Chem Lab Manual

Chemistry 227

General Chemistry Lab 1

Fall 2011

Lab Coordinators: Dr. Eric Sheagley [email protected] Dr. Gwen Shusterman [email protected]

Chemistry 227 General Chemistry Laboratory SYLLABUS – Fall 2010 Lab Packet: All printed material for this lab will be available on Blackboard OR may be purchased at Smart Copy (1915 SW 6th Avenue). Prelab Exercises: Prelab instructions are included in the lab packet. You should answer any questions presented and prepare for the weeks lab before your lab meeting. Pre-labs are due at the beginning of the lab period. Materials: You will need chemical splash safety goggles. These are available from the bookstore. You will need a bound carbonless copy notebook (not loose paper) for recording data. You are responsible for all laboratory equipment checked out to you. If you break glassware, you will pay the cost of the glassware, as listed on the equipment check in sheet. Dress for lab: You must wear shoes that cover your entire foot, including the heel. They should fit up near your ankle; leather is preferred but any non-porous material is okay. Short shorts and short skirts are not allowed. Your clothing must cover your torso and legs down to your knees. Grading: The laboratory is graded on a Pass/No Pass basis. An average of 75% of all points available in the lab is required to pass. Laboratory reports are due at the beginning of the lab period following completion of the experiment. Lab reports should be typed. Late reports will be docked 5 points per day late. Attendance: Attendance in this lab is mandatory. YOU MUST ATTEND ALL SCHEDULED LABORATORY MEETINGS. If you are not able to attend lab you must notify your laboratory instructor as soon as possible. For a missed lab meeting you must make up the missed lab time during the make-up. The make-up laboratory will take place during week 10 of the quarter, during the regularly scheduled lab period. In addition to completing the make up lab, students are responsible for completing the lab report for the missed lab. Data can be obtained from a lab partner or the lab instructor. The made up work should be clearly labeled and indicate the origin of the data reported. Reports are due the class meeting following the syllabus deadline. FAILURE TO DO BOTH WILL RESULT IN A NO PASS GRADE. If you miss two or more labs your grade will be a NO PASS. NOTE: If you are more than 15 minutes late to lab you will be marked late. Two late arrivals during the term will be counted as a missed lab. In addition, late students may be assigned to lab clean up duties at the conclusion of the lab period. If you are chronically late you will be given a NO PASS.

Chemistry 227 Schedule – Summer 2011 Week 1 Check-in. Announcements and registration adjustments and Lab Safety View the Lab Safety Video. A link is available on D2L. Complete the quiz, also available on D2L, before returning week 2 or you will not be able to participate in the lab Week 2 Who has the same solid that I have? Week 3 (Who has the same solid that I have?) Density: How much sugar is in Coke? Week 4 (Density report due) Alkali Metal Project - Simple Weight Loss and Gravimetric Analysis Week 5 (Alkali Metal report due) Copper Cycle Week 6 (Copper report due) Using Conductivity to find an Equivalence Point Week 7 (Conductivity report due) Atomic Emission Spectra Week 8 (Spectra report due) Beer’s Law: How much dye? Week 9 (Beer’s Law report due) All lab sections do not meet. All sections must do the Web Assignment Lab Web Assignment: Electron Density Explorations Week 10 (Electron Density due) Make-up Lab Check- out

Plagiarism: Experiments will be done in groups sharing the computer for data analysis and acquisition. You may compare data with other groups, but the content of your lab reports MUST be written individually. It will be considered an act of plagiarism if you borrow tables or graphs from another student (learning how to properly create a table or graph is an important skill, learn how to do it on your own!). You cannot paraphrase the internet, your book or any other source without the proper reference. Additionally, it will be considered an act of plagiarism if you borrow data without prior approval from your TA. There are additional resources online to help you avoid plagiarism. Please be sure to check http://www.lib.pdx.edu/instruction/survivalguide/writeandcitemain.htm or http://web.pdx.edu/~b5mg/plagweb.html, and feel free to discuss the issue with your TA or the lab coordinator. Depending on the severity of the offense(s), you will receive, at a minimum, a zero score for the report. Additionally, a report may be made to the Office of Student Affairs. Late Work: Reports are due at the beginning of the indicated laboratory period. Reports will be docked 5 points for each calendar day they are late.

Grading Criteria Unless otherwise noted, every lab is worth 90 points (including the prelab, notebook and technique) Each lab report will be graded according to the following point distribution: Prelab: 10 points Abstract: 10 points Introduction: 10 points Data: 10 points Results: 15 points Discussion: 15 points

In addition to the above points each lab meeting will have an additional 10 points assigned on the following basis: Notebook: 10 points Lab technique: 10 points Be sure to read “Keeping a Lab Notebook”. Both of these criteria will be evaluated by your TA during each lab meeting. At the end of each lab you must check out with your TA so that he or she can assess your lab notebook and verify that you have cleaned your work area. Points assigned to lab technique are assigned by the TA. The basis for assigning these points includes (but is not limited to) general lab technique and methods, safety, general mannerism in lab and cleaning up after yourself.

Labs are graded on a Pass/No Pass basis. In order to pass the lab you must turn in every lab report. More than one absence will result in a No Pass for the class. You must receive 75% or greater of all the points available to pass.

Laboratory Safety Rules and Procedures Safety Rules The guidelines below are established for your and your classmates’ personal safety. Failure to adhere to the guidelines below will result in a loss of Lab Technique points. • Personal Protective Equipment (PPE) is used to protect you from serious injuries or illnesses resulting from contact with chemical hazards in the laboratory. Spills and other accidents can occur when least expected. For this reason it is necessary to wear proper PPE. The PPE for student labs consist of goggles, gloves and clothing. Proper PPE is required for all students or they will be asked to leave the lab •Goggles – Goggles must be worn whenever any experimental work is being done in the laboratory to protect the eyes against splashes. Only indirect-vented goggles are allowed in the student labs and should be worn at all times when any chemical is being used in the lab. These are for sale in the bookstore and stockroom. You should not wear contact lenses in a chemical laboratory. Chemical vapors may become trapped behind the lenses and cause eye damage. Some chemicals may dissolve “soft” contact lenses. The most important aspect of having the goggles fit comfortably is the proper adjustment of the strap length. Adjust the strap length so that the goggles fit comfortably securely and are not too tight. If you find that your goggles tend to fog, you can pick-up anti-fog tissue from the stockroom. • Gloves – Gloves should be worn to protect the hands from chemicals. Gloves are provided through your student fees and are located in the student labs. For health and safety reasons it is important to always remove at least one glove when leaving the student laboratory, this prevents things such as door handles from getting contaminated. • Clothing – Dress appropriately for laboratory work. You must wear shoes that cover your entire foot, including the heel. They should fit up near your ankle; leather is preferred but any non-porous material is okay. Your clothing must cover your torso and legs down to your knees. Short shorts, short skirts, tank tops and halter tops are not allowed. • Eating, drinking and smoking are prohibited in the laboratory at ALL times. Wash your hands after finishing lab work and refrain from quick trips to the hall to drink or eat during lab. If you take a break, be certain to remove gloves and wash hands before ingesting food or drink. • Never work alone in the laboratory or in the absence of the instructor. • Headphones may not be worn in lab.

Safety Procedures • Know location of safety equipment; fire extinguisher, fire blanket, first aid kit, safety shower, eyewash fountain and all exits. • In case of fire or accident, call the instructor at once. • Small fires may be extinguished by wet towels. • If a person’s clothing catches fire, roll the person in the fire blanket to extinguish the flames. • In case of a chemical spill on the body or clothing, stand under the safety shower and flood the affected area with water. Remove clothing to minimize contamination with the chemical. • If evacuation of the lab is necessary, leave through any door that is safe, or not obstructed; doors that lead to other labs may be the best choice. Leave the building by the nearest exit and meet your TA on the field next to Hoffmann Hall. This would also be the meeting place in the event of an earthquake or other emergency. It is good to know the nearest exits of your lab on the first day of class. • Spilled chemicals must be cleaned up immediately. If the material is corrosive or flammable, ask the instructor for assistance. If acids or bases are spilled on the floor or bench, neutralize with sodium bicarbonate, then dilute with water. Most other chemicals can be sponged off with water. • Avoid contact with blood or bodily fluids. Notify the instructor or stockroom personnel if ANY blood is spilled in the lab so that proper clean up and disposal procedures may be followed. • If a mercury thermometer is broken, do not attempt to clean up yourself. Notify students around you, so that mercury is not spread, then notify your lab instructor or stockroom personnel. The stockroom is equipped for proper clean up and disposal of mercury.

Laboratory Procedures and Protocol General Etiquette: • Leave all equipment and work areas as you would wish to find them. • Keep your lab bench area neat and free of spilled chemicals. Your book bag, coat, etc., should be kept in the designated area at the entrance to the lab, not at your bench. • All chemical waste must be disposed of in proper containers. Proper disposal of chemicals is important for student safety and proper disposal. Putting chemicals into the wrong containers can lead to injury from unexpected chemical reactions. Mixing waste can also

make it more difficult or expensive for PSU to dispose of them. Only chemicals should go into waste jars. Waste jars for each experiment will be provided in the lab. They will be labeled specifying which contents should be placed inside. It is important that you replace the lids to the waste containers. When done with the waste jar, make sure it is placed in a secondary container. Do not put anything down the sink unless you are explicitly told to dispose of it this way. Your instructor will provide specific disposal guidelines when needed. Following these guidelines assists us in lowering the environmental impact of the labs. There are several locations for very specific waste. i. Chemical waste – these containers are ONLY for chemical waste generated in the lab. They are each specifically labeled for each lab and waste type. READ THE LABELS. ii. Contaminated paper waste – this is ONLY for paper towels used for clean-up of chemical spills. iii. Broken glass – this is ONLY for broken glassware. iv. Gloves – this is ONLY for used gloves. v. Normal trash – this is for all other trash that is not chemically contaminated, glass, or gloves. • Clean your bench and equipment Clean all your glassware- dirty glassware is harder to clean later. Wash with water and detergent scrubbing with a brush as necessary. Rinse well with water. Do not dry glassware with compressed air, as it is frequently oily. The water and gas should be turned off and your equipment drawer locked. • Clean the common areas before you leave the lab. Point deductions for the entire class will be imposed if the instructor or stockroom is not satisfied. • Return any special equipment to its proper location or the stockroom.

Handling Chemicals: Obtaining reagents: • Read the label CAREFULLY. The Chemicals are organized by experiment in secondary containment bins. Make sure the chemical name and concentration match what is required by the experiment! • Do not take the reagents to your bench. • We recommend always picking up bottles by the label. If all students do this, then any unnoticed spills when pouring will not cause possible problems for the next user. Remember to wear gloves while working with reagents. • Do not put stoppers or lids from reagents down on the lab bench. They may become contaminated. Be sure that the lids or stoppers are replaced.

• Do not place your own pipet, dropper, or spatulas into the reagent jar. Pour a small amount into a beaker and measure from that. Please pour on the conservative side to minimize waste and cost of labs. You can always go back for more. • Do not put any excess reagent back in the reagent jar. Treat it as waste and dispose of it properly. • When weighing chemicals on the balances, never weigh directly onto the weighing pan. Weigh into a weighing boat or beaker. Any spills on the balances MUST be cleaned up immediately. If you are unclear how to clean a spill, notify your instructor. The balances you are using are precision pieces of equipment and costs up to $4000. • All chemicals should be treated as potentially hazardous and toxic. Never taste a chemical or solution. When smelling a chemical, gently fan the vapors toward your nose. • Any chemicals that come in contact with your skin should be immediately washed with soap and copious amounts of water.

Laboratory Procedures • Never pipet any liquid directly by mouth! Use a rubber bulb to draw liquid into the pipet. • Never weigh hot chemicals or equipment. • When heating a test tube, always use a test tube holder and be certain never to point the open end of the test tube toward yourself or another person. • Handling glass tubing or thermometers: to insert glass tubing into a rubber stopper, lubricate the glass tubing with a drop of glycerin, hold the tubing in your hand close to the hole, and keep all glass pieces wrapped in a towel while applying gentle pressure with a twisting motion. • To prepare a dilute acid solution from concentrated acid, acid should be added slowly to water with continuous stirring. This process is strongly exothermic, and adding water to acid may result in a dangerous, explosive spattering. • Use the fume hood for all procedures that involve poisonous or objectionable gases or vapors. • Never use an open flame and flammable liquids at the same time.

Keeping a Lab Notebook In keeping a lab notebook, there are certain principles that should be followed. These boil down to being clear and complete in your entries in your lab notebook. There are also certain conventions for lab notebooks that are universally followed. High on this list are the following: Use a notebook with pre-numbered pages Record entries in ink Keep entries reasonably neat and organized Never tear pages out of your lab notebook (other than the carbonless copy pages) What Kind of Notebook Should I Use? For this class you must use a notebook with carbonless copy pages. General Guidelines • Write your name on outside front of notebook • Use black ink, fine-tipped ball-point pen (this will photocopy clearly) • At the front of the notebook, leave a few pages for a Table of Contents • Each lab should have a brief introduction and description of procedure • Generally use only the right hand page for most text • Use facing left page for working graphs, manual calculation, and working notes • Prepare data tables in advance - with columns for calculated results and notes • Working graphs done in lab notebook to monitor progress Usage and Structure You should record all your work in your lab notebook. That is the proper place for all lab planning and observations. Nothing should be recorded on odd scraps of paper, etc. The overriding principle for a lab notebook is to record in it all the pertinent information about your lab work. This boils down to clear descriptions of what you did and what you observed as a result. It is a working tool, and a reference for other researchers who might want to read your notebook and reproduce your work. (This applies to notebooks in learning laboratories: Your lab instructor may want to look at what you did in order to understand your results. This is often the case. So, it needs to be clear.) The word “clear” here is crucial. In order to be clear, data must be recorded in well-thought-out tables, clearly labeled. Descriptions of procedures must be clear and concise; to the point.

Structure for your Lab Notebooks: For each lab in this class you should have the following sections in your lab notebook: Title Purpose Procedure and Observations It is also often helpful to include a Result section Note: When preparing your notebook for lab only write on the right hand page. Title: With your lab notebook laid open, on the right hand page write down the title of the experiment, and the date. In general, you will use the right-hand page for all your writing. The left-hand page is reserved for recording scratch work. Don’t use this space until you need to. One example of how to use the left-hand page: if your work requires simple calculations using your measurements, use the left-hand page to do the calculations. If unexpected results occur later, sometimes you can look back at your scratch work and discover the error. (“Oh, I subtracted wrong! We put in 10.5 grams of copper sulfate, not 9.5 like we thought!”) Better to discover the error after the fact than never to discover it at all. Purpose: Below the title, write the purpose of the experiment in one or two sentences. This section serves to remind you and notify the reader what the experiment is about. Procedure and Observations: This next section will be labeled Procedure and Observations. As the name suggests, write down what you actually do and what you observe. This section is where you should have preprepared tables for data collection. Set up this section by dividing the page into a right and left column. In the left hand column write your procedure and in the right column next to the procedure, record observations and data or measurements. Results and Discussion: You might want to include a final section that is labeled Results and Discussion. In this section, you would describe what results you got, what conclusions you have reached, ideas for continuing work, etc.

An example of a prepared notebook follows.

Writing Style in the Lab Notebook For certain entries in your lab notebook, such as the Introduction before each experiment, you should strive to write as logically and clearly as possible. It is also a good idea to write in the third person passive voice, to get into the habit, and so that in many cases you can copy entries from your lab notebook into your reports without the need for major revisions/rewrite. However, this is a working document. It is not expected that you write perfect prose in your notebook – it is a first draft. Just do the best you can. Also, as a working document, with many entries being written while an experiment is in progress (your observations) it is understood that many entries will be brief – but still record crucial observations. Example Notebook entry: “Added 10 mL of 1M HCl – solution turned red instantly; pcpt.↓ a few secs later→ clr soln.” When written into a lab report or journal article, this would be expanded a bit and made grammatically correct. “10 mL of 1.0 M HCl were added to the clear reaction mixture. This immediately resulted in a crimson solution, and a red precipitate formed a few seconds later, leaving a clear solution.”

Adapted courtesy of Keith James

Report Guidelines For each experiment performed this term you will turn in a type written report (at the end of each lab you will find a summary of which sections to include in the report for that lab). The reports are due at the beginning of class the week following completion of the experiment. Below is a description of what should be included in each section. The sections are presented here in the order they should appear in your lab report. It is expected that you will complete each experiment and do the necessary calculations and analysis during the scheduled lab period each week. You may discuss the calculations and analysis with your lab mates. Your written lab report should be your own individual work!! The lab report sections should be complete but CONCISE. For most experiments this term, your report should be 1-2 pages long.

Writing Style You will write you reports using a formal scientific writing style. A lab report must be written in the third person, passive voice. Also, it must be in the past tense. It should not contain personal pronouns such as, “I”, “we” or “he” neither should it contain proper names of persons. Good: “50 mL of 1.0 M HCl were poured into a 125 mL Erlenmeyer flask” Bad: “I poured 50 mL of hydrochloric acid into a flask.” Also bad: “Joe Shmoe poured 50 mL of hydrochloric acid into a flask.” This is not the correct form of 3rd person. It includes Joe’s name. Also bad: “We are going to put 50 mL of acid into the flask.” Uses future tense; also, “we”. After you write your report, there is one more thing to do before you print it and hand it in: Proofread it! Read it out loud. If is doesn’t sound right, it isn’t. Fix it. Then do it again until it is right. You will enjoy writing reports more if you take pride in what you hand in.

Abstract: This is like a condensed version of your lab report. It is a stand-alone document. Abstracts are, in fact, often published separately from the articles they describe. A library search of the literature generally involves reading abstracts. This is done with the aim to identify articles that need to be read in full, and eliminate many others whose abstract makes it clear that they are not relevant to the study at hand. So, the abstract needs to be brief, but complete. There are three questions that should be answered in any good abstract 1. What did you do? 2. How did you do it? 3. What did you find?

Introduction: Here, you want to address WHY you did this experiment. Your introduction begins with a statement of the purpose of the experiment. You already did this in the Abstract, but remember the Abstract is a stand-alone document. What you said there doesn’t count now; you will have to repeat yourself a bit. Next, provide any relevant background, to put the experiment into context. Include any key concepts needed by the reader to understand your experiment. This means that your Introduction will often include some explanation of the theory behind the experiment. If general equations are important to the explanation, include them here.

Data: This is section is where your observations and experimental data belong. (In published papers, a data section is usually not included, but, this is a class so this section will be included.) When ever possible, data should be presented in tables. Any tables should include descriptive column headings, including units. Also, include gridlines for clarity. Tables should not be divided across page boundaries. In this section you would also include observations and descriptions of other pertinent events. This section is not where the calculations, interpretation and discussion of your results belong.

Results: The results section is where you should show sample calculations and report all of your results. For every type of calculation you should show one sample calculation. Each calculation should have a descriptive title, i.e. “Calculating the density of Coca-Cola”. The calculation section should be annotated. The annotation is provided to describe why each calculation is useful and relevant to the lab activity. The description should not be any longer than two or three sentences and should help you describe your results in your discussion section. Sample calculations may be written by hand attached as an appendix to your report. The results of all calculations should be summarized in a table where appropriate.

Discussion: In this section, you will discuss interpretations of the experimental results. This is where you get to present your thinking process. For any labs that have questions to answer, this is also where the answers get written up. The discussion is one of the most important parts of the lab report! It is your chance to show that you UNDERSTAND what you did in the lab. Explaining how you reached your results from the raw data can go a low way. This DOES NOT mean to include detailed procedures or to explain your calculations is words. It DOES mean that a general description of the experiment can be useful in explaining your results. In this section you also discuss error analysis. This does not necessarily mean trying to explain what went wrong. (Maybe nothing did go wrong!) It means discussing the limitations of your experiment. For example, if you are doing calorimetry in a coffee cup, and the cup feels warm to your hand, it means that some heat is escaping. Also, if you are reading a 5 degree temperature change with a thermometer that you can only read to the nearest 0.5 degree, there is a significant

uncertainty in the exact magnitude of the temperature change. You could easily have a 10% error, or even more, and this needs to be taken into account. It at least needs to be mentioned, to show that you were aware of the issue. This is a limitation of the apparatus, not an error on your part. And, yes, if something did go wrong (your lab partner forgot to write down the exact molarity of your reagent), then that should go here, too, along with an explanation of how you attempted to correct for the error. (In this case, you may have had to re-do the experiment.)

Adapted courtesy of Keith James.

Example Lab Report Following is an example of a lab report prepared according to the previous report guidelines. Sample calculations can be written on a separate paper and attached to the report.

Calibration of a 10 ml Volumetric Pipette a 10 Abstract: A 10 ml volumetric pipette was calibrated by determining the mass of water delivered by the pipette. The mass of water was then converted to volume using the density of water. The volume of the pipette was determined to be 9.98 +/- 0.02 ml when the mass of water was determined on a pan balance and 9.998 +/- 0.002 ml when determined with an analytical balance.

Introduction: A volumetric pipette is designed to deliver a stated volume of liquid, however, the actual amount of liquid any individual pipette delivers may vary slightly from this ideal stated volume. In order to determine the actual volume an individual pipette delivers, it will be calibrated. In this case, calibration refers to the comparison of the actual amount of liquid delivered by the pipette to the standard value of the pipette (10 ml). Because delivered volume is being calculated, another measurable quantity must be used to verify the volume delivered by the pipette. In this case, the relationship between mass and volume (density) will be used. Mass is an easily measurable quantity that can be determined with a high degree of accuracy due to the availability of electronic balances. Mass can then be converted to volume by the use of density. D (density) = m (mass)/ V (volume) Because the density of water at a variety of temperatures is readily available, here, water will be used to calibrate the volume of the pipette.

Data: Table 1: Mass Determined by Pan Balance (+/- 0.01g) Run # mass beaker (g) mass beaker + water (g) 1 27.88 37.83 2 27.88 37.82 3 27.88 37.86 Table 2: Mass Determined by Analytical Balance (+/0.0001g) Run # mass beaker (g) mass beaker + water (g) t(transfer) 1 27.2349 36.5618 2:29:00 2 27.2348 36.7813 2:32:00 3 27.2335 36.8251 2:41:20

t(weigh) 2:30:30 2:33:20 2:42:30

Diameter of beaker: 3.9 cm +/- 0.1 cm Mass of water evaporated in 60 seconds: 0.0016g +/- 0.0002g Temp of water: 20.5 ºC +/- 0.2 ºC Density of water: 0.9980 g/ml

Results: Calculation of the volume of water: In this calculation, the average mass of water for the three trials as determined by the pan balance was divided by the know density of water at 20.5 ºC. The data are summarized in Table 3. Volume = 9.96 g / 0.9980 g/mL = 9.98 mL Table 3: Mass of Water as Determined by a Pan Balance (+/- 0.01g): Run # mass water weighed (g) Volume water (ml) 1 9.95 2 9.94 3 9.98 Average 9.96 9.98 Error +/- 0.02 +/- 0.02 Calculation for the mass evaporated: To correct for evaporation of water in the time it takes to measure the mass of the water delivered by the volumetric pipet, the mass of water that evaporated was estimated. The rate of evaporation of water in the 50 mL beaker in 60 seconds: 0.0016g. The data are summarized in Table 4. Mass evaporated = rate of evaporation x time of evaporation = (0.0016 g/ 60 s) x 90 s = 0.0024 g

Calculation of the mass transferred: The mass of water initially transferred was the sum of the mass of water evaporated and the mass of water present at the time of weighing. The data are summarized in Table 4. Mass transferred = mass water weighed + mass (transferred) = 9.9769 + 0.0024 = 9.9793 g Table 4: Mass of Water as Determined by an Analytical Balance (+/-0.0001 g) Run # mass water weighed (g) t(evap) (s) mass (evap) (g) mass (transferred) (g) Volume (ml) 1 9.9769 90 0.0024 9.9793 2 9.9735 80 0.0021 9.9757 3 9.9756 70 0.0019 9.9775 Average 9.978 9.998 Error +/- 0.002 +/- 0.002

Discussion: The mass of water delivered by a 10 ml volumetric pipette was determined on both a pan balance and an analytical balance. The mass of water was then converted to volume using the density of water. In the case of the analytical balance, the rate of evaporation of water (which is a systematic error) was taken into consideration. In this case the mass of water that evaporated from the time the water was delivered to the beaker to the time of weighing was added to the weighed mass of water delivered by the pipette. This correction was not necessary when the pan balance was used since the accuracy of the pan balance is +/- 0.01 g and the evaporation rate of water under experimental conditions was found to be 2.7 x 10-5 g/s. The use of the analytical balance increased both the precision and the accuracy of the calculated volume of the pipette (9.98 +/- 0.02 ml with the pan balance and 9.998 +/- 0.002 ml with the analytical balance). The improvement in the results can easily be seen by the percent error which was calculated to be 0.2 % with the pan balance and 0.02 % with the analytical balance. The largest source of error in this experiment most likely came from the difficulty in accurately filling the pipette to the mark with water which introduced random error into the experiment.

Chemistry 227 Pre-Lab: Who has the same solid that I have? Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. Some of the chemicals you will use this year are very hazardous. One way of determining the risk of using various compounds is to read the Material Safety Data Sheet (MSDS). These provide a wealth of information regarding the safety risks of each compound. Do a web search with the key words “MSDS and Lead Nitrate”. Read through the MSDS sheet and determine which steps need to be taken in case of accidental skin exposure (your most common risk in the lab). 2. Look up the MSDS for both Hydrochloric Acid and Sodium Hydroxide. What steps need to be taken if there is skin exposure? 3. In this lab, there are many possible unknown compounds, including ammonium iodide, sodium acetate, Silver nitrate, calcium carbonate, lithium carbonate, aluminum chloride, potassium iodide. Pick one of the listed compounds and read the MSDS on it. Additionally, search the internet for two interesting factoids regarding the chemical compound you choose.

Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted format (be sure to include space for observations).

At the start of your lab, remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA.

"Who has the same solid that I have?" Science is generally a cooperative collaborative affair; most discoveries are not made by just one person. It is important for scientists to be able to communicate their data and be prepared to share data or samples. In this experiment you will be trying to determine who else in your lab section has the same unknown compound as you. This will be done while learning some very basic techniques that are used for the analysis of some compounds. As you will see as you progress throughout the year in chemistry, compounds are can be classified in many different ways (ionic or molecular, acidic or basic, metals or nonmetals…..). Think of how you might classify water on the basis of easily observable properties. We know that water is clear, colorless, freezes at 0°C, boils at 100°C, dissolves most salts, has a density of 1 g/ml and is composed entirely of hydrogen and oxygen atoms in a definite ratio (this list could go on and on). To enable of our classifications, it is important to be able to accurately determine and compare the chemical and physical properties of compounds. You will be given a sample of an inorganic solid and will determine your samples properties such as: the solids relative solubility, its relative melting point, the electrical conductivity of the substance and its aqueous solution, the acidity/basicity of the compound’s aqueous solution, its appearance in a flame and its reactivity. Your goal is to identify other students in class who have the same compound that you have. Comparisons of different samples may be made doing a side by side analysis utilizing the various techniques. Guided by your TA, your lab section will determine a method for sharing or reporting your observations. You should identify all people in the lab section having the same substance and run some confirmatory tests to verify. Below you will find a list of properties that you will investigate during this lab exercise: 1) Melting points: A substance’s melting point temperature will depend on the bonding type or intramolecular forces in the sample. As a result, some compounds may have melting points much greater than 200°C, while others may have much lower melting points. Upon further heating some compounds may decompose into simpler compounds or burn. 2) Conductivity of aqueous solutions: When dissolved in water, some compounds dissociate in to ions. These dissolved ions allow the electros to move through the solution and thus conduct electricity. 3) Crystalline or amorphic: As a result of the types of bonds in the compound, a substance may form very regularly shaped crystals. Others are less able to form regular patterns so their solids are less geometric. Additionally, crystalline compounds are hard and brittle because the ions are locked tightly into place by their electronic interactions. As a result, it’s difficult to move these ions apart, and when they do move apart, the whole crystal typically breaks. 4) Flame test: Atoms emit visible characteristic colors when excited by an energy source. This light is a characteristic signature of the element, which is a consequence of the electronic structure of the element. These emitted colors can be used to identify the atomic composition of a substance. For instance, potassium produces a violet color while lithium will emit a vibrant red.

5) Acidic, basic or neutral aqueous solutions: Because of the nature of their composition, certain substances when dissolved in water will make the solution acidic or basic when they dissociate into ions. Some ions in solution have the ability to act as acids while others act as bases. When a substance is dissolved in water, these properties can easily be tested using pH paper. 6) Reactivities: Each compound has a characteristic reactivity that may or may not be easily elucidated. In mixing an aqueous solution of the unknown with an aqueous solution containing another compound reactivity patterns may become visible. Reactions are usually visualized by looking for the formation of a solid, gas or change in color. This lab is based upon the journal article "Who Has the Same Substance that I Have?": A Blueprint for Collaborative Learning Activities, Brian P. Coppola, Richard G. Lawton, Journal of Chemical Education 1995 72 (12), 1120 and Identification of Ionic and Molecular Compounds, http://tinyurl.com/3jf6oq6, n.p. n.d. Web. 24 Aug. 2011

MATERIALS gas burner, striker, well plate, magnifying glass, conductivity meter (all available in the lab room).

PROCEDURES 1. Physical characteristics a. Obtain a small (pea sized) sample of your assigned unknown. b. Using a magnifying glass examine the sample and record your observations 2. Determine conductivity in the solid state a. Using the sample obtained previously, test for electrical conductivity using the conductivity meter supplied by touching the prongs to the sample. [Be sure the prongs are dry!] b. Record your observations 3. Determine the solubility of each unknown a. Add ½ of your “pea sized” sample of the unknown to a small test tube. Add 2 cm or approximately 1 fingers width of deionized water to the test tube. b. Mix with your microspatula. c. Record your observations as S = soluble, IN = insoluble (soluble means that a clear solution has formed, insoluble means that the sample is cloudy or that there is undissolved solid left in the test tube) d. Do not discard the contents of the test tube! 4. Determine the conductivity the solution made in step 3 a. As a control, use the conductivity meter to monitor the conductivity of deionized water b. Test your unknown aqueous solution for conductivity. Record your observations c. Do not discard the contents of the microplate!

5. Flame test a. Dip a Nichrome wire into a beaker of water and place it in the hot part (the blue inner cone) of the flame until it glows red-orange. Do at least 5 times or until the color of the flame is consistent. b. Dip the wire into your unknown and place it in the hot part of the flame substance. Observe the color of the flame when the liquid is evaporating. Note: some substances will not show a positive flame test. c. Record your results 6. Determine if the aqueous solution is acidic basic or neutral. a. Dip your micro-spatula in your solution and wipe it on a piece of pH paper. The paper is normally orange. It will turn red if the solution is acidic or blue if it is basic b. Record your observations 7. Determine the reactivity or your unknown. a. Using a dropper, equally divide your solution amongst three wells in your well plate. b. Add 5 drops of 1M hydrochloric acid (HCl) to the first well. (The symbol M represents molarity, a unit of concentration. The greater the molarity, the greater the concentration). c. Add 5 drops of 0.1 M lead(II) nitrate (Pb(NO3)2) to the second well d. Add 5 drops of 1.0 M sodium hydroxide (NaOH) to the third well. Waft a water moistened piece of pH paper over the third well to monitor if a gas (ammonia) is produced. Note: aqueous solutions of ammonia are basic, which is why the pH paper would turn blue in the presence of ammonia gas. e. Record your observations. f. Clean your test tube, well plate and dropper as directed by your TA 8. Determine the relative melting point for each unknown compound a. Carefully light the gas burner b. Using a scoopula, obtain a small portion (1/2 of the original amout from step 1) c. Carefully heat the sample on the edge of the scoop. d. Approximate the distance from the flame and monitor the time required to melt. (Any substance that will melt under normal lab conditions will do so quickly don’t heat any substance longer than 20 seconds!) e. Record your observations f. Clean the scoop as directed by your TA 9. Clean up. All remaining solutions and solids should be placed in the properly labeled waste jar. Your lab area should be wiped clean and all glassware and equipment should be placed in your lab drawer.

DATA Guided by your TA, you will determine a table to report your results and observations.

For the TA The most important practical aspect of setting up this laboratory is to ensure that the identification is based on the experimental data that are collected by the students. Please discuss contamination and how not to contaminate the stock solutions and unknowns. Possible unknowns include: ammonium iodide, sodium acetate, silver nitrate, calcium carbonate, sugar, lithium carbonate, aluminum chloride, citric acid, potassium iodide

Who has the same solid that I have? Lab Report: Your report for this lab should include the following sections: Abstract: Your first abstract for this lab should be written as part of a post-lab discussion led by your TA Introduction: Address WHY you did this experiment Data: Include your data table Results: A results section is not necessary for this report. Discussion: Write a paragraph explaining the results of this experiment, discuss the following points:  Give a brief summary of the results  Discuss how you matched your unknown sample with the other(s) in your lab section.

Chemistry 227 Pre-Lab: How much sugar is in a can of coke? Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. A solution has a mass of 109.5 g and a volume of 100.0 ml, what is the density of the solution? 2. An object has a density of 0.25 g/ml, would you expect this object to float on water? 3. How many milliliters are in 1 pint? 4. What is the purpose of constructing a calibration curve? 5. The Coke was left open for several days. Why do you think this was an important step?

Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted format (be sure to include space for observations) and preparing any tables necessary for data collection.

At the start of your lab, remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA.

How Much Sugar is in a Can of Coke? GOALS: 1. Determine the amount of sugar (in grams) in a can of coke 2. Learn how to make solutions quantitatively 3. Learn how to make and use a calibration curve

INTRODUCTION: If you were to measure out identical volumes of Coke and diet Coke, you would find that the two liquids have different masses. This difference in the mass of the two liquids is best discussed by looking at the mass per unit volume (or density) of the two liquids. Density is a convenient quantity because it is independent of the volume used. The density of two different solutions can be compared without needing to have the same volume of the two solutions. In other words, density is an intensive property. Intensive properties are independent of the amount of substance. M Density = Mass  Volume D V When comparing Coke and diet Coke, it is found that Coke is more dense than its sugar free relative. To understand why, a molecular view of the two substances must be examined. Obviously, the main difference between the two is the presence of dissolved sugar in Coke which is absent in diet Coke. The presence of the sugar makes Coke more dense than diet Coke. To a first approximation, Coke can be represented as a solution of sugar dissolved in water. As the amount of sugar dissolved in a given volume of water increases, so does the density of the resulting solution. This provides an ideal means by which the mass of sugar in Coke can be determined. The relationship between the amount of dissolved sugar and density of the sugar water solutions will be determined through the use of a calibration curve. Calibration curves are constructed using known quantities. Calibration curves allow you to determine the content of an unknown by comparing it to known values. In this case, you will prepare solutions of known volume with a known amount of dissolved sugar. After calculating the density of each solution, you will prepare a graph of dissolved sugar vs. density. You will then graphically determine the relationship between the two quantities. Once the relationship between density and sugar content is determined, you will calculate the density of sugar and use this relationship to determine the amount of sugar in Coke.

SCIENTIFIC GRAPHS: This experiment will also serve to introduce you to scientific graphing. Here, we will introduce what must be included in any scientific graph. Whenever you are asked to produce a graph from laboratory data all of the following criteria must be met. 1. All graphs must have a title 2. Axes must be labeled with a name and units 3. The graphed data must take up the full space of the graph 4. The minimum size of a graph should be ½ of a standard sheet of paper 5. When showing a fit line to the data, the line and equation (if applicable) should be shown on the graph 6. The independent variable is the x-axis and the dependent variable is the y-axis Figure 1 shows an example of an acceptable scientific graph of raw data. Figure 2 demonstrates the proper way to represent a linear fit on a graph. Figure 1: The Relationship Between Mass and Volume for Water 60 50

Mass (g)

40 30 20 10 0 0

10

20

30

40

50

60

Volume (ml)

Figure 2: The Relationship Between Mass and Volume for Water 60 y = 1.0015x + 0.009

50

Mass (g)

40 30 20 10 0 0

10

20

30 Volume (ml)

40

50

60

Graphing Using Microsoft Excel: An excellent tutorial on graphing with excel can be found at the following website: http://www.ncsu.edu/labwrite/res/gt/gt-menu.html This is a list of the basic steps necessary to graph data and do linear regressions with excel: Basic Graphing: 1. With the program open, enter the data to be graphed in the cells. Enter x data in one column followed by y data in an adjacent column. 2. Click and drag the mouse to highlight all the data to be graphed. 3. Click on the chart wizard icon

4. Choose XY (Scatter) for the chart type and the unconnected points icon for the Chart sub-type 5. Click next. A preview of your chart will appear. If it appears correctly, click next. 6. Enter a chart title and the axis labels and click finish 7. With the chart selected you can also access the title and axis labels by selecting ‘Chart’ then ‘chart options’ from the drop down menu

Adding a Linear Trendline to a Graph: 1. With the graph selected, select ‘Chart’ then ‘add trendline’ from the drop down menu. 2. Select ‘linear’ as your regression type 3. Select the ‘options’ tab in the popup window 4. Select the ‘display equation on chart’ button and click ok

PROCEDURE: Calibration Curve: You will make up five sugar water solutions. Each solution should have a different amount of dissolved sugar covering a range from about 1 – 8 g of sugar per 50 ml of solution volume. To make the solutions in a quantitative manner, they must be prepared in volumetric flasks. Volumetric flasks are designed to accurately contain a specific volume. Volumetric flasks are marked with a marked fill line. When filled to the marked line the flask accurately holds the stated volume. When filling volumetric flasks, it is best to bring the fluid to the line carefully by using a wash bottle or eyedropper to assure that the flask is not overfilled. Errors most likely arise from students not taking caution to ensure that fluid is only added below the line and not poured down the side of the flask. Water on the sides of the flask above the fill line will result in solutions that are more dilute than intended.

To accurately know the mass of sugar dissolved per 50 ml of each of your five know solutions, follow the following series of steps: 1. Weigh the empty flask with stopper and record the mass. 2. Weigh out the desired mass of sugar in a weigh boat. (This mass does not need to be recorded.) 3. Add the sugar to the flask. Weigh the flask containing sugar with the stopper and record the mass. 4. Add water to the flask until it is approximately half way to the fill line. Swirl the flask to dissolve the sugar. Do not shake the flask. Shaking or violent swirling will cause water to collect above the fill line and result in an inaccurate solution. 5. Once the sugar has completely dissolved, add water carefully to the fill line. Cap the bottle and shake (or invert a few times) to ensure that the solution is thoroughly mixed. 6. Weigh the capped flask containing the solution and record the mass. The mass of sugar used for each solution is found by subtracting the mass of the empty stoppered flask from the mass of the stoppered flask containing sugar. The mass of the solution is found by subtracting the mass of the empty stoppered flask from the mass of the stoppered flask containing the solution.

Below is an example of an acceptable table to present the data from this experiment. In this first experiment, a table is supplied for you. In later experiments, you will be expected to produce your own data and results tables. Mass of Empty Flask + Stopper

Mass of Flask + Stopper With sugar

Mass of Flask + Stopper With solution

Mass of Sugar

Mass of Solution

Density of Solution

Grams of sugar per ml

Flask 1 Flask 2 Flask 3 Flask 4 Flask 5 Using the data in the above table, construct a graph of density (y) vs. mass of sugar per ml of solution (x) and fit the data to a linear relationship. Report the equation for the line on the graph. Graphs should be prepared by computer. Your TA will assist you with this. This graph represents the relationship between the density of the sugar water solution (something that can be measured) and the amount of dissolved sugar in the solution (something that cannot be measured directly).

Determine the amount of sugar in a can of Coke: Weigh and record the mass of a dry, clean 50 ml volumetric flask before carefully filling the flask to the fill line with flat Coke. Weigh and record the mass of the flask containing Coke. Determine the density of the coke. Put the used coke in the provided waste jar.

RESULTS: When a linear relationship exists between two quantities (density and amount of sugar) it is only necessary to measure one of the quantities (density) and know the relationship (found from your calibration curve) before the other quantity (amount of sugar) can be found. By finding the density of coke on your calibration curve, you can read across to the y axis and determine the amount of sugar dissolved in each ml of Coke. In order to find the mass of sugar in one can of Coke, you will need to consider the volume of a can of Coke (12 ounces). One liter contains 33.8 fluid ounces. Calculate the percent error in your determined value.

Supplemental information for the density of coke lab. In this lab one should find a liner relationship between the density of a sugar solution and the number of grams of sugar in each mL of solution. Some people are confused by the fact that both axes of the graph have units of g/mL. One thing to keep in mind is that one is density and the other is “a quantity of sugar in every mL of solution”. So if I wanted to know how many grams were in 12 ounces of coke, all I would need to do is to determine the density of coke. With density, I could do one of two things. First, I could interpolate the graph to estimate the

mass in each mL of coke or second, I could substitute the density of coke into the linear equation that describes the relationship of the density of sugar solutions to the mass in each mL of solution. Both methods are fine but the second is a much cleaner method. I encourage you to determine how to get excel create a trend line and supply the linear equation.

Density of Blah Blah Blah Solutions

1.4 1.2 1 (g/mL)

Density of Blah Blah Blah Solution

1.6

y = 0.05x + 1

0.8 0.6 0.4 0.2 0 0

2

4

6

8

10

Grams of Blah Blah Blah in Each Milliliter of Solution (g/mL)

In the above graph, I plotted the density of blah solutions vs. the mass of blah in each milliliter of solution (I used blah so there would be no confusion between what I am showing here and what you should do with your data). I will show both methods using my fictitious data as to how one should determine the mass of blah in 355 mL of a drink containing blah. Using the linear equation for my data provided by excel (y = 0.05x + 1). y represents the density of the solution and x represents the grams of blah in each mL. If I empirically determine the density of my drink to be 1.37 g/mL, that value can be substituted in to the above equation to give 1.37 = 0.05x + 1. Solve for x and one calculates there are 7.4 grams in each milliliter. To get the total mass in the drink requires multiplying the volume of drink by the 7.4 g/mL, a value equaling 2600 g of blah. The second method of interpolation requires nicely drawing lines perpendicular to one another. The first line (red) is drawn from the y-axis at the measured density of the drink to where it crosses the data trend line. The second (blue) goes from that line intercept to the x-axis (the number of grams in each mL of the drink). From that one may estimate that there are just about 7.4 g of blah in each mL of drink.

Sugar in Coke Lab Report: Your report for this lab should include the following sections: Abstract: Your first abstract for this lab should be written as part of a post-lab discussion led by your TA Introduction: Begin with a statement of the purpose of the experiment Provide any relevant background and key concepts, equations and an explanation of the importance of calibration curves Data: Include your data table Results: Include a copy of your calibration curve Report the amount of sugar in 1 mL and 1 can of coke Calculate the percent error for your amount of sugar in a can of coke Be sure to attach hand written sample calculations to the back of your report Discussion: Discuss the experiment and any possible sources of error

Chemistry 227 Pre-lab: Which Alkali Metal Project Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. What is the law of conservation of mass? 2. How many moles of CO2 can be produced by the reaction of 1.53 g of lithium carbonate? (assume sufficient HCl for a complete reaction) 3. In the synthesis of barium carbonate from an alkali metal carbonate a student generated 3.723 g of barium carbonate from 2.001 g of their alkali metal carbonate. The reactants were M2CO3 and barium chloride. Write the balanced chemical reaction for this synthesis. 4. How many moles of barium carbonate were produced? 5. How many moles of alkali metal carbonate were reacted? 6. What is the molar mass of the alkali metal carbonate? Hint: remember the units of molar mass are g/mol. 7. The chemical formula for the alkali metal carbonate is M2CO3, what is the molar mass of M? 8. Look up the MSDS for barium chloride. Is the substance toxic? What steps need to be taken if there is skin exposure or accidental ingestion?

Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted format (be sure to include space for observations) and preparing any tables necessary for data collection.

At the start of your lab, remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA.

Which Alkali Metal Carbonate? The Problem In a search for a good cleaning formulation (as in laundry detergent or a degreaser for metal parts) you may find one of the alkali metal carbonates useful. In natural deposits, these carbonates may occur as deposits of a single alkali metal carbonate (such as, lithium carbonate) or as several of the alkali metal carbonates codeposited. You are an analytical chemist and have received a sample of a pure alkali metal carbonate from a newly-discovered deposit. Your task is to determine which alkali metal carbonate composes the sample. You will determine an experimental atomic weight for the alkali metal in the carbonate you have and, from that value, will determine which alkali metal is present. You will also evaluate sources of error as you compare your experimental values with the expected value for the atomic weight of the alkali metal. You will use three different methods to ensure “triangulation” and confidence in your results. Based on your experience you will be able to recommend which procedure you would use if you had time and resources for only one technique. Although you will work independently you will form a research team with two or three other people so you can compare procedures and results. Collaboration is usually the case in working laboratories. At the end of the project you will prepare a report giving your experimental results. This will include the identification of your alkali metal carbonate, error discussion, and a recommendation about which single method you would recommend to the lab.

Gravimetric analysis 1 In this experiment you will perform a synthesis and use reaction stoichiometry to identify your unknown alkali metal carbonate. The reaction involves your aqueous carbonate reacting with barium chloride (BaCl2). The product is an insoluble barium carbonate. You will isolate it and weigh it.

PROCEDURE: Be sure to discard all waste as directed by your TA 1. Add 1.00 g sample of your carbonate to a 250 mL beaker. Add 100 mL of water and stir until the carbonate is completely dissolved. 2. To precipitate the barium carbonate, add 20.0 mL of 1.0 mol/L BaCl2 to the sample. Stir. CAUTION: You should use gloves to handle the barium compounds. 3. Heat the BaCO3 formed to “digest” the precipitate (cause to form larger aggregates). This involves boiling the solution for 5 minutes with little agitation. 4. Weigh two pieces of filter paper. Be sure to record the mass. 5. Filter the barium carbonate using 2 pieces of filter paper in a Buchner funnel using vacuum filtration as demonstrated by your TA. 6. Wash the precipitate with water. 7. Wash the precipitate with ethanol. 8. Carefully remove the solid and filter paper and place your product on a watch glass and allow to dry until the end of the lab period. CAUTION: The filtrate (solution left after filtration to isolate barium carbonate) contains excess Ba2+. Dispose of in the proper waste container (see lab instructor for proper procedure). DO NOT DUMP DOWN SINK. 9. Weigh the dry solid at the end of the lab. Analysis Use the following questions to lead you to the identity of M: Determine the mass of barium carbonate produced? How many moles of barium carbonate? Use stoichiometry to determine how many moles of M2CO3 were present in the 1.00 g sample you started with. Determine the molar mass of the unknown metal carbonate and the atomic weight of M. Where are the errors most likely to enter into the experiment?

Simple weight loss In this experiment you will make use of the principle of the conservation of mass to determine the identity of your alkali metal carbonate. The metal carbonate will react with the acid added to produce carbon dioxide (CO2) gas. Applying the law of conservation of mass you can determine the mass of CO2 evolved. This will allow you to use the chemical equation to determine the atomic mass of the alkali metal. The chemical reaction is M2CO3

+

2 HCl

----->

CO2 + 2 MCl + H2O

Use the outlined procedure below to help you design your experiment and set up your data table. Answer any questions you encounter along the way. You should perform the procedure three times and should obtain a deviation of less than 10%.

PROCEDURE: 1. Place 1.00 g of your unknown in a pre-weighed 250 mL beaker. 2. You will need to add about 40 mL of 1 M HCl; you should determine the actual mass of HCl added. Pour the HCl slowly onto the unknown metal carbonate. 3. Measure the mass of the beaker after the reaction. Analysis: Use the following questions to lead you to the identity of M: Determine the mass of CO2 produced. How many moles of CO2 were produced? How many moles of carbonate produced the CO2? How many moles of M were present in the 1.00 g of carbonate? Determine the molar mass of the unknown metal carbonate and the atomic weight of M. Where are the errors most likely to enter into the experiment?

Experiment adapted from: Dudek, E. P. J. Chem. Educ. 1991, 68,948.

Alkali Metal Lab Report: Your report for this lab should include the following sections: Abstract: The second abstract written in this class may be a group effort within your research team of 2-3 students Introduction: Begin with a statement of the purpose of the experiment Provide any relevant background and key concepts, useful chemical equations and an explanation of the techniques used Data: Report the number of your unknown Include a data table for your 3 simple weight loss trials Include the data from your gravimetric analysis Results: Include a results table that summarizes your results from both methods (be sure this includes all trials of each method) Report your findings of the molar mass of your metal carbonate, the molar mass of the metal and the identity of the metal Be sure to attach hand written sample calculations to the back of your report Discussion: Discuss the experiment and any possible sources of error In addition, answer the following question as part of your discussion: 1. Which method do you feel was the most successful (be sure to support your answer with an explanation)

Chemistry 227 Pre-Lab: Copper Cycle Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. Balance the equations for reactions 1 and 2 of the lab. 2. Write the net ionic equation for reaction 2 of the lab. 3. List FOUR types of aqueous reactions. 4. Look up the MSDS for nitrogen dioxide gas. What steps need to be taken to avoid exposure? 5. This is one of the most dangerous labs for the term because of the risk of exposure to some many dangerous substances. Concentrated nitric acid is a particularly nasty solution; what happens when skin comes in contact with it. What safety precautions need to be taken to avoid exposure?

Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted format (be sure to include space for observations) and preparing any tables necessary for data collection.

At the start of your lab, remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA.

A Cycle of Copper Reactions GOALS: 1. 2. 3. 4.

Cycle solid copper through a series of aqueous reactions Learn about and identify different types of aqueous reactions Practice writing ionic and net ionic reactions Calculate percent recovered copper

INTRODUCTION: This experiment will cycle elemental copper through a series of five reactions summarized below:

Cu(NO3)2

NaOH

Cu(OH)

HNO3 heat

Cu Zn, HCl CuSO4

H2SO

CuO

The cycle will both begin and end with pure elemental copper. At different stages of the cycle, copper will be present in different forms. At times copper will be present in solid compounds and other times in ionic form. Each chemical change that copper undergoes is observable as a change in the physical properties of the solution (or precipitate). As you perform each reaction be certain to observe and record all physical changes. At this point in the class, you should have been introduced to three different types of aqueous reactions precipitation reactions, acid-base reactions and oxidation-reduction (or redox) reactions. In precipitation reactions, soluble cations and anions combine to form an insoluble compound. In acid-base reactions, an acid and base react to produce water and a salt. Redox reactions involve the transfer of electrons. As you go through the series of reactions you should be able to classify each reaction (with the exception of reaction 3) as one of the three above described types of aqueous reactions.

Reaction 1: The first reaction takes place according to the following unbalanced chemical equation: HNO3 (aq) + Cu (s)  Cu(NO3)2 (aq) + H2O (l) + NO2 (g) In this first reaction, elemental copper is reacted with concentrated nitric acid. The result of this reaction changes copper from its elemental state to an aqueous, ionic state (Cu2+). Reaction 2: The second reaction then converts the aqueous Cu2+ into the solid copper II hydroxide (Cu(OH)2) through reaction with sodium hydroxide according to the following unbalanced chemical equation: Cu(NO3)2 (aq) + NaOH (aq) Cu(OH)2 (s) + NaNO3 (aq) Reaction 3: The third reaction takes advantage of the fact that Cu(OH)2 is thermally unstable. When heated, Cu(OH)2 decomposes (breaks down into smaller substances) into copper II oxide and water according to the following equation. Cu(OH)2 (s) + heat  CuO (s) + H2O (l) Reaction 4: When the solid CuO is reacted with sulfuric acid, the copper is returned to solution as an ion (Cu2+) according to the following equation. CuO (s) + H2SO4 (aq) CuSO4 (aq) + H2O (l) Reaction 5: The cycle of reactions is completed with his reaction where elemental copper is regenerated according to the following equation. CuSO4 (aq) + Zn (s) ZnSO4 (aq) + Cu (s) Here, zinc and copper exchange states in acidic solution. Hydrochloric acid is used to dissolve any excess zinc. The solid copper can then be collected, dried and weighed.

PROCEDURE: Be sure to discard all waste as directed by your TA Reaction 1: Caution: Concentrated nitric acid is hazardous. Avoid getting it on your skin or clothing. If you do get any on your skin or clothing, wash it off immediately with water. Do not breathe vapors. Weigh out about 0.5 g of copper. Be sure to record the actual amount used to the nearest milligram. Place the copper at the bottom of a 250 mL Erlenmeyer flask. In a graduated cylinder, carefully measure out 10.0 mL of concentrated nitric acid. DO THIS NEXT STEP IN THE HOOD!! NO2 gas is toxic! In the fume hood, add the nitric acid to the flask containing the copper. The nitric acid should completely cover the copper. Be sure to record all observations. Remaining in the hood, swirl the flask until all the copper has dissolved. Once the reaction is complete and the gas has dissipated, add DI water to the flask until it is about half full. Once you are sure that all the gas has been removed in the fume hood, you may return to your workbench.

Reaction 2: While stirring with a glass rod, slowly add 30 mL of 6.0 M NaOH to the flask. Be sure to record all observations.

Reaction 3: With stirring, slowly heat the flask on a hot plate until the solution just begins to boil. At this point you should notice that the blue Cu(OH)2 has been converted to the black CuO. If the conversion does appear complete (all of the blue Cu(OH)2 has disappeared), heat the flask a little longer. Do not let the solution boil vigorously. Remove the flask from the hot plate once the conversion is complete and allow the CuO to settle. In a clean beaker, heat ~ 200 ml of distilled water. Once the CuO has settled, carefully decant the supernatant liquid. Add ~50 mL of the hot water to the flask, allow the CuO to settle again and decant the water. Repeat the wash with a second aliquot of the hot DI water. Finally, Remove as much of the water as possible without losing the desired product (CuO). Be sure to record all observations.

Reaction 4: Add 15 ml of 6.0 M H2SO4 to the flask. All of the black CuO should be gone at this point. Be sure to record all observations.

Reaction 5: Caution: Concentrated hydrochloric acid is hazardous. Avoid getting it on your skin or clothing. If you do get any on your skin or clothing, wash it off immediately with water. Do not breathe vapors. DO THIS NEXT STEP IN THE HOOD. Hydrogen gas is generated which is extremely flammable. There should be no open flame in the room. Add all at once 2.0 g of 30-mesh zinc stirring until the supernatant liquid is colorless. When the evolution of H2 gas has become slow, decant the supernatant liquid. Remain in the hood and add 5 ml of distilled water followed by 10 ml of concentrated hydrochloric acid. The hydrochloric acid removes any excess zinc according to the following equation. Zn (s) + 2 HCl (aq) ZnCl2 (aq) + H2 (g) If hydrogen gas emission stops before all the solid zinc has been removed, more acid can be added. Once the evolution of hydrogen gas has become very slow, the flask may be returned to the workbench. Warm (hot plate), but do not boil, the solution. Once the hydrogen gas evolution has completely stopped, remove the flask from heat and decant the liquid. Transfer the solid copper to a clean beaker. Using a wash bottle to wash the copper into the dish can facilitate the transfer. Wash the copper at least twice with about 5 ml of distilled water each time. Decant the water after each wash. Wash the copper with a additional 5 ml of methanol. Allow the copper to settle and decant the methanol (be sure to put the methanol in the proper waste container). Gently heat the copper on a hot plate to evaporate any remaining methanol and dry the copper. Once dry, remove the copper from the hot plate and allow to cool. Once cool, determine the mass of the copper. Be sure to record all observations.

RESULTS: Once the mass of recovered copper is known, the percent recovered can be calculated from the following formula: Percent recovery = (mass of copper recovered/initial mass of copper)*100%

REFERENCE: 1. Condike, GF, J. Chem. Ed. 1975, 52, 615.

Copper Cycle Lab Report: Your report for this lab should include the following sections: Abstract: Your third abstract in this class needs to be written individually Introduction: Begin with a statement of the purpose of the experiment Provide any relevant background and key concepts, useful chemical equations and an explanation of the techniques used Data: Include your observations for each reaction Include the initial and final mass of copper Results: For each reaction, give to total balanced equation and the net ionic equation (these may be hand written within the report) Identify each reaction as a precipitation, acid-base or redox reaction (reaction 3 is a decomposition reaction) Report your percent recovery of copper Be sure to attach hand written sample calculations to the back of your report Discussion: Discuss the experiment and any possible sources of error Discuss your percent recovery and possible sources for loss of copper during the reaction cycle.

Chemistry 227 Pre-lab: Equivalence Point by Conductivity Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. When the conductivity probe is placed in a solution of Ba(OH)2, do you expect the conductivity to be high or low? 2. Does the conductivity increase or decrease as you add H2SO4 to the solution of Ba(OH)2? 3. Is the conductivity in the flask greater or less than the original, when you have added an equal number of moles of H2SO4 to the moles of Ba(OH)2 originally present? 4. If you add excess H2SO4, past the equilvalence point, what happens to the conductivity in the flask? 5. Write the balanced chemical reaction for the titration of strontium hydroxide with sulfuric acid. Could you use conductivity to determine the equivalence point of this reaction? Why or why not?

Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted format (be sure to include space for observations) and preparing any tables necessary for data collection.

At the start of your lab, remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA.

Using Conductivity to Find an Equivalence Point OBJECTIVES In this experiment, you will  Hypothesize about the conductivity of a solution of sulfuric acid and barium hydroxide at

various stages during the reaction.  Use a Conductivity Probe to monitor conductivity during the reaction. See the effect of ions, precipitates, and water on conductivity.

INTRODUCTION In this experiment, you will monitor conductivity during the reaction between sulfuric acid, H2SO4, and barium hydroxide, Ba(OH)2, in order to determine the equivalence point. From this information, you can find the concentration of the Ba(OH)2 solution. You will also see the effect of ions, precipitates, and water on conductivity. The equation for the reaction in this experiment is: – – Ba2+(aq) + 2 OH (aq) + 2 H+(aq) + SO42 (aq)   BaSO4(s) + H2O(l) Before reacting, Ba(OH)2 and H2SO4 are almost completely dissociated into their respective ions. Neither of the reaction products, however, is significantly dissociated. Barium sulfate is a precipitate and water is predominantly molecular. As 0.02 M H2SO4 is slowly added to Ba(OH)2 of unknown concentration, changes in the conductivity of the solution will be monitored using a Conductivity Probe. When the probe is placed in a solution that contains ions, and thus has the ability to conduct electricity, an electrical circuit is completed across the electrodes that are located on either side of the hole near the bottom of the probe body. This results in a conductivity value that can be read by the interface. The unit of conductivity used in this experiment is microsiemens per centimeter, or µS/cm. Prior to doing the experiment, it is very important for you to hypothesize about the conductivity of the solution at various stages during the reaction. Would you expect the conductivity reading to be high or low, and increasing or decreasing, in each of these situations?    

When the Conductivity Probe is placed in Ba(OH)2, prior to the addition of H2SO4. As H2SO4 is slowly added, producing BaSO4 and H2O. When the moles of H2SO4 added equal the moles of BaSO4 originally present. As excess H2SO4 is added beyond the equivalence point.

MATERIALS NEEDED magnetic stirrer (if available) Vernier Conductivity Probe 50 mL buret

Measuring Volume Using a Buret 1. Obtain and wear goggles. 2. Measure out approximately 60 mL of ~0. 02 M H2SO4 solution into a 250 mL beaker. Record the precise H2SO4 concentration in your data table. CAUTION: H2SO4 is a strong acid, and should be handled with care. Obtain a 50 mL buret and rinse the buret with a few mL of the H2SO4 solution. Use a utility clamp to attach the buret to the ring stand as shown here. Fill the buret a little above the 0.00 mL level of the buret. Drain a small amount of H2SO4 solution so it fills the buret tip and leaves the H2SO4 at the 0.00 mL level of the buret. Dispose of the waste solution from this step as directed by your instructor. 3. Measure out about 60 mL of the Ba(OH)2 solution. 4. Using a 100 mL graduated cylinder, measure out 25.0 mL of the Ba(OH)2 solution and transfer the solution to a clean 100 mL beaker. Then add 15 mL of distilled water to the beaker. CAUTION: Ba(OH)2 is toxic. Handle it with care. 4. Arrange the buret, Conductivity Probe, beaker containing Ba(OH)2, and stirring bar as shown here. The Conductivity Probe should extend down into the Ba(OH)2 solution to just above the stirring bar. Set the selection switch on the amplifier box of the probe to the 0-2000 µS/cm range. 5. Connect the Conductivity Probe to the computer interface. Prepare the computer for data collection by opening the file “Lab 6: Eqiv. pt” from the Chemistry 227 folder of Logger Pro. 6. Before adding H2SO4 titrant, click and monitor the displayed conductivity value (in µS/cm). Once the conductivity has stabilized, click . In the edit box, type “0”, the current buret reading in mL. Press ENTER to store the first data pair for this experiment. 7. You are now ready to begin the titration. This process goes faster if one person manipulates and reads the buret while another person operates the computer and enters volumes. a. Add 1.0 mL of 0.02 M H2SO4 to the beaker. When the conductivity stabilizes, again click . In the edit box, type the current buret reading. Press ENTER. You have now saved the second data pair for the experiment. b. Continue adding 1.0 mL increments of H2SO4, each time entering the buret reading, until the conductivity has dropped below 200 µS/cm. c. After the conductivity has dropped below 200 µS/cm, add one 0.5 mL increment and enter the buret reading. d. After this, use 2-drop increments (~0.1 mL) until the minimum conductivity has been reached at the equivalence point. Enter the volume after each 2-drop addition. When you have passed the equivalence point, continue using 2-drop increments until the conductivity is greater than 50 µS/cm again. e. Now use 1.0 mL increments until the conductivity reaches about 2000 µS/cm. 8. When you have finished collecting data, click directed by your teacher. 9. Print a copy of the table. 10. Print a copy of the graph.

. Dispose of the beaker contents as

PROCESSING THE DATA 1. From the data table and graph that you printed, determine the volume of H2SO4 added at the equivalence point. The graph should give you the approximate volume at this point. The precise volume of H2SO4 added can be determined by further examination of the data table for the minimum voltage. Record the volume of H2SO4. 2. Calculate moles of H2SO4 added at the equivalence point. Use the molarity, M, of the H2SO4 and its volume, in L. 3. Calculate the moles of Ba(OH)2 at the equivalence point. Use your answer in the previous step and the ratio of moles of Ba(OH) 2 and H2SO4 in the balanced equation (or use the 1:1 – ratio of moles of H+ to moles of OH from the equation). 4. From the moles and volume of Ba(OH)2, calculate the concentration of Ba(OH)2, in mol/L.

EQUIVALENCE POINT DETERMINATION: An Alternate Method An alternate way of determining the precise equivalence point of this titration is to perform two linear regressions on the data. One of these will be on the linear region of data approaching the equivalence point, and the other will be the linear region of data following the equivalence point. The equivalence point volume corresponds to the volume at the intersection of these two lines. 1. Drag your mouse cursor across the linear region of data that precedes the minimum conductivity reading. Click on the Linear Fit button, . 2. Drag your mouse cursor across the linear region of data that follows minimum conductivity reading. Click on the Linear Fit button, . 3. Choose Interpolate from the Analyze menu. Then move the mouse cursor to the volume reading when both linear fits display the same conductivity reading. This volume reading will correspond to the equivalence point volume for the titration. This lab was modified from lab 26 Using Conductivity to fond an Equivalence Point from Chemistry with Computers, Third Edition, Vernier

Using Conductivity to Find an Equivalence Point Lab Report: Your report for this lab should include the following sections: Abstract: Your abstract must be written individually Introduction: Begin with a statement of the purpose of the experiment Provide any relevant background and key concepts, useful chemical equations and an explanation of the techniques used Data: Report the molarity of sulfuric acid used Include a copy of the data table from LoggerPro Results: Include a copy of your titration graph from LoggerPro Report your determined concentration of Ba(OH)2 from both analysis methods Be sure to attach hand written sample calculations to the back of your report Discussion: Discuss the experiment and any possible sources of error In addition, answer the following question as part of your discussion: 1. Which of the two analysis methods do you feel was the most accurate (be sure to support your answer with an explanation)

Chemistry 227 Prelab: Atomic Emission Spectra Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. What wavelengths of the electromagnetic spectrum correspond to visible light? 2. Why do atoms exhibit line spectra? 3. When light is emitted from the hydrogen atom, is the atom moving from a higher energy state to a lower energy state or a lower energy state to a higher energy state? 4. What is the equation for the energy of the hydrogen atom? Give the units associated with the energy equation you report.

Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted format (be sure to include space for observations) and preparing any tables necessary for data collection.

At the start of your lab, remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA.

Atomic Emission Spectra Experiment Goals   

To view the hydrogen emission spectrum and other light spectra To measure the wavelengths of the bright lines in the visible emission spectra of hydrogen and mercury, and calculate the energy of each line. To gain an understanding of the quantized nature of the hydrogen atom

Supplies   

Simple transmission grating spectroscope Hydrogen and other element emission lamps White, red, and green light sources

Definitions 1. A spectroscope is an intrument that allows you to analyze light in some way. In this case, your spectroscope acts like a prism, splitting the light into different wavelengths. 2. A continuous spectrum is one in which a rainbow of colors is seen when viewed through a spectroscope or prism. 3. A line spectrum, when viewed through a spectroscope, appears as one or more sharp narrow lines. 4. A band spectum is intermediate in appearance. It is between a line spectrum and a continuous spectrum. A range of nearby wavelengths are emitted.

Background The understanding of the internal structure of the atom was advanced when Niels Bohr explained the cause of the emission spectra of atoms using the concept of the quantization. He stated that electrons in the atom could exist at finite energy levels. Electrons within an atom can be excited to higher energy states through various means, including heating the atoms or a electric discharge. Exciting the electrons causes the atoms to emit electromagnetic radiation as the excited electrons relax into a lower energy state. The emitted light can be passed through diffraction grating to separate it in to its individual components generating an atomic emission spectrum, a line spectrum characteristic to the particular sample of atoms. For instance, the emission spectrum of hydrogen consists of only four visible lines: red, bluegreen, violet and deep violet (this your eyes may not be sensitive to). Each color corresponds to the transition of an electron from an excited state, a higher principle energy level, to a lower

principle energy level, possibly the ground state or some allowed energy level in between. As the electron drops, it emits light in the form of a photon which may or may not be in the visible region. The energy associated with the emitted light equals the difference between the energy of the electron in the higher energy level minus the energy of the electron in the lower energy level. In this experiment you will use a spectroscope and gas discharge lamp to measure the wavelength of each bright line in the visible atomic emission spectra of both hydrogen and mercury. You will then use these measurementsto calculate the photon energy for each bright line. The final results will be the wavelength (in meters) and the photon energy of each bright line measured for hydrogen and mercury.

How To Use Your Instrument WARNING!! The power supply to the discharge lamps is 5000 volts ! DO NOT TOUCH Figure 2 shows the spectroscope you will be using. It is a small box, with a transmission grating in one side and a narrow slit directly opposite the grating. To observe a spectrum you point the slit toward a light source and look through the grating. You will see am image of the spectrum along the back wall of the box, just over the wavelength scale. To determine the wavelength of the light you observe, you will make use of the distance of the slit to the image, which together with the distance from the slit to the grating will allow you to calculate the wavelength. Using the scale on the back wall of the spectroscope, measure the distance from the slit to the image of each line, this distance will be called a in Equation 1. The distance from the slit to the grating is b, and the distance between the grating lines is d (in cm). (Note, there are 13,400 grating lines per inch.)

Equation. 1 Things That Might Be Useful  E = - (2.178 x 10-18 J) (1/n2final - 1/n2initial)  E = hc/ 1 nm = 1 x 10-9 m

 = ad/(a2 + b2)1/2

Spectrum of a Single Electron Element – Hydrogen • Record the line color and its position for 3 or 4 brightest lines observed in table 1 and calculate its wavelength using equation 1 from your lab manual. Be sure to include a sample of your work for one of the calculations. Table 1: Bright Line Spectra for Elemental Hydrogen. Line Color

Spectral Line Position (cm)

Wavelength (nm) Caculated using Eq. 1

Clearly show your work here for your calculation of wavelength:

Spectrum of Multi Electron Elements and Other Miscellaneous Spectra Qualitatively observe several other spectra and make notes below about there differences to hydrogen.

Qualitatively compare the 3 brightest lines from the emission spectra of the elemental mercury lamp to the spectra of a fluorescent lamp and an incandescent bulb.

Observe the solar spectra (if daylight is available).

Data Analysis 1) For the first four electronic transitions in the Balmer Series, calculate the change in energy of the electron (∆E), the predicted energy of the emitted photon (Ephoton) and the predicted wavelength of the emitted photon (λphoton). Put the calculated values in table 2 and be sure to clearly show an example of each calculation in the space provided. Table 2: Calculated Values for the Balmer Series of Hydrogen. Electronic Transition

ΔE (J)

λphoton (nm)

Ephoton (J)

n3 → n2 n4 → n2 n5 → n2 n6 → n2 Clearly show the following calculations for the n3 → n2 transition. -- the change in energy of the electron, ΔE

-- the predicted wavelength of the emitted photon, λphoton

2) Based on your theoretical calculations, match the electronic transitions in the Balmer Series to the spectral lines you observed in Table 3. Then calculate the percent error between your experimentally determined and calculated wavelengths. Table 3: Comparison of Experimental and Accepted Wavelengths from the Balmer Series. Electronic Spectral Line Accepted λ (nm) from Experimental λ (nm) Percent Error Transition Color Observed from Table 1 Table 2

n3 → n2 n4 → n2 n5 → n2 n6 → n2

Below, clearly show your percent error calculation for the n3 → n2 transition.

3) It is not possible to observe the n7 → n2 transition in the Balmer Series. Why do you think that is?

4) Emission spectra are sometimes referred to as atomic fingerprints. Is it possible to use them to identify elements in a unknown sample? Provide evidence; think about the Hg spectrum and that of a fluorescent bulb.

5) Why might sodium vapor lights cast a different color than fluorescent lamps?

6) Calculate the ionization energy of the hydrogen atom? Think about taking an electron from its ground state, n = 1, to a position/energy level far, far away from the nucleus, n = ∞.

Atomic Emission Spectra Lab Report:

There is not a formal lab report for this lab. Complete the above pages and submit them to your TA.

Chemistry 227 Pre-Lab: Determining the Concentration of a Solution Part A Answer the following questions in your lab notebook (be sure to show work for any calculations): 1. You are given a colored solution that is labeled 1M. You need to prepare a solution from this that is 0.5 M. Describe your procedure in detail. 2. What is the relationship between absorbance and transmittance? 3. Allura red is a commonly used red food dye. Does allura red transmit or absorb red light? 4. If 5.00 ml of a 0.5 M solution is diluted to a final volume of 100.0 ml, what is the concentration of the dilute solution?

Part B Prepare your lab notebook for the lab. This includes stating the purpose of the experiment, summarizing the procedure in a bulleted format (be sure to include space for observations) and preparing any tables necessary for data collection.

At the start of your lab, remove the copies of the pages where you completed the above work from your lab notebook and turn them into your TA.

Determining the Concentration of a Solution: Beer’s Law OBJECTIVES In this experiment, you will    

Prepare Allura red standard solutions. Use a Colorimeter to measure the absorbance value of each standard solution. Find the relationship between absorbance and concentration of a solution. Use the results of this experiment to determine the concentration of Allura red in gatorade.

INTRODUCTION The primary objective of this experiment is to determine the concentration of Allura red in a commercially available mouthwash. You will be using the Colorimeter shown in Figure 1. In this device, light from the LED light source will pass through the solution and strike a photocell. A higher concentration of the colored solution absorbs more light (and transmits less) than a solution of lower concentration. The Colorimeter monitors the light received by the photocell as either an absorbance or a percent transmittance value.

Figure 1

Figure 2

You are to prepare five Allura red solutions of known concentration (standard solutions). Each is transferred to a small, rectangular cuvette that is placed into the Colorimeter. The amount of light that penetrates the solution and strikes the photocell is used to compute the absorbance of each solution. When a graph of absorbance vs. concentration is plotted for the standard solutions, a direct relationship should result, as shown in Figure 2. The direct relationship between absorbance and concentration for a solution is known as Beer’s law. The concentration of Allura red in an unknown solution (gatorade) is then determined by measuring its absorbance with the Colorimeter. By locating the absorbance of the unknown on the vertical axis of the graph, the corresponding concentration can be found on the horizontal axis (follow the arrows in Figure 2). The concentration of the unknown can also be found using the slope of the Beer’s law curve.

MATERIALS NEEDED Vernier Colorimeter one cuvette two 10 mL pipets pipet pump or pipet bulb

PROCEDURE 1. Obtain and wear goggles 2. Add about 30 mL of Allura red stock solution to a 100 mL beaker. Add about 30 mL of distilled water to another 100 mL beaker. Be sure to record the concentration of the stock solution of allura red. 3. Label five clean, dry, volumetric flasks 1-5. In these flasks, you will prepare solutions of Allura red varying in concentration from approximately 6 x 10-6 M to 2 x 10-5 M. You may wish to check your concentrations and calculations with your TA before making the solutions. Make the solutions by pipetting the correct quantity of the Allura red stock solution into the volumetric flask. Be careful to avoid getting liquid above the fill line. Thoroughly mix each solution by inverting the flask several times and shaking. 4. Connect the Colorimeter to the computer interface. Prepare the computer for data collection by opening the file “Lab 8: Beer’s Law” from the Chemistry 227 folder of Logger Pro. Set the colorimeter to a wavelength of 470 nm. 5. You are now ready to calibrate the Colorimeter. Prepare a blank by filling a cuvette 3/4 full with distilled water. To correctly use a Colorimeter cuvette, remember:    

All cuvettes should be wiped clean and dry on the outside with a tissue. Handle cuvettes only by the top edge of the ribbed sides. All solutions should be free of bubbles. Always position the cuvette with its reference mark facing toward the white reference mark at the top of the cuvette slot on the Colorimeter.

6. Calibrate the Colorimeter. Fill a cuvette with water and place it in the colorimeter. In Logger Pro click on “Experiment” then, from the drop down menu select “calibrate” and “lab pro colorimeter”. In the popup window check the box next to “one point calibration”. Click “Calibrate now” and enter “100” in the box provided (100% T). Click “Keep” and then “Done”. You should now see an absorbance reading of 0.000. . 7. You are now ready to collect absorbance data for the five standard solutions. Click Empty the water from the cuvette. Using solution 1, rinse the cuvette twice with ~1 mL amounts and then fill it 3/4 full. Wipe the outside with a tissue and place it in the Colorimeter. After closing the lid, wait for the absorbance value displayed on the monitor to stabilize. Then click type the concentration of the standard solution in the edit box, and press the ENTER key. The data pair you just collected should now be plotted on the graph. NOTE: When entering values, 2 x 10-6 can be entered as 2E-6. You may need to click on the Autoscale button to rescale the graph. 8. Discard the cuvette contents as directed by your teacher. Rinse the cuvette twice with solution 2, and fill the cuvette 3/4 full. Wipe the outside, place it in the Colorimeter, and close the lid. When the absorbance value stabilizes, click , type the concentration of the standard solution in the edit box, and press the ENTER key.

9. Repeat the Step 8 procedure to save and plot the absorbance and concentration values of solutions3-5. Wait until Step 12 to do the unknown. When you have entered all of your standard solutions, click . 10. Be sure to record the absorbance and concentration data pairs that are displayed in the table. 11. Examine the graph of absorbance vs. concentration. To see if the curve represents a direct relationship between these two variables, click the Linear Fit button, . A best-fit linear regression line will be shown for your five data points. This line should pass near or through the data points and the origin of the graph. (Note: Another option is to choose Curve Fit from the Analyze menu, and then select Proportional. The Proportional fit has a y-intercept value equal to 0; therefore, this regression line will always pass through the origin of the graph). 12. Obtain a small amount of Gatorade in a small clean beaker. Use the pipette to deliver 10 mL of the Gatorade to a clean, dry volumetric flask. Finish preparing your unknown by diluting the Gatorade to a total volume of 100 mL and mix thoroughly. Rinse the cuvette twice with the unknown solution and fill it about 3/4 full. Wipe the outside of the cuvette, place it into the Colorimeter, and close the lid. Read the absorbance value displayed in the meter. to read (Important: The reading in the meter is live, so it is not necessary to click the absorbance value.) When the displayed absorbance value stabilizes, record its value. 13. Discard the solutions as directed by your teacher. Proceed directly to Steps 1 and 2 of Processing the Data.

PROCESSING THE DATA You may use Microsoft Excel to plot the data and obtain a linear relationship between the data, or you may use the following method: 1. Use the following method to determine the unknown concentration. With the linear regression curve still displayed on your graph, choose Interpolate from the Analyze menu. A vertical cursor now appears on the graph. The cursor’s concentration and absorbance coordinates are displayed in the floating box. Move the cursor along the regression line until the absorbance value is approximately the same as the absorbance value you recorded in Step 12. The corresponding concentration value is the concentration of the unknown solution, in mol/L. 2. Print a graph of absorbance vs. concentration, with a regression line and interpolated unknown concentration displayed. To keep the interpolated concentration value displayed, move the cursor straight up the vertical cursor line until the tool bar is reached. Enter your name(s) and the number of copies of the graph you want. 3. Use the calibration curve to determine the concentration of allura red in the diluted solution. Calculate the concentration of allura red in the undiluted Gatorade.

Allura Red has the following chemical structure:

With the help of your TA, calculate the molar mass of allura red and determine the number of molecules of allura red you would consume if you drank one 20 ounce bottle of Gatorade.

This lab was modified from lab 11 Determining the Concentration of a Solution: Beer’s Law from Chemistry with Computers, Third Edition, Vernier

Beer’s Law Lab Report: Your report for this lab should include the following sections: Abstract: Your abstract must be written individually Introduction: Begin with a statement of the purpose of the experiment Provide any relevant background and key concepts, useful chemical equations and an explanation of the techniques used Data: Include a table showing the concentrations of your standard solutions, how they were made and their absorbance values Results: Include a copy of your calibration curve State the concentration of allura red in your diluted Gatorade solution and the undiluted Gatorade Be sure to attach hand written sample calculations to the back of your report Discussion: Discuss the experiment and any possible sources of error In addition, answer the following questions as part of your discussion: 1. How many molecules of allura red you would consume if you drank one 20 ounce bottle of Gatorade

Chemistry 227 Electron Density Pre-lab Before coming to lab this week, complete the following questions: Read sections 9.6 and 10.5 in your text book.

1. What are the three type of bonding classifications?

2. What is the description polar bond referring to?

3. The interactions between water molecules can be described as electrostatic or coulombic, where areas of positive charge are attracted to areas of negative charge. Draw a cartoon of how three water molecules might be arranged in space, given this electrostatic interaction.

Chemistry 227 Lab

Electron Density Lab The Electron Density Lab can be accessed two ways. It is on the desktop on all the lab computers. Just click to open it and follow the instructions. You can also access the lab and its images by going to the web site: web.pdx.edu/~shusteg. This is Dr. Shusterman’s website for her course. Scroll down the main page and you will see a link for CH 227 students. This file is in pdf format. You need Acrobat Reader to view it. This is a free viewer and there is a link on the page to the site where you can download it. All campus computers should already have this software and most of you already use it to view lectures. If you have not already done so, you might also find it useful to work through the Electron Density Tutorial also at Dr. Shusterman’s website. This gives practice in the interpretation of electron density models. Report your answers here (turn in to TA after Thanksgiving):

Predict Polarity: A) CH4

B) NH3

C) CH3CH2OH

D) CH3NH2

E) CHClCHCl

F) CCl2CH2

Which is most soluble in water?

How could you separate these compounds?