every step, and given a starting mass of copper wire in Step 1 of 0.875 g? Show calculations or explain your reasoning.

3. What is the percent yield if a reaction between 0.743 g copper (II) oxide and 12.0 mL of 1.00 M hydrochloric acid is determined to produce 0.547 g of copper (II) chloride experimentally? Show calculations.

Honors. How well could you isolate each copper-containing product in Steps 1 through 5 (or would you have been able to isolate, had you tried)? What were some of the challenges in the isolations that you did perform?

Reflection:The Reflection explains what you learned about chemistry as a result of this lab. Looking back on this lab, how did the percent yield values compare to your expectations? What trends did you observe on percent yield; what other trends might you now suspect? What is the importance of balancing chemical equations before performing stoichiometric calculations? Did your ideas about stoichiometry, limiting reactants, or percent yield change as you worked through this lab?

Experiment 3: Qualitative Analysis


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Overview: **Note that Experiment 3 is a 2-week lab.** In Experiment 3, students will collect qualitative (non-numerical) data about the physical and chemical properties of selected substances, based on careful observations and use of systematic procedures. Students will use the first week to build testing skills and will work collaboratively to collect information about known substances, including appearance (color, texture, graininess, etc.) of the substance, color of the substance during a flame test, solubility (in water, acid, and/or base), pH (of an aqueous substance), conductivity (of an aqueous substance), and reactivity with provided 0.10 M solutions of barium chloride, calcium nitrate, iron (III) nitrate, lead (II) nitrate, silver nitrate, potassium nitrate, sodium bromide, sodium carbonate, sodium chloride, sodium hydroxide, sodium iodide, sodium nitrate, potassium oxalate, sodium phosphate, and sodium sulfate. (These solutions are intended for use with microwell plates.) The principal task of the first week will be to gather the data that will permit development of a flowchart for testing an unknown chemical and determining its identity (this is the pre-lab assignment for the second week of this experiment).

Students will see the same substances in the second week; however, the substances will now be in the guise of numbered “Unknowns” (numbers will range from 1 - 32 with possible repeats as well as possible voids; for example, there might be six vials of boric acid and no calcium carbonate). For the second week, individual students will start with an unknown solid sample, attempt to follow their flowchart plan or logical outline for testing the sample, and make an unambiguous identification of the unknown from comparison with the data they have collected previously. Students must also work within the constraint of a small sample size of only 1.0 gram of each unknown sample; thus, students must plan how to use their samples efficiently.


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Background: Qualitative Analysis is a field of chemistry concerned with identifying the substance(s) present in a sample and seeks to answer the question, “What is it?” (Qualitative Analysis is commonly contrasted with Quantitative Analysis, which seeks to determine, “How much is present?” and which we will pursue in Experiments 4 and 6.)

Qualitative Analysis may give a significant amount of information. For example, solubility in water is commonly classified as “soluble,” “sparingly soluble,” or “insoluble;” however, you might find some substances that you wish to classify as “exceedingly soluble,” or you might compare solubility in cold water to solubility in hot water. It is recommended that students observe physical appearance (of the original solid substance and any reaction product) with care and an eye for detail; use the most precise language possible to describe colors, etc. Your ability to make well-organized, careful, consistent, and systematic observations will be a key to your success in this experiment.

Qualitative Analysis also typically involves identification based on a small sample size; thus, procedures need to be highly efficient with respect to sample usage. An ideal identification flowchart will lead to unambiguous identification in a small number of steps, with some sample left over for verification.

Qualitative Analysis procedures may support identification of periodic trends or other categorization techniques; these, in turn, may be used to gain efficiency in classifying the sample. Do some tests necessarily precede others? (If so, why?) Which test(s) are best for differentiating among substances? (Perform these first!) How can you make a determination using the fewest possible tests yet still guarantee an unambiguous identification? Will some tests make others unnecessary? Are there some results that point to the need for a follow-up test?


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Goals:

Design an efficient procedure (and communicate the procedure in the form of a flowchart or logical outline) for determining which substance among 12 possible choices is present in a sample.

Test your flowchart or outline: accurately identify your unknown sample.

“Along-the-way” Goals:

Compare evidence collected on an unknown sample to evidence collected the previous week on a set of known samples.

Write complete and net ionic equations for precipitation reactions, including phases.

Obtain and evaluate a comprehensive body of data for essential, distinguishing tests and characteristics.

Build skills in solubility testing, conductivity testing, conducting a flame test, and using a microwell plate to test reactivity.


Bunsen burners (and lighters)conductivity test kits (Note: LED lights are directional; also, they must be paired with resistors.)microwell platespH paper (red and blue)pipets

Chemical ListThese chemicals are all available to you, but you may find that all are not required for accurate identification.

“First Week Knowns and Second Week Unknowns” (provided as solids each week):

boric acid sodium carbonate calcium carbonate sodium chloride calcium nitrate sodium hydroxide potassium iodide sodium iodide silver nitrate sodium phosphate sodium acetate sodium sulfate

1.0 M hydrochloric acid solution is available for testing solubility in acid (and for cleaning nichrome wire for flame test)

1.0 M sodium hydroxide solution is available for testing solubility in base

0.10 M solutions of the following are available for testing with the microwell plates:


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barium chloride sodium bromide copper (II) nitrate sodium carbonateiron (III) nitrate sodium chloride lead (II) nitrate sodium hydroxidesilver nitrate sodium iodide

potassium oxalate

Pre-Lab Requirements Due the FIRST Week of this Experiment

Pre-lab Exercises:

1. Suppose you have been assigned to test reactivity of calcium nitrate. You make a solution of calcium nitrate to test in your microwell plate. Review the list of chemicals supplied in 0.10 M solutions for testing chemical reactivity: barium chloride, copper (II) nitrate, iron (III) nitrate, lead (II) nitrate, silver nitrate, sodium bromide, sodium carbonate, sodium chloride, sodium hydroxide, sodium iodide, and potassium oxalate. Would you test your calcium nitrate solution with all of the above solutions or just some (if so, which ones)? Briefly explain your reasoning.

2. For the reactions possible in question 1, choose two that you expect to form a precipitate. Write balanced equations for each of the two reactions you selected. Include phases. Refer to your text and/or Table 4.1 shown below for solubility rules, as needed.


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3. Write net ionic equations for each of the reactions you selected in question 3. Include phases.

Honors. Suppose sucrose (table sugar) were present on the list of “First Week Knowns and Second Week Unknowns”. Given that sucrose is a covalent compound, you would expect the addition of sucrose to distilled water to form a solution that would not conduct electricity. What other compound on the list “First Week Knowns and Second Week Unknowns” (p. 40) might you expect to have low (or no) conductivity when placed in distilled water? (There is more than one correct response.) Explain your reasoning for the substance you select.

Original Beginning Question:

Consider the goals of the laboratory experiment and the Beginning Question Seeds below. Write an Original Beginning Question of your own that you would be interested in investigating, or choose one of the seeds from below and elaborate on it. (Beginning Question recommendations for all labs are provided on page 7.)

Beginning Question Seeds: Which physical tests or 0.10 M solutions give the most differentiated results or reaction

products {and how can we use this information to minimize the number of steps required in our flowchart or outline}?

What tests (physical and/or chemical) demonstrate similarities in groups of the periodic table?

Which physical or chemical tests are least informative and least useful in constructing a flowchart or logical outline?

Which physical or chemical tests provide the most clear and consistent data across the class?

Which known/unknown substances are most challenging to identify?

Selected Question:

Contribute to the online discussion of Beginning Questions. Write the Selected Question in your lab notebook. It is each student’s responsibility to follow the timetable given by your TA for this discussion.


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Hypothesis: Write your educated guess for what your findings will be in answer to the Selected Question. State a reasonable explanation and/or give some background in support of your hypothesis.

Procedure:

Review your notes from the Skill Building Inter-session as needed for reminders about solubility testing, conductivity testing, pH testing with blue and red litmus paper, chemical reactivity testing with microwell plates, and flame tests. Describe (as though to someone who did not have access to your Skill Building notes) a brief, first week procedure for complete testing of the 12 known samples. Include approximate amounts of known sample and other reagents that will be used for each test. Consider ways to standardize the processes; also consider which tests may be skipped for each of the 12 known samples.

Safety Notes:

Use common sense and/or prior experience to suggest at least one safety note specific to this lab. (Hint: you may wish to consider the inherent hazards of chemicals being used.)

In-Lab Requirements for the FIRST Week of this Experiment

NOTE: Be prepared to work together as a class and stay as long as it takes to get all the data you will need to (individually, at home, or collectively during post-lab discussion) develop a complete flowchart/outline that you will use on the second week of lab.

Pre-Lab Discussion:For the first 2-3 minutes of class, discuss your pre-lab work with your table group. Share your most brilliant, innovative ideas and ask for help on anything that has been problematic for you.

Spend about 5-10 minutes talking as a whole class. Consider and agree on: Summarize the procedure to be used; determine how to standardize it across the class as

much as possible. How will the data be organized? (“Free” sample data tables are provided to use, modify, or

ignore in favor of creating your own from scratch.) What are some of the hypotheses we are making, and what data must be collected to prove

or disprove our hypotheses? How will we divide labor? What are the collective safety concerns?

Students may modify or completely re-write their Hypotheses, Procedures, and Safety Notes as a result of this discussion. Standardized methods are strongly encouraged.


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Conduct the first week experiment and carefully note your own First Week Data and Observations in your lab notebook. Report all data for class evaluation. Be sure to stay and get the complete set of whole class data. As always, clean up; wipe down your lab bench area, etc. when your work is complete. If you finish early, find other groups to help.

Post-Lab Discussion (takes place in lab during first week):Determine the chemical and physical properties that can uniquely characterize each substance (find the subset of distinct properties that is necessary and sufficient to identify each substance). Your TA will introduce some ideas on how to construct an efficient flowchart or logical outline to be used the second week of this experiment (this is the pre-lab assignment for next week).

Pre-Lab Requirements Due the SECOND Week of this Experiment

Complete a flowchart or logical outline to guide your second week procedure. This is a graphical representation of what you will do with your unknown sample, in sequence, and it needs to lead to all 12 possible outcomes. (Any student with a missing or blatantly incomplete flowchart will sit out of the experiment until the flowchart is completed.) This is not a trivial assignment; it will take time to complete. Be sure yours is neatly organized and clearly legible, preferably typed. Print two copies; one to submit to your TA and the other to tape or staple into your lab notebook.

Example flowchart to distinguish between Cu(OH)2, CuCl2, and Cu(NO3)2:


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Example logical outline to distinguish between Cu(OH)2, CuCl2, and Cu(NO3)2:

In-Lab Requirements for the SECOND Week of this Experiment

Pre-Lab Small Group Mini-Discussion:Compare flowcharts with other students. Any differences will be potential discussion points for your reflection.

Follow your flowchart/outline, record precise data and observations, and identify your unknownObtain 1.0 g of an unknown and be sure to record your unknown # in a prominent location. Carefully and systematically record your Second Week Data and Observations. (There are no

I. Is the original, dry sample blue?A. Yes It is copper. Go to II.B. No It is not copper. Go to IV. (IV = to be determined for other species…)

II. Dissolve approx. popcorn kernel size of sample in ~3 mL distilled water. What is the pH of the resulting solution? (Dab a clean stir rod into the solution, then tap it onto red litmus paper.)

A. Red litmus paper turns blue This indicates presence of a base & COMPOUND IS COPPER (II) HYDROXIDE.B. Red litmus paper does not change color This indicates the absence of a base. Go to III.

III. Mix a few drops of the solution prepared in II above with a few drops of silver nitrate. Does a precipitate (commonly abbreviated “ppt”) form?

A. Yes (white ppt) Precipitate is silver chloride & COMPOUND IS COPPER (II) CHLORIDE.B. No By process of elimination, the COMPOUND MUST BE COPPER (II) NITRATE.

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calculations for this lab.) Work individually; however, consultations (especially if experts were assigned during the first week) are recommended. You may choose to include additional, confirmation tests once your original flowchart or logical outline procedure has provided an identity. If you stray from your flowchart, record careful notes on procedures performed as well as resultant data and observations.

Report your unknown number and suspected identification to the board or class-wide computer-based data table. Your TA will then provide the actual identification for comparison purposes.

Honors. Identify two unknowns (rather than one) during this experiment.

Clean up; wipe down your lab bench area, etc. If you finish your work early, find other groups to help.

As always:Your lab safety goggles must stay on until everyone in the lab has put away all glassware. Please double check that your lab drawer is locked.

Post-Lab Discussion (takes place in lab):How did it go? Across the class, were there different identifications made for the same unknown (same actual substance, just numbered differently)? What were the difficulties for this experiment? Were there any surprises? How do your findings relate back to the selected question? As usual, discussions will be student-directed and may address the above questions and/or anything that is relevant or helpful.


Results:Provide data (tables/other) demonstrating pertinent information for your claim and reflection discussion. Be sure to state your unknown number, your identification of it (name and chemical formula), and its actual name and chemical formula. Honors: do this for each unknown that you tested.

Claim:


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Craft at least one Claim directly addressing the Selected Question and your Hypothesis. As always, claims are to be the result of independent student thought and interpretation of the data, though they may be guided by the whole class discussion. Copying another student’s wording (whether exactly or with minor paraphrasing) will result in zero points for both students.


Post-lab questions:

1. Explain what information collected during the first week of this experiment guided you most strongly in developing your procedure (flowchart or logical outline) for the second week. Also comment on any tests you elected not to use at all in your procedure, and why.

2. Write balanced equations for any two chemical reactions that you observed in this lab (first and second week, collectively). For each, write the full molecular equation with phases and show the net ionic equation with phases. (You may show the complete ionic equation, as well, if it is helpful to you to do so, but it is not necessary. Consult section 4.6 of your text as needed for details/reminders on writing net ionic equations.) Also note that your selected pair of chemical reactions should (statistically speaking) be different from other students’ pairs.

Honors. How do you predict iron (III) chloride could be identified, uniquely from the other substances, if it were added to the Second Week Unknowns? Cite evidence from the First Week testing results to support your prediction.

Reflection:The Reflection explains what you learned about chemistry as a result of this lab.

Explain in a few sentences how you would assess your procedure (your own flowchart or outline). (e.g., Was it effective and efficient? Did you find any “holes” or room for improvement as you used it? What comparisons were you able to make with other students’ versions?)

How does this lab fit into the “big picture,” or how will it inform your ideas for future experiments? (e.g., Did you observe any broader trends? Did you observe any periodic trends or consistencies among groups of elements? What substances were nonreactive? What do you think makes them nonreactive, based on your knowledge of chemistry?) As always, make connections between this lab and the rest of what you know about chemistry.

Experiment 4: Titrations of Acids

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Overview: In Experiment 4, students will perform titrations of acidic solutions with known acid identities but unknown molarities. Students will first perform a standardization of stock sodium hydroxide solution by titrating against the primary standard, solid potassium hydrogen phthalate, or KHP. (See Background for additional details on this.) Once the concentration of the sodium hydroxide solution is known, students may measure what volume of sodium hydroxide solution is required to completely and exactly neutralize a given sample of acid. This measurement is performed by titrating, using a couple drops of phenolphthalein indicator


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solution to demonstrate the endpoint of the reaction. If the volumes of both the acid solution and the sodium hydroxide are measured precisely and accurately, and the molarity of the sodium hydroxide solution has been determined through standardization, students may readily determine the molarity of the sample acid solution.

Finally, students will use pH meters to obtain the initial pH of each acid solution. They will use this data in conjunction with titration data to further evaluate their acids.

Background - Titration: Titration is a common yet powerful method used in analytical chemistry and may be described as the process of gradually releasing one substance, the titrant, shown to the right at Substance A, which is the stock sodium hydroxide solution in this lab, into a flask containing a precisely measured volume of a second substance, the analyte, shown on the right as Substance B. The analyte for this experiment is an acid sample with known identity but unknown molarity.

Typically, a few drops of appropriate indicator solution are added to the analyte flask before the titration begins that will produce a color change near the stoichiometric equivalence point of the reaction between the titrant and the analyte. Thus, the experiment provides a visual prompt when the flask contains stoichiometrically equivalent amounts of each substance, and the volume of titrant may be determined with high precision from the buret.

A word about high precision buret readings: Remember that the buret has tick marks at every tenths place; therefore, volume measurements should be recorded to the hundredths place, since we always estimate one digit beyond the available tick marks.

Given the balanced acid-base neutralization reaction: x HA + y BOH reaction products where x and y are coefficients of the balanced equation, we may solve for the molarity of the unknown acid by considering the necessity for moles of each reactant to be in stoichiometric equivalence, as shown below:

volumeof acid×molarityof acid=volumeof base×molarity of base× xmolesacidy molesbase

Background - KHP and Standardization: KHP, potassium hydrogen phthalate, is a weakly acidic compound known for its air-stable properties, combined with its non-hygroscopic property or avoidance of water absorption and adsorption. Water is surprisingly well able to adhere to many substances; a substance that can dissolve readily in water but which does not absorb water vapor in air is relatively rare. These properties make KHP an excellent choice for determining the concentration of basic solutions. Since the basic solution may, in turn, be used as a standard against which to evaluate the concentration of an unknown acid solution, KHP is known as a primary standard.

Substance A (or titrant)

Substance B (or analyte)


moles of acid moles of base

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KHP has the structural formula shown to the right and the empirical formula KHC8H4O4. The compound will fully dissociate in water, producing potassium cations and hydrogen phthalate anions in solution. The HP− ion (where P stands for phthalate, not phosphorus) is weakly acidic and will react with sodium hydroxide as shown:

K + + HP− + Na+ + OH− K + + Na+ + P2− + H2O (Note that spectator ions K + and Na+ are shown to illustrate overall charge balance.)

The reaction above is a neutralization reaction and is considered to proceed completely to the right. The molar mass of KHP is 204.22 g/mol.

Background - Phenolphthalein and endpoint v. equivalence point: Phenolphthalein is an acid-base indicator that is colorless in acidic solutions and turns pink-fuchsia in the pH range from 8 -10. Section 8.6 in your text (Zumdahl & Zumdahl, 7th Edition) shows several color pictures of interest.

The equivalence point of an acid-base titration is defined as the point at which the moles of acid are in stoichiometric equivalence with the moles of base. This point may be thought of as a pH value and simultaneously as a value representing the volume of titrant added. Figures 8.9 and 8.10 on page 326 of the text illustrate examples of titration curves and equivalence points. Note that the equivalence point of a strong acid-strong base titration occurs at a pH of 7, while the equivalence point of a weak acid-strong base titration occurs at a pH larger than 7. (Do you know the reason for this? Think about the properties of the respective reaction products.)

The endpoint of a titration is defined as the point at which a permanent color change is observed. Once again, this point may be thought of a pH value and simultaneously as a value representing the volume of titrant added.

The equivalence point and the endpoint of a single titration may have significantly different pH values; however, if the indicator solution is chosen wisely, the two points will have very similar values for volume of titrant (for this experiment, NaOH) added. It is helpful that the titration curve tends to be steep near the equivalence point; in fact, the equivalence point may be determined experimentally as the steepest portion of the curve or as the inflection point where the curve transitions from being concave up to being concave down.

In this experiment, we will not collect data to compile a titration curve; we simply find the observable endpoint and trust that our endpoint is an excellent approximation of the equivalence point. (You will observe that a single drop of sodium hydroxide solution can produce a remarkable change in color as you approach and reach the endpoint; this is confirmation that a small change in volume may result in a large change in pH.)

Background - Percent dissociation: (See also Section 7.5 in the course text, 7th Ed.)In evaluating acids, the percent dissociation describes the relative concentration of acid molecules that have dissociated to produce ions in comparison to the initial concentration of acid molecules. For this experiment we may assume that all H+ (or hydronium ion, H3O+), is present as the result of dissociation from the acid. In other words, we assume that the H+ production from water may be considered negligible in comparison with that from the acid.


Structure of KHP

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The titration allows us to determine the molarity of the acid, [HA]0, and we may use pH of the initial acid solution (before titration) to determine the concentration of H+, which we take to be the concentration of acid that was dissociated.

Thus, for monoprotic acids,

Percent dissociation=amount dissociated (mol

L)

initial concentration(molL

)×100%=¿¿

It may be further demonstrated for weak monoprotic acids where [HA] ≈ [HA]0 that

K a=( percent dissociation100% )2

×[HA ]0

Goals:

Use titration methods to determine initial concentrations of at least two acid solutions.

Use titration data and pH data to calculate percent dissociation for all acid samples.

(Honors) Compute Ka values for the weak acid samples; compare calculated values to true values, and report percent error.

“Along-the-way” Goal:

Perform standardization using potassium hydrogen phthalate, KHP, as the primary standard.


burets (to be checked out from the stock room; return to stock room with stopcock in the open position)volumetric pipets, assorted sizes, may be checked out from the stockroompH meters (will be set up in the lab; TA will explain usage)


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Chemical List

~ 0.05 M sodium hydroxide solution (Shake well before use to disperse possible concentration gradients in bottle.)potassium hydrogen phthalate, KHPphenolphthalein indicator solution

Solutions with unknown molarities:Note that unknown acid solution molarities will fall in the range 0.01 M - 0.2 M.hydrochloric acid - Solutions A, B, and Cacetic acid - Solutions D, E, and F


Pre-lab Exercises:


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Note that pre-lab questions are fairly extensive for this experiment. As compensation, there will be no post-lab questions, and points normally applied to post-lab questions are counted toward pre-lab questions for grading purposes on this lab.

1. Calculate the molarity of sodium hydroxide solution if 0.395 g of KHP (dissolved in ~50 mL heated distilled water) takes 21.34 mL of the sodium hydroxide solution to reach the endpoint.

2. Calculate the molarity of hydrochloric acid if a 20.0 mL sample of the acid requires 12.85 mL of 0.875 M sodium hydroxide solution to reach the endpoint.

3. If you expect the stockroom solution of sodium hydroxide to be 0.0500 M, and you want to use 12.00 mL of the solution to provide a guaranteed 4 significant figures in your volume measurements, determine how many grams of KHP you should use to perform your standardization.

4. If you expect the unknown acid molarities to be between 0.010 M and 0.20 M and the sodium hydroxide solution to be 0.050 M, what volume of unknown acid solution would you place into the flask? State your reasoning or assumptions that you make (perhaps typical of those given in question 3, perhaps not) and show your calculation.

Bonus. State two possible errors when using burets that would reduce the accuracy of your titration. Explain specifically whether the result would be a calculated acid molarity that is smaller than the actual value or a calculated acid molarity that is larger than the actual value.




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Beginning Question Seeds: How does percent dissociation change with varying concentration of a weak acid? How consistent are calculated Ka values for a weak acid across titrations performed on the

same solution? …performed on different solutions with varying concentrations? What factors (e.g., initial pH, strong v. weak acid, acid molarity, etc.) affect titrant volume

required to reach an end point? Are class results consistent enough to justify a claim?

Selected Question:



Procedure:

Review your notes from the Skill Building Inter-session as needed for reminders about using burets for titration. Describe (as though to someone who did not have access to your Skill Building notes) a brief procedure for performing the necessary titrations. Note the availability of “free” data tables on Webcampus; these may be helpful in guiding your Procedure.

Your Procedure should include specific amounts of KHP to use for standardization and amounts of acid solution to use for titration; your responses to the pre-lab exercises may be useful here. Your Procedure should also note specific glassware to be used.

For pH meter work, it is acceptable to simply state that pH meter will be used with TA assistance.

Safety Notes:

Use common sense and/or prior experience to suggest at least one safety note specific to this lab.

In-Lab Requirements

NOTE: Be prepared to work together as a class and stay as long as it takes to get all the data you will need.


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Spend about 5-10 minutes talking as a whole class. Consider and agree on: What are some of the hypotheses we are making, and what data must be collected to prove

or disprove our hypotheses? Agree on and summarize the procedure to be used. How will the data be organized? (Once again, “free” sample data tables are provided to use,

modify, or ignore in favor of creating your own from scratch.) How will we divide labor? What are the collective safety concerns?

Students may modify or completely re-write their Hypotheses, Procedures, and Safety Notes as a result of this discussion.

Conduct the experiment and carefully note your Data, Observations, and Calculations in your lab notebook. Report all data for class evaluation. Be sure to stay and get the complete set of whole class data. As always, clean up; wipe down your lab bench area, etc. when your work is complete. If you finish early, find other groups to help.

Post-Lab Discussion (takes place in lab):Discussions will be student-directed and may address results, calculations, and anything that is relevant or helpful.



Reminder:Remember that one representative calculation should be shown for each calculated value (base molarity, acid molarity, and percent dissociation), in the Data, Observations & Calculations section of your lab manual. (Honors: add Ka to the above list.)


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(This does not count toward the two page maximum.)

Results:Provide a standardization table summarizing the KHP mass, sodium hydroxide beginning and final volume, net volume, and calculated sodium hydroxide molarity for all standardization trials conducted by the class. (If, for any reason, you wish to discard a portion of the data, explain your reasons in a sentence or two.) Calculate and note the average sodium hydroxide molarity; use this value for future calculations of acid molarities.

Provide a second, titration table summarizing the acid solution identity, “solution name” (Solution A, B, C, D, E, or F), initial pH of the acid solution, acid volume titrated, sodium hydroxide beginning, final, and net volume, calculated acid molarity, and percent dissociation for all trials conducted by the class. Provide any additional information needed to support your claim and reflection discussion.

Honors (other sections optional): Report Ka and compare with “true” Ka value for acetic acid = 1.8 x 10−5; report percent error.



Post-lab questions:

None; pre-lab questions were extensive for this lab; points normally applied to post-lab questions are counted toward pre-lab questions for this lab.


What did you learn about titration as an analytical chemical method, as a result of this lab? (e.g., What is a "good" titration? How do you know if you performed a “good” titration? What are some potential pitfalls when performing titrations?) What did you learn about acids as a result of this lab? (e.g., What trends were you able to observe? Compare your experimental results to theory presented in your text, lecture, or from other reputable resources.)

Experiment 5: Calorimetry and Thermodynamics

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Experiment 5: Calorimetry and

Thermodynamics

Overview: In Experiment 5, students will utilize both calorimetry and solubility as a means of calculating thermodynamic properties. In particular, students will use observable phenomena to increase their understanding of abstract thermodynamic concepts.


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Background: Please read “A Simplified Method for Measuring the Entropy Change of Urea Dissolution” by Charles Liberko and Stephanie Terry1, found on the course Web Campus webpage. Note that two independent experimental procedures are necessary to evaluate the entropy of urea dissolution: a calorimetry experiment that leads to enthalpy and a solubility equilibrium experiment that leads to Free Energy.

Background - Urea: Urea is found naturally in the body (it is the main nitrogen-containing chemical in urine.) Friedrich Wöhler, in 1828, was the first person to produce urea from inorganic starting materials, and urea was the first organic chemical that was synthesized inorganically. (This disproved a popular theory of the time called vitalism.) Urea has chemical formula (NH2)2CO and structural formula as shown to the left.

Background - Calorimetry: (Section 9.4 in the text contains additional resources.)A coffee cup calorimeter is a simple, cost-efficient device that attempts to isolate heat transfer and permits measurement of the temperature change created by a given process. For this lab, the process of interest is actually the sum of two steps: the initial dissolution wherein the urea complex breaks down into single molecules and the subsequent solvation wherein water molecules interact with and stabilize the urea molecules. Note that the dissolution step is typically endothermic, while the solvation step is typically exothermic; the total process may be either endothermic or exothermic; for the purposes of this lab, we have chosen to call both steps together the solvation of urea.

You will assemble a coffee cup calorimeter as shown to the right, or with modifications of your choice. (Note that two, nested Styrofoam cups are indicated; however, the actual number of cups may be varied.) You will place a measured mass of water into the coffee cup(s) and then add a measured mass of urea. You will use a thermometer to measure the temperature change of the solution due to the solvation of the urea.

Be aware of potential sources of heat gain or loss in the real system (e.g., the thermometer, the air, etc.) Also take some time to consider the assumptions being made (e.g., a common assumption for this experiment is that all heat is evenly distributed throughout the solution; we also assume that Styrofoam is a perfect insulator, etc.) How does it help us be effective experimenters to consider these assumptions--what might we say about stirring or not stirring, etc.?

Background - Solubility Equilibria: (Section 8.8 in the text contains additional resources.)The method described in the Liberko & Terry article for determining the equilibrium constant of urea solvation is one of many possible choices. Another method is to make a saturated

1 Liberko, C. A. & Terry, S. J. Chem. Educ. 2001, 78, 1087.

Coffee cup calorimeter


C

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solution with excess solid visibly present, so that solvated product is shown to be in equilibrium with solid reactant. Once equilibrium has been established, students may determine the urea concentration of the solution; the equilibrium constant is equal to this concentration, divided by the unit mol/L to provide the familiar, unitless equilibrium constant.

One method for determining the concentration of the saturated urea solution is to pipet a precisely measured volume of the solution--taking care not to draw any solid into the pipet--into a pre-weighed evaporating dish and heat the evaporating dish until all water has been driven off (using the method of heating to constant mass). This allows determination of mass of solute per liter of solution.

For this lab, we are interested in the equilibrium constant of a solubility reaction. If the reactant were an ionic compound, releasing multiple ions into solution, the equilibrium constant would likely go by the name “Solubility Product Constant,” designated Ksp. Ksp is simply a specialized form of Keq; it is not unusual in literature sources to see the two terms used interchangeably.

It is worth noting that the solubility product constant is significantly different from solubility. Solubility is commonly expressed either as grams of solute per liter of solution or more commonly as grams of solute per 100 g solvent, while the solubility product constant is unitless. Similarly, the numerical values of the two are different, which is reasonable since they, in fact, are measurements of different things.

Background - Thermodynamic Properties: Free Energy, ΔG°, may be related to the equilibrium constant, Keq, through the equation ΔG° = −RT ln(Keq), where R is the ideal gas constant, expressed as 8.3145 J K−1− mol−1, T is temperature in Kelvin, and Keq may be determined experimentally. Note that Free Energy and solubility product constants each follow different conventions to signify spontaneous reactions (ΔG° < 0 and Keq > 1, respectively); these relationships are consistent with the above equation.

Enthalpy, ΔH, is equivalent to heat, q, during a constant pressure reaction in which no work is performed. Heat refers to the heat gained by the system or qsystem, which is equal and opposite to the heat gained by the surroundings. In the solvation of a chemical species, the atoms being solvated are the system, and the solution (and more generally the rest of the universe) is the surroundings. Heat gain by a solution is more readily measured and is given by the equation qsurroundings = m s ΔT, where m is the total solution mass in grams, s is the specific heat capacity of the solution, assumed to be the value of the specific heat of water, 4.184 J C−1 g−1, for dilute solutions, and T is the change in temperature (final temperature minus initial temperature) inΔ degrees Celsius (or in Kelvin; values are the same when taking the difference). Enthalpy of solvation, then, is given by H = Δ qsystem = −qsurroundings = −m s ΔT. Temperature change may be measured experimentally using a coffee cup calorimeter. Though not sophisticated, the coffee cup calorimeter allows good thermal isolation of the solution from its reaction vessel and has demonstrated close agreement with literature values for enthalpy in past experiments.

A quick warning about units: the equations above will each produce values with units of J/mol; however, Free Energy and Enthalpy are most often reported in units of kJ/mol.


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From the Free Energy and Enthalpy values developed from experimental evidence, students will calculate the entropy (ΔS°) of urea solvation, which will be reported in units of J mol−1 K−1. Entropy is determined from the equation ΔG° = H° − T S°, where T is temperature in Kelvin.Δ ΔFinally, students will qualitatively and quantitatively compare their calculated values to the accepted values. Based on this comparison, students will evaluate the assumptions they have made and consider means of correcting for those assumptions.

Goals:

Calculate the Free Energy (in kJ/mol), Enthalpy (in kJ/mol), and Entropy (in J

mol K )

of urea solvation.

Perform a class comparison, such as an average and standard deviation of outcomes resulting from following the same procedure or an analysis of trends in outcomes resulting from a varied procedure. Evaluate and rationalize (attempt to make sense of) the findings. (If the average and standard deviation are computed, and the standard deviation is large, what specific factors do you think contributed to its value--explain why. Or if no trends were detected from a varied procedure, why do you think they were absent; what does this say about the factor being varied?)

Compare calculated values to literature values (including performing a percent error calculation), and evaluate potential sources of error (including assumptions).


Styrofoam coffee cups, assorted sizes(lids may be provided or may need to be improvised: paper towels, etc. may be used)pipets (assorted sizes may be checked out from the stockroom)

Chemical List

urea

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Pre-lab Exercises:

1. Write the chemical equation for urea solvation--your own or the one found in the Liberko & Terry article. Also write the expression for Keq that corresponds to the equation.

2. The following solubility data is collected: 10.00 mL of saturated urea solution at 21.3°C has an initial total mass of 11.684 g, and the final solid urea mass (after reaching constant mass from evaporating off the water) is 6.198 g.Find:

a) Keq at 21.3°C andb) ΔG° at 21.3°C.

3. Inspect the following calorimetry data for the dissolution and solvation of sodium hydroxide.

Trial #1 Trial #2mass of sodium hydroxide 2.053 g 1.965 gmass of water 50.5 g 99.9 ginitial temperature of water 21.0 °C 21.2 °Cfinal temperature of water and sodium hydroxide

31.3 °C 26.5 °C

a) Calculate how many moles of NaOH are participating in the reaction for each trial. Be sure your answer has the correct number of significant figures.

b) Calculate the heat, q, (in kJ) absorbed by the solution for each trial.

c) Enthalpy always refers to enthalpy of a reaction and is stated in units of kJ per “mole of reaction” which means per mole of NaOH dissolved in this case. Calculate the enthalpy of dissolution for each trial; watch the sign and use units of kJ/mol.

d) The true value of enthalpy of dissolution for sodium hydroxide, ΔH°, is known to be -44.51 kJ/mol. Calculate the percent error for each trial.

e) Given that enthalpy at standard state, ΔH°, is based on solution molarities of exactly 1 M, which trial more closely represents the standard state?

f) Copy the following sentence into your lab notebook and state ALL possible answer choices that apply:

From the temperature increase and the ease of sodium hydroxide solvation, it may be summarized that the solvation of NaOH is a(n) ___________________ process.

(i) endothermic(ii) exothermic(iii) spontaneous(iv) non-spontaneous


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Honors. What does the structural formula of urea suggest about its solubility in water? Your explanation should include types of bonds (polar, non-polar, hydrogen, etc.).



Beginning Question Seeds: How does sample size of urea (range: 0.2 g - 2.0 g) affect the accuracy of enthalpy

measurements using the coffee cup calorimeter, if water mass is kept constant at 100.0 g? What effect does the number of nested coffee cups (range: 1 - 4) have on the accuracy of

enthalpy measurements? How does urea solubility data change with changing water temperature? (How might the

function be shaped compared to a straight line?)

Selected Question:



Procedure:

Describe a brief procedure for both calorimetric and solubility testing of urea solvation. Include approximate amounts (masses or volumes, as appropriate) of urea and other reagents that will be used for each test. Also include details about the glassware or other items to be used.

Safety Notes:

Use common sense and/or prior experience to suggest at least one safety note specific to this lab. (Hint: a MSDS for urea is located on WebCampus.)


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In-Lab Requirements



or disprove our hypotheses? Agree on and summarize the procedure to be used. How will the data be organized and reported for whole class use? How will we divide labor? What are the collective safety concerns?


Experiment:Conduct the experiment and carefully note your Data, Observations, and Calculations in your lab notebook. Report all data for class evaluation. Be sure to stay and get the complete set of whole class data. As always, clean up; wipe down your lab bench area, etc. when your work is complete. If you finish early, find other groups to help.




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Reminder:Remember that one representative calculation should be shown for each calculated value (heat gained by solution, enthalpy, molarity of the saturated solution, free energy, entropy, and percent error), in the Data, Observations & Calculations section of your lab manual. These calculations should be written neatly, show all work and include units and systematic unit cancellations. Answers should be to the correct number of significant figures.(This does not count toward the two page maximum.)

Results:Provide a table summarizing class results for any variable that was purposefully altered in the experiment and the calculated enthalpy, free energy, and entropy of urea solvation. Provide any additional information you deem important.

Add to the above table (or provide a second table, modified as is appropriate, based on your selected question) the percent error of each enthalpy, free energy, and entropy value, based on comparison with the literature values as reported by Liberko & Terry. Also, explicitly state the literature values you are using for comparison; this may be written in paragraph form, placed as a footnote in the table, or take the form of a small table of their own. Provide any additional information needed to support your claim and reflection discussion.



Post-lab questions:

1. Is the overall solvation of urea endothermic or exothermic? Explain how your data supports your answer.

2. Is the overall solvation of urea spontaneous or non-spontaneous? Explain how your data supports your answer.

3. What do you believe are your 2-3 most significant sources of error for this lab? Explain your reasons for selecting these sources.


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What did you learn about calorimetry and/or enthalpy as a result of this lab?

What did you learn about solubility equilibria and/or free energy as a result of this lab?

Do your best to rationalize the class findings. (e.g., What broader trend were you able to observe, if any? How do you make sense of the overall data? Depending on the Selected Question, this discussion may echo your Claim or may expand beyond it.)

Honors: Discuss assumptions that were made for this experiment. (e.g., How valid do you think these assumptions were? How might they have affected your results? Did the experimental procedure and/or results justify the assumptions?)


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Experiment 6: Redox Titration

Experiment 6 : Redox Titration

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Overview: In Experiment 6, students will be presented with whole citrus fruits (oranges, lemons, limes, and grapefruit) and perform titrations to collect data about the amount of ascorbic acid (Vitamin C) present in samples of fresh-squeezed juice. Students may also bring in their own samples of other juices (e.g., alternate fruits than those listed above, organic fruit juices, reconstituted frozen juice concentrate, juice box juice, 3% juice products, etc.) or Vitamin C supplements. Note that any sample that enters the lab ceases to be a food item and becomes a non-consumable science supply.

Acid-base titration is a familiar concept from Experiment 4. Redox titration is conceptually and procedurally very similar; in both cases, a pair of reactants--with a known, balanced chemical equation--produces a color change near the equivalence point of the reaction. Instead of an acid-base pair, however, we have an oxidizing agent-reducing agent pair. In Experiment 6, the titrant, present in the buret, is the oxidizing agent, an iodine solution prepared by the stockroom. The analyte, present in the flask, is the reducing agent, a sample containing ascorbic acid (Vitamin C).

Standardization is necessary for redox titrations for the same purpose as it is for acid-base titrations: to determine the molarity of the titrant solution. Pure ascorbic acid, from a chemical supply company, will be used as the analyte for the standardization.

Most importantly, this experiment reflects ongoing analytical processes in industry today.

Background: The following background is borrowed verbatim from Sowa, S.; Kondo A.E. J. Chem. Educ. 2003, 80, 550. (Supplemental materials)

Vitamins are essential micronutrients in our diet. We need them in small amounts to stay healthy. Although they were originally thought to be “vital amines,” we now know that there are two categories of vitamins: fat-soluble (A, D, E, K) and water-soluble (the B complex and C). Most vitamins act as cofactors with metabolic enzymes to maintain our life processes.

Vitamin C, or ascorbic acid, is a water-soluble molecule that acts as an antioxidant, i.e. if oxidizing agents attack a cell, then the vitamin C acts as a target in place of another more important target such as a protein. Vitamin C also acts as a cofactor for an important enzyme that synthesizes collagen, the most abundant protein in vertebrate animals (such as humans). Collagen is a fibrous protein in connective tissue that literally holds us together. Vitamin C


titrant

analyte

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deficiency results in a disease called scurvy, which is characterized by skin lesions, bleeding gums, and tooth loss.

Historically, vitamin C played an important role in the discovery of the “New World.” Sailors on long voyages with no available fresh fruit or vegetables commonly developed scurvy, which was described by Jacques Cartier during his exploration of North America in 1535. In 1768, James Cook introduced the addition of lime juice to the diets of British sailors, who then became known as “limeys.” By 1795, Lind advocated fresh citrus as a cure for scurvy. Dutch and Hungarian soldiers also used sauerkraut as a source of vitamin C, and American sailors carried cranberries. The chemical structure of vitamin C was not discovered until 1932. More recently (1970’s) the two-time Nobel laureate Linus Pauling controversially advocated large doses of vitamin C as a prevention for the common cold.

Vitamin C can be measured using a technique called titration. Because vitamin C is easily oxidized, we can use this chemical reactivity to detect it in solution. If we react vitamin C with iodine, the vitamin C is oxidized and the iodine is reduced. Iodine (I2 molecules) form a purple complex with starch in solution, but reduced iodine, iodide (I−) is colorless in the presence of starch.

If we add iodine to react with the vitamin C in an acidic solution that also contains starch, then once we’ve oxidized all of the vitamin C, the I2 will remain, and a purple color will be produced. This color change is the basis for the titration reaction, and the purple color is used to determine the endpoint.

Goals:

Determine and compare the molarity of ascorbic acid (Vitamin C) in a variety of samples.

Explain which species are oxidized and which are reduced; provide a practical description of the term “antioxidant.”

Evaluate results for consistency and repeatability. (e.g., Qualitatively, describe the degree of “scatter” in molarity values for each type of fruit juice--or other sample type--tested; quantitatively, calculate averages and standard deviations--reported to the first significant figure only, and report average ± standard deviation values to a reasonable number of significant figures, given the place value of the standard deviation.)

“Along-the-way” Goal:


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Perform standardization using pure ascorbic acid to determine the molarity of the stockroom iodine solution (or standardize the stock

solution).


burets (to be checked out from stockroom)manual juicersgloves (available from stockroom, wear when handling iodine solution)additional pipets (graduated or volumetric) may also be checked out from the stockroom

Chemical List:ascorbic acid, pureiodine solution, ~ 0.02 M concentrationstarch solution (used as indicator)0.2 M acetic acid, provided in a dropper bottle, used to acidify ascorbic acid samples

Fresh Produce Provided by the Stockroom:grapefruitlemonslimesoranges

Note that additional items may be brought by students; all items are to be treated as science supplies and may not be consumed.


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Pre-lab Exercises:

1. If you expect the stockroom iodine solution to be 0.0200 M, and you want to use 30.00 mL of the solution, determine how many grams of ascorbic acid you should use to perform your standardization. (The molar mass of ascorbic acid is 176.12 g/mol.)

2. In this experiment, what chemical species is being reduced? Write the balanced half-reaction showing the reduction of this species; include H+ ions and electrons as needed to balance chemicals and charge. Note oxidations numbers for the reactant and the product of interest.

3. In this experiment, what chemical species is being oxidized? Write the balanced half-reaction showing the oxidation of this species; include H+ ions and electrons as needed to balance chemicals and charge. Note oxidations numbers for the reactant and the product of interest.

4. (Re-read the background section as needed.)

a. What color is “reduced iodine” in the presence of starch?

b. What specific color change (from what to what) will you see when you reach the endpoint?

c. What complex (combination of species) causes the color?



Beginning Question Seeds: How much random variation is there in ascorbic acid molarities of the same type of fruit

compared to the differences detected among different fruit types? (e.g., Will the values for average molarity of lime juice ± its standard deviation have no overlap, some overlap, or significant overlap with the values for average molarity for orange juice ± its standard deviation?)

How does the ascorbic acid concentration for Emergen-C (brought by students, prepared per instructions) compare to the ascorbic acid concentration from fresh squeezed juice? (Or, what percentage by weight of dry Emergen-C is pure ascorbic acid?)

How does the ascorbic acid concentration for fresh oranges compare between regular, non-organic oranges (supplied by the stockroom) and organic oranges (brought by students)?


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Note that students are not required to use fresh fruits provided by the stockroom; they may elect to use all four types of fruit, or none at all, or anything in between. Students who write beginning questions relating to products not supplied by the stockroom must be prepared to bring in those products, to become non-consumable science supplies.

Selected Question:



Procedure:

Write a brief procedure for performing the necessary titrations, including the standardization. Your Procedure should include specific amounts of pure ascorbic acid to use for standardization, any filtering (for removal of pulp, perhaps) or other actions to be taken on the juice samples, and amounts of juice (or other substance) to use for titration; your responses to the pre-lab exercises may be useful here. As usual, your Procedure should also note specific glassware to be used.

Safety Notes:

Use common sense and/or prior experience to suggest at least one safety note specific to this lab. The MSDS for iodine is available on WebCampus and may be a useful tool.


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In-Lab Requirements

NOTE: Be prepared to work together as a class and stay as long as it takes to get all the data you will need.



or disprove our hypotheses? Agree on and summarize the procedure to be used. How will the data be organized? How will we divide labor? What are the collective safety concerns?






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Reminder:Remember that one representative, example calculation should be shown for each calculated value (molarity of iodine solution, molarity of juice sample, average molarity, standard deviation of molarity), in the Data, Observations & Calculations section of your lab manual. (This does not count toward the two page maximum.)

Results:Provide a standardization table summarizing the ascorbic acid mass, iodine solution beginning and final volume, net volume, and calculated iodine solution molarity for all standardization trials conducted by the class. (If, for any reason, you wish to discard a portion of the data, explain your reasons in a sentence or two.) Calculate and note the average iodine solution molarity; use this value for future calculations of ascorbic acid molarities.

Provide a second titration table summarizing the sample identity, sample volume titrated (or mass, if a dry sample product was brought in to lab), iodine solution beginning, final, and net volume, and calculated ascorbic acid molarity (or grams for a dry sample).

Provide a results section, formatted however you wish--or extend a table from above--stating for each type of sample tested, the average molarity and standard deviation for ascorbic acid molarity of that sample type. Adjust as needed for dry samples or different styles of selected questions.




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Post-lab questions:

1. a) What was your most consistent data set? What do you think made it so consistent?b) What was your least consistent data set? Suggest a plausible reason for the lower consistency.c) What effect do these forms of scatter have on your evaluation of the data?

2. The human body contains reactive oxygen species (including oxygen ions) that may damage tissue. Scholarly articles (e.g., Fukumura et al, J Physiol Sci. 2012, 62, 251-7. and Guaiquil et al, J Biol Chem, 2001, 276, 40955-40961.) suggest that antioxidants may be beneficial to cancer patients. Based on what you have learned in this lab, explain in your own words what an antioxidant is and suggest how it functions in the human body.

Honors. Locate the abstract for either above-mentioned article. (Google scholar is a helpful tool for this; the first article title is, “Effect of ascorbic acid on reactive oxygen species production in chemotherapy and hyperthermia in prostate cancer cells,” and the second is, “Mechanism of Vitamin C Inhibition of Cell Death Induced by Oxidative Stress in Glutathione-depleted HL-60 Cells.”) State which article you looked up (“first” or “second” is sufficient) and what a major finding was.


What did you learn about redox reactions or about antioxidants as a result of this lab? (One option for addressing this is to compare your experimental results to theory presented in your text, lecture, or from other reputable resources such as scholarly articles.)

What was unexpected about this lab? (e.g., Did some of the results surprise you? Compare redox titration to acid-base titration--was the color change trickier, or were there other differences that were unexpected?)


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http://www.ncbi.nlm.nih.gov/pubmed/22392350

Experiment 7: Atomic Spectroscopy

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Overview: In Experiment 7, students will examine the atomic emission spectra of several elements. (The emission spectra of atoms are typically much easier to view than the absorption spectra; observing a small number of lines is easier than observing the “black lines” in a continuous visible spectrum.) Students will gain experience in the calibration process and will collect data allowing them to calculate the Rydberg constant.

Background: Photons are elementary particles that are the force carriers of the electromagnetic force, i.e., electromagnetic radiation. As was demonstrated by Einstein’s treatment of the photoelectric effect, the energy of a single photon (Ephoton) is related to its frequency (ν) by the relationship:

Ephoton = hν (1)

where h is the Planck constant (6.626 × 10-34 J s) and is related to the speed of light (c) by:ν = cλν (2)

Thus, the wavelength of light ( ), its frequency, and the energy of a photon are all interrelated.λ

The electromagnetic spectrum extends from very high frequencies (high energy, short wavelength) to very low frequencies (low energy, long wavelength); see Figure 1. Only a small portion of the electromagnetic spectrum is visible to the eye. Visible light covers the region with wavelengths of 400 - 700 nm ( = 7.5 × 10ν 14 − 4.3 × 1014 Hz). This region of the spectrum promotes so-called electronic transitions. Essentially, these can be thought of as the photon of light possessing equal energy to that required for the transition of an electron from a low-energy orbital to a higher energy orbital.[2] When a photon with energy equal to the electron’s transition energy is absorbed by the atom or molecule, it causes the promotion of an electron to a higher energy orbital, effectively absorbing and storing the energy imparted by the photon.

The inverse of this process is called emission. In an emission event, the electron begins in an excited state and decays to a lower energy state. As the energy of the atom is now lower, and energy must be conserved, the “lost” energy is emitted as a photon. Because the absorption and emission events are the inverse of one other, the energies involved will be related. As shown in Figure 2, the energy of the photon absorbed in promoting an electron from state x to y will be the same energy as the photon emitted in going from state y to x.

2[] This statement is not technically correct. The transitions are occurring from one electronic state to another, not one orbital to another. However, many visualize (to a first approximation) the transitions as occurring between orbitals.


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The energy levels of atoms (and molecules) are not continuous, but quantized. That is, there are only specific energies that an atom can absorb or emit. (This can be seen from the relatively narrow lines present in an atomic emission spectrum and can be related to the principal quantum number, n.) Further, these energies vary with the identity of the atom (or molecule).

Figure 1: The Electromagnetic Spectrum

Increasing Energy


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Figure 2: An atomic absorption (A) and emission (B) event.

The Rydberg equation is an empirically derived formula that can be used to predict the wavelength of light absorbed or emitted by hydrogen. It takes on the form:

(3)

where R is the Rydberg constant (R = 1.09677581 × 107 m−1), and n1 and n2 are integers with n2 > n1. Note that a small subset of integer values for n1 and n2 will yield wavelengths in the visible range.

Background - the Spectroscope: The spectroscope is essentially the same instrument originally used to record atomic spectra in the late 1800s and early 1900s. Although optical instrumentation has become more advanced, the principles behind modern devices have their roots in the simple spectroscope.

The light scale will possess tick marks labeled as nanometers. It will be tempting to assume that the values you see are accurate. Experienced spectroscope users have found that this scale is not truly reliable, but it may be converted into reliable readings by the process of calibration. In other words, the scale on the spectroscope device provides only a rough, (some have even


Figure 3 displays a schematic of a spectroscope. The spectroscope works by allowing light into the instrument through the entrance slit. The light then hits the diffraction grating, is separated into its component wavelengths, and is reflected onto the light scale. (The eyepiece, not shown, is on the same side as the diffraction grating. Also not shown is a focusing knob, which should be used to obtain vivid, sharp lines of color--rather than thick, blurry lines.)

Figure 3: Schematic of a Spectroscope

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said “arbitrary” or “meaningless”) first approximation of the actual wavelengths. To gain precision and accuracy, it is necessary to develop a correction protocol or calibration.

Before you begin the calibration, you should look inside the spectroscope and sketch the scale as accurately as possible into your notebook. While wearing gloves, place the mercury lamp in the power supply and turn it on. For optimal lamp lifetimes, the lamp should only be on ~30 seconds at a time. Position the spectroscope such that the slit is pointed at the center (brightest section) of the mercury lamp. This is a good time to adjust the entrance slit knob to obtain the sharpest, highest resolution lines (resolution is related to the width of the slit opening). Next, you will record where the atomic emission lines are on your sketch of the scale and note the colors observed. Record other, qualitative notes, as you deem appropriate. Turn off the lamp and allow it to cool before removing it.

You will need to develop a calibration equation that maps the observed line values onto the known values--the emission lines for mercury are known to be 404.7 nm (violet), 435.8 nm (blue), 546.1 nm (green), and 579.0 nm (yellow). Then you will use your equation to correct the observed values for all future readings. (You are welcome to use a graphing program to develop the calibration equation.)

Alternately, the emission lines for helium may be used: 402.6 nm (violet), 447.1 nm (violet), 501.6 nm (green), 587.6 nm (yellow), 667.8 nm (red), and 706.5 nm (red). It is possible that all lines noted here may not be observed due to some having a lower intensity; alternately, other even less intense lines exist but have not been included here.

Goals:

Calibrate the spectroscope based on known emission lines of mercury and/or helium.

Compare your calibrated hydrogen emission wavelength values with the expected, theoretical values from the Rydberg equation (theoretical values are developed in pre-lab exercise #1). Evaluate the degree of agreement.

Graph a scatter plot with x values set to 1/n22, where n2 values are the theoretical values

of n2 for hydrogen emission developed in pre-lab exercise #1, and the y values are 1/ ,λ where λ values are the calibrated, experimentally observed wavelengths of hydrogen. The x and y pairs should be ordered such that each inverse-wavelength y value corresponds to its theoretically matched 1/n2

2 x value. Perform a best-fit analysis of the slope to estimate your own version of the Rydberg constant--name it after yourself, if you like. (A useful equation/expression is developed in pre-lab exercise #1.) Compare this value to the known, literature/textbook value for the Rydberg constant.

“For Fun” Goal:

Record the atomic spectrum of at least one other element besides mercury and hydrogen, either helium to perform an alternate calibration or one or more of the other elemental lamps available: oxygen, bromine, chlorine, neon, argon, krypton or xenon.


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Honors: you are required to observe helium for a post-lab question.


SpectroscopeHydrogen, helium, and mercury lamps (12 of each available for class use)For optional use: oxygen, bromine, chlorine, neon, argon, krypton and xenon lamps (1 of each)Power source for lamps


Pre-lab Exercises:

1. The Rydberg equation yields four emission lines with wavelengths that fall in the visible region (400 - 700 nm); these are called the Balmer series and occur when n1 = 2. Determine the four values of n2 that yield visible wavelengths. Calculate and state the wavelength for each value of n2. Show work or state resources used.

2. Convert the Rydberg equation, shown as equation (3) on page 77, into an equation of the form y = mx + b, where y = 1/λ and x = 1/n2

2. What is the expression for the slope? What is the expression for the y-intercept?

3. The calibration step is critical; it is helpful to perform a practice calibration as a pre-lab question. If the following figure represented your observations when viewing the mercury lamp, what would your calibration equation be? Colors from left to right are violet, blue, green, and yellow. (As stated previously, you are welcome to use a graphing or spreadsheet program to develop the calibration equation.)

Not sure how to start? Hint: Consider a line of the form y = mx + b; the variable x may be placed in the equation and get mapped on to the variable y; spreadsheet programs are very good at finding the best fit lines--using functions slope and intercept in Excel, for example--for any set of x and y data.


Example:

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Please note that once a calibration equation is developed, it should be used to interpret all subsequent experimental data.


Consider the goals of the laboratory experiment; they are fairly specific and don’t lend themselves to a great deal of original questioning. Feel free to select one of the Beginning Question Seeds below, modify it if you wish, or write an Original Beginning Question of your own that you would be interested in investigating--with kudos for originality and effort.

Beginning Question Seeds: How do the calibration equations compare for mercury calibration and helium calibration? Does the number of visible emission lines for noble gas element lamps relate to the atomic

number (e.g., will larger atoms have more complex spectra)? How close is the Rydberg constant developed from hydrogen emission line data to the

literature value provided in this lab? (Does the range given by the class average ± class standard deviation cover the literature value?)

Can the Rydberg equation be used for any other elements besides hydrogen?

Selected Question:



Procedure:

Important Notes:

In order to have an atom emit light, it must be in an excited state. Therefore, energy must be “pumped” into the system to promote an electron from ground to excited states. For this lab, we use electricity in vapor lamps as the energy source; the electricity excites the gaseous atoms in the vapor, and they spontaneously emit light as they return to a lower


Example:

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state. In these vapor lamps there will only be one element present that will be excited, and therefore all of the lines are attributable to that element alone.

CAUTIONS! It is VERY important that you do not touch the lamps with your bare hands. This is done to prevent oils from your skin getting on the lamp, which can damage the lamp itself. These lamps get hot, so allow them to cool down before touching them. These lamps are run by high voltage power supplies, so do not touch any exposed wires, and follow your lab instructor’s directions on how to change lamps. Also, you must wear your safety goggles. This is important (among other reasons) because the lamps can give off a small amount of UV light (the plastic in your safety goggles absorbs the UV light).

Write a couple of bullet points summarizing main points from above and describing the number of trials to be performed for each measurement that you determine to be necessary.

Safety Notes:

Write at least three safety notes.


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In-Lab Requirements

NOTE: As usual, be prepared to work together as a class and stay for the full 2 hours, 50 minutes.



or disprove our hypotheses? Agree on and summarize the procedure to be used. How will the data be organized? How will we divide labor? What are the collective safety concerns?





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Reminder:Be sure that your Data, Observations & Calculations section shows a record of the data you used to obtain your calibration equation, a note about how the equation was obtained (e.g., “data input into Excel; slope and intercept determined by the program” or show work if calculation is performed by hand) and the final (labeled) equation for your calibration. Also show a representative calculation for getting a calibrated wavelength from an experimentally observed wavelength and a sample of any comparison calculation performed between the calibrated, experimental value and the theoretical value.

Similarly, include a record of the input values you used to obtain your own version of the Rydberg constant, a note about how the equations was obtained (e.g., “data input into Excel; slope and intercept determined by the program”) and the final, (labeled) value of your version of the constant. Also show a sample of any comparison calculation performed between your value and the literature/textbook value.

(As always, this does not count toward the two page maximum.) Results:Provide a table for each source you observed, summarizing the observed and calibrated wavelengths.

Summarize your version of the Rydberg constant and its comparison with the literature/textbook value.

Include any additional information pertinent for your claim and reflection discussion.

Claim:Craft at least one Claim directly addressing the Selected Question and your Hypothesis. As always, claims are to be the result of independent student thought and interpretation of the


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data, though they may be guided by the whole class discussion. Copying another student’s wording (whether exactly or with minor paraphrasing) will result in zero points for both students.


Post-lab questions:

1. Given that each emission line represents an energy transition, select one observed wavelength, state the sample from which it was obtained, its raw wavelength value, calibrated wavelength, and the energy released by the emission of a single photon. (Review equations 1 and 2 on page 75, as needed.)

As has been noted previously, the statistical likelihood is expected to be slim that your calculation will match that of any of your lab partners.

2. Describe each of the following from your laboratory observations:

a) The emission line with the largest wavelength

b) The emission line with the smallest wavelength

c) The emission line that resulted from the highest energy transition (i.e., largest Ephoton or ΔE)

d) The emission line that resulted from the lowest energy transition (i.e., smallest Ephoton or ΔE)

Honors. The National Institute of Standards and Technology (NIST) maintains an atomic spectra database at physics.nist.gov/PhysRefData/ASD/lines_form.html. Visit the website and view the data for helium. (Input He in the field for Spectrum, 400 in the Lower Wavelength field, and 700 in the Upper Wavelength field. Verify that the units are nm. Finally, click the Retrieve Data button.) Scroll through the data. Notice that the second column from the left shows the wavelength values in nm. Find a wavelength that matches one you observed in lab, and record the NIST values for:

a) Observed Wavelength Air (nm)

b) Rel. Int.

c) Lower Level Conf., Term, J

d) Upper Level Conf., Term, J

e) Note that each Lower/Upper Level Conf. represents an electron configuration. Write a summarizing sentence of your selected NIST data demonstrating your sense of how it all fits


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together (e.g., “The 1s6d 1s2p electronic transition emits a photon of light at 414.3761 nm with an intensity of 3.”)


Discuss how your results compared to theoretical results. (e.g., Do your results confirm/reinforce the validity of the theoretical treatment or cast doubts on it? Are there reasons that one element may follow the results closely while others may not? Explain the degree of certainty or the number of significant figures you would claim for your estimated Rydberg constant. How does this compare to the known value? How do you think researchers were able to compute the known value to so many digits?)

Experiment 8: Particle in a Box Simulations


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Special Note: The grading for this lab is unique: 15 points will be assigned for the pre-lab and 15 points for the in-lab. There is no post-lab, and the lab reports are due at the end of the lab period. The pre-lab assignment will be due at the beginning of the lab period per your normal procedures, and all answers to the in-lab portion of this experiment must be submitted at the end of lab.

A rubric for this lab is available on the course Webcampus page for your reference.

It is recommended that you bring a textbook and a laptop computer to lab. The stockroom will have a limited number of laptops that may be borrowed (also, @One also has laptops available for 2-day check-out.)

Overview: Experiment 8 is not a traditional wet lab but rather a computational lab. Students will use the 1D particle in a box concept to understand the absorption spectrum of a series of dyes (highly colored organic molecules). In this lab, we will be looking at the electronic absorption spectra of so-called “organic conjugated dyes” (a conjugated system is a series of alternating single and double bonds). Another term for an electronic absorption spectrum is the UV-Vis spectrum; this is the region of light that molecules absorb to promote low energy electronic transitions. Students will analyze spectroscopic data to find the peaks and model the peak values using a parameterized form of the 1D particle in a box solution to the Schrödinger equation. Students will then take the information learned and formulate the structure of a dye that will give an absorption at a specific wavelength.

A significant percentage of modern chemical research relates to theoretical (or computational) chemistry topics. Computational chemistry commonly seeks to model interactions of matter and light, such as absorption and emission events. One measure of a successful model is that, while it is typically developed from a relatively small, simple system, it can be applied to a larger or more complex system without significant loss of accuracy.


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Background:

Spectra showing electronic transitions--when an electron moves from one energy level to another--may be obtained for molecules as well as for atoms. For the series of dye molecules being studied in this lab, important electronic transitions occur in the visual range of the spectrum (which leads to the physical appearance of color); however, other molecules may have their most intense electronic transitions in other ranges of the spectrum. Molecular spectroscopy studies the interactions of electromagnetic energy with molecules; these interactions are more complex than the equivalent interactions with single atoms that were investigated in the previous lab on atomic spectra. While every type of matter has a unique, “fingerprint-like” spectrum, all atomic spectra appear as (relatively) sharp lines, called line spectra, because the electronic transitions are not “muddied” by molecular vibrations and rotations that cause peaks to broaden.

In contrast to the line spectra produced in atomic spectroscopy, we see broad peaks in molecular spectroscopy. These are presented as graphs where the x-axis scans across a wavelength range of interest, and the y-axis shows relative strength of absorbance (Abs.) at that wavelength. See Figure 1, below. The peak maxima correspond to the individual electronic transitions (movement of an electron from some ni level to some n f level). These peak maxima are abbreviated λmax since absorbance is maximized for that wavelength; each such wavelength possesses a corresponding energy which may be taken as the energy of the electronic transition. Notice that the absorbance spectrum shown in Figure 1 has more than one transition (each is a “local maximum” on the graph, but a less intense transition may appear to be “on the shoulder” of a more intense transition).


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Figure 1: Absorption Spectrum for a generic molecule. The values for λmax are 291 nm, 325 nm, and 358 nm.

Background - Dye Structures and Conjugated Bonds:

Conjugated compounds contain alternating single and double bonds. A double bond contains 4 total electrons with 2 electrons in localized, sigma bonds and 2 electrons in delocalized, pi bonds. The series of alternating single and double bonds form a so-called π-system, in which the collection of electrons in pi bonds may be thought of as overlapping 2p-orbitals that allow the electrons to “move” throughout the π system.

Note the conjugated systems in Figure 2 and the structures of dye1, dye2, and dye3 that will be investigated in this lab. The simplest molecule in Figure 2, 1,3-butadiene (H2C=CH-CH=CH2), has two double bonds; thus, it has four electrons in the π-system. The energy levels of these pi electrons are shown in Figure 3.

250 270 290 310 330 350 370 390 410 430-0.2

-0.1

0

0.1

0.2

0.3

0.4

0.5

Wavelength (nm)

Abs

.


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Quantum mechanics can give us some insight into the theoretical energy values for the energy levels shown rather generically in Figure 3. Spectroscopy can show us where 1,3-butadiene absorbs light, and we can convert empirical data about the wavelength into energy ΔE values. Agreement between the empirical values and the theoretical ones provides confirmation of the model behind our theoretical calculations.


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We will also want to know something about the length of the pi system for our calculations. The summed length of the bonds in the π system defines a “box-length” in a 1D particle in a box problem. For 1,3-butadiene, the pi system contains two double bonds and one single bond (the C=C bond length is 1.34 Å and the C-C bond length is 1.54 Å); thus, to a first approximation, we define the “box length” as 4.22 Å for 1,3-butadiene.

Background - Quantum Mechanics and the Particle in a Box Model:

The Schrödinger equation is shown as equation (1) below.

H ψ=Eψ (1)

Solutions to the Schrödinger equation are derived in your textbook and simultaneously solve for wavefunction equations and discrete energy levels. Solving for the Schrödinger equation requires some set of assumptions. Ideally, these assumptions are simple enough that they lend themselves to reasonable, algebraic calculations yet are still able to give close agreement with observed experimental data.

The 1D particle in a box model is a set of very useful assumptions. The electron is said to be a “free particle” with no attractive or repulsive forces acting on it. Further, it is spatially confined to movement along a straight line with length L; by convention, this straight line is called a “1D box” with length L and infinite height, (where the infinite height is described as the walls of the box having infinite potential energy, V (x) = ∞.) . In practice for this lab, the length, L, will be taken to be the total bond length of the pi system. (Note that the pi system lengths of the more complex dye molecules are NOT defined as the sum of all double and single bonds in the entire molecule; the ring structures are typically distinct from the pi system of interest.)

With the constraints of the 1D particle in a box model, the solutions for the energy term in the Schrödinger equation (energy level, En, in terms of the principle quantum number, n) of a 1D particle in a box are:

En=h2n2

8m L2 (2)

where h is the Planck constant (6.626 × 10-34 J s), L is the length of the “box”, and m is the mass of the particle. For this lab m equals me, the mass of an electron.

When an atom or molecule absorbs light, it promotes an electron from a low energy state to a higher energy state. For example, the lowest energy transition of a 1 electron system would involve the promotion of an electron from the n = 1 to n = 2 energy levels. The energy of the photon absorbed must equal the energy difference between the second and first energy levels.


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ΔE=E2−E1=h2

8mL2 (22−12) (3)

or more generally:

ΔE=E f−Ei=h2

8mL2 (n f2−n i2) (4)

where ni is the initial energy level and nf is the final energy level of the particle.

Recall that wavelength, frequency and energy of light are related by:

= cλν (5)

and

ΔE=¿ Ephoton = hν (6)

Thus, the solution for the energy of a 1D particle in a box can be related to the energy of light that can be absorbed due to a transition between two energy levels, which in turn can be related to the wavelength of the absorbed light.

Again, we seek to achieve agreement between the calculated value of ΔE from the 1D particle in a box model and the observed, experimental wavelength.

The spectra being examined have already been recorded and will be provided by the TA during lab. You will analyze the data to find the peaks and model the peak values using a parameterized form of the 1D particle in a box solution to the Schrödinger equation. You will then take the information learned and formulate the structure of a dye that will give an absorption at a specific wavelength.

Goals:

Relate energy to wavelength using the 1D particle in a box solution to the Schrödinger equation.

Perform parameterization to improve agreement between the 1D particle in a box model and experimental

observations. Propose a structure for a dye

that will give an absorption at a specific wavelength.



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Pre-lab Questions (These are the ONLY pre-lab requirement; they are worth 15 points.)

1. For equation 2 (on page 89), the calculated energy should be in units of Joules. To keep the calculations as straightforward as possible, state the value (and units) of the constant me. Also state what units the variable, L, will need to have.

2. Referring to equations 4, 5, and 6, derive an equation that yields λ in terms of h, me, L, c, ni, and n f .

3. For an electron contained in a 10.0 nm box, what wavelength of light would be absorbed to promote the electron from the n = 3 to the n = 5 energy levels? (Use the expression you just created for question 2 above.)

4. (a) Of the transitions depicted in Figure 1 on page 89, which wavelength corresponds to the lowest energy transition? (Hint: this is a straightforward, simple question if you recall the relationship between wavelength and energy.) (b) Assuming the lowest energy transition in Figure 1 corresponds to the promotion of an electron from the n = 2 to the n = 3 energy levels, solve for the length of the box.

5. For 1,3-butadiene, n = 1 and n = 2 are filled (both have two electrons, which is the maximum number in a given energy level for a 1D particle in a box model, as described earlier). Everything from n = 3 and above is unoccupied (no electrons). (a) What is the lowest energy transition you would observe for 1,3-butadiene? Note that to specify this transition, you simply need to state the ni and n f values. (b) What is the lowest energy transition you would observe for 1,3,5-hexatriene? (Revisit Figures 2 and 3 and the discussions and logic behind why 1,3-butadiene has 4 electrons and thus 2 filled shells in its π-system).

6. The λmax for the lowest energy transition of 1,3-butadiene is 217 nm. Calculate the box length L based on this value.

7. This question is rather long and necessitates further explanation.

In reality, the box length for 1,3-butadiene is 4.22 Å (the C=C bond length is 1.34 Å and the C-C bond length is 1.54 Å). Your value (from question 6) should have been longer than this. This is OK. Remember, we are using an approximation that makes a number of physically unrealistic assumptions. Therefore, we will add an adjustment to the 1D particle in a box solution to empirically correct for the inconsistency between our model and the observed value. Note that experimentally observed values are the “true” values we are trying to mimic with our simplified particle in a box model equations. The process of adding a correction term to our equation and solving for the term based on actual data is called parameterizing, and it is commonly used in the sciences to make a model yield physically meaningful results. (A parameterized model must be able to perform well for novel systems, those which were not used in developing the parameter.)

Parameterization of the En values for the 1D particle in a box has been found to be effective by modifying the L term. This modification varies across different classes of molecules (molecules


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that are structurally similar) but may be fairly reliable within a given class. For 1,3-butadiene and 1,3,5-hexatriene, we will replace L in Equation 4(*) with the term (N × l’ + ) where Nα represents the total number of carbon-carbon bonds (both single and double) in the box (N = 3 in the case of 1,3-butadiene); l’ is defined as the average bond length of all bonds within the

“box” region (l’ = 1.34 Å+1.54 Å+1.34 Å

3 = 1.40 Å for 1,3-butadiene); and is a purely empiricalα

parameter that allows us to adapt the model to fit the data. The value that is determined for α should “work” for all molecules that are similar in structure to one another.

(*) Note: the new equation would take the form: En=h2n2

8m(N l’+α )2

(a) Determine the value of for 1α ,3-butadiene. (Let have units of Å.)α

(b) Use the α value from part a to determine λmax for the lowest energy transition of 1,3,5-hexatriene.

(c) Compare the λmax value for 1,3,5-hexatriene from part b to the experimental value of 436 nm. (Perform a percent error calculation and comment briefly.)

(Honors) Compare the agreement with experimental for 1,3,5-hexatriene with the α correction and with α removed (or set to 0.)

In-Lab Requirements


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Pre-Lab Discussion:Submit your pre-lab to your TA, but discuss as long as necessary with your group and with the whole class to be sure you understand the processes for solving the pre-lab questions. Proceed with the in-lab requirements once everyone has a solid understanding of the pre-lab exercises.

Three dyes (all belonging to the “symmetrical polymethine family,” with structures shown in Figure 2 as “dye 1,” “dye 2,” and “dye 3”) have been tested in a spectrophotometer in the Chemistry building on campus.

1. Determine all λmax values for each dye: As a class, plot the spectra for each dye from the data file (provided by your TA); project the spectra on the room’s “big screen.” Identify the lowest energy transition for each dye, and calculate the energy for this transition. Record these in your lab notebook, showing complete calculations. (Work together as a class for one dye, then do the other two independently.) Clearly label your data and calculations with “dye 1,” etc.

Note that the calculated transition energy is a “true,” experimentally observed value; this is the value you hope to match with your parameterized computational model.

2. Determine L for each dye: One of the earliest models of symmetrical polymethine systems by Kuhn3 suggested that pi electrons could only move along and in the direction of the chain connecting the two nitrogen atoms, including one carbon atom beyond each nitrogen. (The benzene rings may be approximated as the walls at which the potential energy, V (x) = ∞.) Due to the high electronegativity of nitrogen, the delocalization of pi electrons decreases in the vicinity of the nitrogen atom. Although this effect is relatively strong, it is not strong enough to stop the propagation of electrons entirely, thus leading to the spread of the electrons to one extra carbon atom along the chain. However, this does not happen to the second carbon atom that is attached to nitrogen since the “branching out” leads to great increase in the potential energy.

For dye 1, a box is shown below that approximately defines the box length for our particle in a box model. Together with the class, calculate the box length, L, for dye 1 assuming that you can sum the length of bonds as though they were arranged in a straight line, with bond lengths as noted below. Then compute L for dye 2 and dye 3, working independently. (Refer to Figure 2.)

Figure 4: Dye 1 structure with the approximate π-system shown in a box. N-C 1.47 ÅN=C 1.28 ÅC-C 1.54 ÅC=C 1.34 Å

3 Kuhn, H. A Quantum‐Mechanical Theory of Light Absorption of Organic Dyes and Similar Compounds. J. Chem. Phys. 1949, 17, 1198-1212.


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3. Determine the number of pi electrons contained in the box for each dye: As a helpful intermediate step before determining the identity of ni and n f corresponding to the lowest possible energy transition of each dye, we will first examine each structure to evaluate the number of pi electrons present in the box described previously. Two pi electrons are present in every double bond; however, we also find that the lone pair of electrons on N (but not N+) act as pi electrons. Note that lone pair electrons are not explicitly noted on most structural drawings but are shown below for clarity. Thus, we may see that dye 1 has 6 π electrons in the box. Determine the number of electrons in the box for dye 2 and dye 3; note all results in your lab notebook.

Figure 5:π electronsfor dye 1,noted withcircles.

4. Determine ni and n f for each dye: For each dye, draw a qualitative orbital energy diagram such as that shown below for 1,3-butadiene, which shows the lowest energy transition, and explicitly identify ni and n f .

For 1,3-butadiene, we may state that ni = 2, and n f = 3; be sure you understand why this represents the lowest energy transition (note, for example, that ni = 2 to n f = 4 is also possible but would require a greater change in energy.)

Figure 6:Qualitative orbital energy diagram for 1,3-butadiene.

:


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5. Determine α for each dye: For each dye, solve Equation 7 shown below to determine a value of α specific to that dye. As in your pre-lab, let have units of Å.α (Note that α values may differ significantly for each dye.)

ΔE= h2

8m(L+α )2 (n f2−n i2) (7)

Honors: 5b. Evaluate the values determined above. Identify a trend, and consider how youα might add a second parameter that would give more uniform results. Discuss your ideas with your table group, and report your favorite approach in your lab notebook.

6. Design a custom dye: Molecules that can absorb light in the near-IR are becoming widely considered in a number of fields including solar cells and telecommunications. Let’s say that you are working for a company and are given the task to design a dye that could be classified within the “symmetrical polymethine family” that you have just been analyzing, except your dye must be able to absorb in the near-IR range. You think back to this lab and realize that the dyes you just investigated can be “tuned.” That is, the absorption spectrum can be changed by increasing and decreasing the box size. Design and draw the dye that would absorb around 1500 nm (give or take 20 nm).

Bonus. Dye 1 appears red. Explain why and predict what colors dye 2 and dye 3 should appear. (Use the class computer to pull up Figure 1 from Experiment 7, if needed.)

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Appendix

Sample Lab

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