Unformatted Document Excerpt
Course Hero has millions of student submitted documents similar to the one
below including study guides, practice problems, reference materials, practice exams, textbook help and tutor support.
Course Hero has millions of student submitted documents similar to the one below including study guides, practice problems, reference materials, practice exams, textbook help and tutor support.
Manual Lab for Distance Chem 105 Chemistry in Context I
First Edition, February 2006
Department of Chemistry
Compiled by Mark A. Griep, Department of Chemistry University of Nebraska-Lincoln
Manual for Distance Chem 105: Chemistry in Context I 1st edition, February 2006
Description of the Chem 105 Lab Portion An introduction to the concepts of chemical research including but not limited to: chemical lab safety, identifying inorganic salts, identifying organic materials, measuring gas volumes, measuring liquid volumes, measuring concentration, determining measurement accuracy, determining reaction stoichiometry, and making pH measurements.
TABLE OF CONTENTS
History & Background of these experiments .......................................................................................2 GENERAL LAB INSTRUCTIONS ....................................................................................................4 Checking In. ....................................................................................................................................... 4 Safety.................................................................................................................................................. 4 Fume Hoods. ...................................................................................................................................... 4 Sinks. .................................................................................................................................................. 4 Broken Glass. ..................................................................................................................................... 4 Spilled Solutions. ............................................................................................................................... 4 Use of Reagents.................................................................................................................................. 5 Washing Glassware. ........................................................................................................................... 5 Tap Water versus Distilled Water. ..................................................................................................... 5 The Psychology of Safety Preparation. .............................................................................................. 5 The Psychology of Preparation Before Research............................................................................... 5 Laboratory Notebook. ........................................................................................................................ 6 Laboratory Reports............................................................................................................................. 6 1. Electrolysis of Water .......................................................................................................................8 2. Qualitative Analysis of Inorganic Ions..........................................................................................12 3. Models of Covalent Structures......................................................................................................15 4. Virtual Radiation Safety................................................................................................................19 5. "Me & Isaac Newton" ...................................................................................................................24 6. Qualitative Analysis of Organic Compounds................................................................................26 7. Vapor Density of Butane...............................................................................................................31 8. Molar Volume of Hydrogen ..........................................................................................................35 9. Cheese Puff Calorimetry ...............................................................................................................38 10. Cost of an Aluminum Atom ........................................................................................................42 11. Household Acids, Bases, and pH ................................................................................................44 12. Peroxide Content in Teeth Whiteners .........................................................................................48
History & Background of these experiments
The experiments in this manual were edited and compiled in February 2006 by Dr. Mark A. Griep, Department of Chemistry. Griep also wrote the "Facts to Know" sections. High School chemistry teacher Greg Cooper tested the experiments during Summer 2005 and wrote the first draft of the Procedures. Some experiments were developed from ideas or demonstrations and others from the experiments tested by Brittney Schirber for the classroom version of Chem 105. Dr. Griep gratefully acknowledges Cooper's and Schirber's efforts.
Most of these experiments are based on exercises developed by many chemists over many years. The following sources will lead you to some recent literature for these experiments. "Psychological Safety Preparation" was developed by Dr. Mark Griep with some things adapted from Dr. Kingsbury's "High School Safety Web Pages" at www.chem.unl.edu/safety/hslcon.html Expt 1 "Water Electrolysis" experiment was adapted from the UNL Chem 109 Lab Manual but is described in all introductory chemistry textbooks. The following references provide examples and tips on doing this experiment: 1) Eggen, Kvittingen (2004) J. Chem. Educ. 81, 13371338; 2) Zhou (1996) J. Chem. Educ. 73, 786-787; 3) Manjkow, Levine, Rowley (1986) J. Chem. Educ. 63, 809-810. Expt 2 "Qualitative Analysis of Inorganic Compounds" was adapted from two sources. The water analysis was developed by Brittney Schirber from a suggestion by Dr. C. William McLaughlin. The qualitative analysis scheme was provided by Greg Cooper. The principles behind these experiments are described in all introductory chemistry textbooks. The following references provide many examples and tips on doing this experiment: 1) Oliver-Hoyo, Allen, Solomon, Brook, Ciraolo, Daly, Jackson (2001) J. Chem. Educ. 78, 1475-1478; and 2) Petty (1991) J. Chem. Educ. 68, 942-943 Expt 3 "Models of Covalent Structures" was created by Greg Cooper and Dr. Griep for this manual. Expt 4 "Virtual Nuclear Safety" was adapted from two sources. The household products are part of UNL Chem 109 classroom demonstrations. The inverse-square dependence was suggested by Dr. Mark Griep. The inverse square law applies to radiation, sound, light, electric fields, and gravity. You can find out more by searching for "inverse square law" on the internet. Expt 5 "Me & Isaac Newton" report was created by Dr. Mark Griep for this course. You can learn more about his use of movies to teach chemistry by reading an article he wrote with Marjorie Mikasen: Griep & Mikasen "Based on a True Story: Using Movies as Source Material for General Chemistry Reports" in the Journal of Chemical Education October 2005 issue. Expt 6 "Qualitative Analysis of NSAIDs" was developed by Brittney Schirber from the various experiments suggested in two textbooks: "Systematic Identification of Organic Compounds" 6th Ed., 1980 by Shriner, Fuson, Curtin, & Morrill and "Multiscale Operational Organic Chemistry" 2002 by Lehman. The following references provide examples and tips for this sort of experiment: 1) Cooley, Williams (1999) J. Chem. Educ. 76, 1117-1120; 2) Griswold, Rauner (1991) J. Chem. Educ. 68, 418-420. The following reference provides a detailed analysis of DMSO solubility: Balakin et al. (2004) J. Biomolecular Screening 9, 22-31. Expt 7 "Vapor Pressure of Butane" was adapted from the UNL Chem 109 Lab Manual. Expt 8 "Molar Volume of Hydrogen" was adapted from the microscale experiment listed on Dr. David Brooks' website at http://dwb.unl.edu/Chemistry/MicroScale/MScale15.html. Expt 9 "Cheese Puff Calorimetry" was adapted from the UNL Chem 109 classroom demonstrations. Expt 10 "Cost of an Aluminum Atom" was adapted from Bieron and Dinan's "Avogadro Goes to Court" posted on the National Case Study Teaching in Science website (http://www.sciencecases.org/avogadro/avogadro.asp) Expt 11 "Nitrate Content in Soil" colorimetric assay by was adapted from numerous websites. For background, there is an article titled "Human Alteration of the Global Nitrogen Cycle: Causes and Consequences" published in Issues in Ecology in Spring 1997. Expt 12 "Peroxide Content of Tooth Whiteners" was developed by Brittney Schirber from an idea proposed by Dr. C. William McLaughlin for UNL Chem 105. For some background, there is an article titled "Shedding Some Light on Teeth Whiteners" in Chemical & Engineering News released on February 10, 2003.
GENERAL LAB INSTRUCTIONS
Read these before doing any work in your home laboratory.
When the new semester begins, every student checks the contents of their drawer of equipment to make sure that everything is there that should be and that it is in proper order. Likewise, the distance student should check the entire contents of their box an familiarize yourself with it. It would also be a good idea to create a special storage place for the material.
General Lab Safety--Make sure you know the location of your fire extinguishers (if you don't own one, you should purchase one and keep it under your kitchen sink) and evacuation exits. Your life may depend on it. Goggles--Whenever you are working on these experiments, you are required to wear protection for your eyes. Goggles provide the best protection so you will be provided with a pair of goggles. The goggles were chosen to fit over eyeglasses. Put your goggles on before beginning the exercise and keep them on until you complete the experiment. Gloves--In an actual laboratory, researchers often wear plastic gloves to prevent their skin from contaminating their experiments and to protect themselves from the occasional toxic substance. Nitrile gloves are preferred by many people because the latex ones can cause allergic reactions with skin. The negative side to using nitrile gloves is that they are thicker and more slippery than latex gloves. The experiments in this lab were chosen for their low toxicity and you don't really need to wear gloves. Nevertheless, we have included two pairs of small, medium, and large gloves. Hair and Clothing--Even though it is a clich that scientists wear labcoats and tie their hair in buns, they do it for safety reasons. Long flowing hair, dangling jewelry, loose clothing, and open-toed shoes are not appropriate attire for safe experimentation.
Chemical fume hoods are used whenever you want to control chemical emissions. They are used to hold smelly reagents and to carry out reactions that release noxious fumes. There is only one experiment that you will do that requires good ventilation and we ask that you perform it on your patio.
We have selected lab materials that are safe enough to wash down your kitchen sink or thrown in the garbage. It is important to keep in mind, however, that when you work with chemicals in the real world there are special Hazardous Materials disposal guidelines and that there are some things that cannot be discarded down the sink.
Sweep up any broken glass immediately. Do not throw the shards in with your regular garbage. Instead, discard the shards in a labeled "Sharps" container made from an empty milk carton. Make one of these now so that it is ready when you need it.
Spilled Solutions. 4
If you spill strong acid or base on your countertop, you should neutralize it before you wipe it up. If you spill a strong acid, then add enough baking soda (a weak base) to soak up all of the solution. If you spill a strong base, then add enough vinegar (a weak acid) to lower the pH
Use of Reagents.
Always read the label carefully before using any reagent to make certain that you are obtaining the correct material. If you do not read the label with sufficient care, you may obtain the wrong reagent and as a result you will have to repeat part or all of the experiment. If you wish to use a reagent, pour a little more than the desired amount into a clean beaker (or test tube or graduated cylinder) and then pipette from this smaller supply. Please be careful not to contaminate or waste the reagents.
Your glassware is new when it was packaged but will certainly have packing dust on it upon its arrival. The first day that you check the box contents is a good one to wash the glassware. We have included a small amount of glassware cleaning soap with your kit. The key to washing glassware is to use very little detergent. You want to rinse thoroughly so that you don't leave any residue whatsoever. Too much detergent is a common source of contamination in undergraduate chemistry labs that causes low yields. You should also reserve one cotton dishtowel for your glassware. This will prevent your other one from contaminating your experiments. Use a basin with lots of tap water and only a few drops of detergent. After scrubbing your glassware in the basin, rinse with plain tap water and then rinse with distilled water. Put your clean test tubes in a rack upside-down so you can tell which are clean and which are used. Do not place detergent directly on the glassware.
Tap Water versus Distilled Water.
When an experimental procedure requires the addition of water, it is understood that distilled water is to be used. For these home labs, we have relaxed this requirement for most of the experiments. In one or two of them, however, you will need to purchase a small bottle of distilled water. Water distillation involves heating the water to the point of boiling and then collecting the condensate. This action removes all chemicals from the water except those that have the same boiling point as water.
The Psychology of Safety Preparation.
Safety is a state of mind. The following rumination might help put you in the proper mood. Treat all solutions as though they are corrosive toxins. Wear Gloves and Goggles. Point reaction tubes and containers so that they point away from you and others. Always add things to water. Never add water to strong acids or bases because the heat generated may splatter. Treat all gases and vapors as though they will knock you unconscious if you breathe them. Do not smell flasks directly. Waft the vapor carefully toward yourself. Remember that hot glassware looks just like cold glassware. Do not drink or eat while doing your labs. Water looks the same whether something is dissolved in it or not. Most serious accidents in high schools and colleges occur when students are handling solutions in glassware
The Psychology of Preparation Before Research.
Familiarize yourself with the entire written lab material before you start. Do the pre-lab "Safety & Purpose" exercises. Make a list of questions you'd like answered from the experiment. Don't rely solely on the ones we've listed. The purpose of doing the labs is so that you have a "hands on" opportunity to verify important chemical principles. Record all your questions and data in a sturdy bound laboratory notebook. In medicine and industry, if an observation is not recorded at the time
you make that observation, it doesn't count. Record your color changes, temperature changes, physical state changes, etc. when you see them happen. It is sometimes necessary to make changes to a lab procedure. Record the changes you decide to make. After you've completed an experiment, look over your observations to make sure that they make sense. If you think you made a mistake, repeat that portion of the experiment. Never erase or otherwise obliterate something you have written. If you wish to ignore it, then draw a single line through it and give an short explanation next to it. Finding your mistakes is the best sort of self-instruction. Clean your glassware at the end of the period so that it will be ready for the next.
Each person must keep a laboratory notebook. The purpose of a lab notebook is to provide experience in the direct recording of experimental data in organized form. It should be possible for someone not familiar with the experiment to be able to clearly understand the procedures used and what data were obtained. The lab notebook should be brought to each session and the data collected during that session should be directly recorded. When working with a partner, it is acceptable to write all of the data in one book during the lab period and then later copy the data into the other notebook. The information should clearly indicate that it is a copy and from whom it was obtained. Your lab notebook will be checked and graded at the end of the semester. Every experiment in the notebook should be recorded in ink as follows: 1. Date, Title and Partner's name(s). 2. Procedure Used (refer to the page number in the lab manual). Note any modifications to the procedure in the lab manual. 3. Raw Data collected while the experiment takes place. Data that has been recorded on scraps of paper and then later copied into the notebook is unacceptable practice. 4. It is recommended that you attach Printouts, photos or their photocopies, TLC plates or their photocopies, etc to your notebook after properly labeling them to their date, your name and partner's name, expt # and page in your notebook to which they correspond.
The reports should be succinct. For the purposes of this course, only the following elements need to be included in your report: 1. Your name, partner's name(s), experiment number and title. 2. Purpose (or Objectives). A very brief statement about what you were trying to discover. Use as few sentences as possible. 3. Experimental Methods. The procedures used to obtain the data. It is mentioned here because it is an integral part of formal reports in that it allows other researchers to know exactly what they would have to do to obtain your results. However, you may indicate this part to the page number(s) in the lab manual in your lab report because we all know what procedure you used. All you need to do is mention that you followed the procedure in the book or the text and then note any modifications to the written procedure. 4. Results. Present the observations and calculations that you made. Very few measurements in chemistry give absolute answers. It is nearly always necessary to estimate the uncertainty, which in turn requires an explanation as to the possible sources of the uncertainty. Identify each figure by a short and appropriate legend. Explain your reasoning for any unknowns that were identified. Because this is a lab course, sometimes the experiments do not work. If a lab does not work it is up to you to hypothesize why it did not and state what you would do differently next time. 5. Discussion. Read the experiment's title and purpose before you begin your discussion so that you are discussing the right problem. Bring everything (data, calculations, analysis) together in
your discussion. Find a high-quality source on the internet that provides relevant material for your results and cite it in your discussion. 6. Acknowledgments. Estimate the extent to which you contributed to generating the data. If you received significant help in data analysis or interpretation, if it was a team effort, give credit to your fellow students.
The reports are due one week after you begin an experiment.
1. Electrolysis of Water
I. Objectives To decompose water to produce two volumes hydrogen gas and one volume oxygen gas. To separate a molecule into its component elements. II. Safety Considerations Wear eye protection at all times. Be careful when using any power source because there is a shock hazard. You will be producing oxygen and hydrogen gas. Oxygen supports combustion and hydrogen is extremely flammable. Keep away from any flame or heat. Electricity will be flowing through water. Although 9 volts is not very much you may feel a tingle if you place your fingers in the salt solution. Salt in the eyes will burn.
III. Materials From Your Home 1 Tablespoon of salt 1 Stockpot or Saucepan (or any container that will hold at least 6 cm of water) Ruler
Provided for you 2 small test tubes Beaker, small 1 9V battery Sharpie Narrow range pH paper.
IV. Facts to Know Water electrolysis into hydrogen and oxygen gases is one of the most surprising experiments because it is so simple to do. That's why we've chosen it as our first experiment of the semester. William Nicholson and Anthony Carlisle first reported water electrolysis in 1800. They had just learned of Allesandro Volta's pile (the first battery was made of a stack of zinc and copper discs, one on top the other, separated only by oily pads) and were doing their best to reproduce the results he described in a letter. When they used a drop of water to help improve one of their electrical connections, they noticed that a gas was being generated in the water at the point of contact with the wire. After some work, they proved that it was hydrogen gas. Next, they placed a volume of water in line with the current and saw that a gas was generated at both points where the wires met the water. They determined that the other gas was oxygen. Finally, they showed that the rate of gas generation increased in proportion with the electrical current. The reactions that take place at each electrode are as follows: At the anode: 2 H2O --> O2(g) + 4 H+ + 4 e-- At the cathode: 4 H2O + 4 e-- --> 2 H2(g) + 4 OH-- Also taking place is the acid and base neutralization reaction: H+ + OH-- --> H2O The electrolysis net reaction is: 2 H2O --> O2(g) + 2 H2(g) This net reaction equation indicates that twice as much hydrogen gas will be produced as
oxygen gas. You will use your apparatus to determine how close experiment meets with theory. There are many practical reasons that your ratio won't be exactly two. Some of the reasons are: 1) oxygen gas is more soluble in water than hydrogen gas (31 mL oxygen is dissolved per liter in contrast to 18 mL hydrogen per liter at 20 C and 100 kPa); 3) dirty electrodes may contain impurities that will generate hydrogen peroxide at the anode; 4) the acid or base that is produced at the electrodes will corrode them so that they produce the gases more slowly, which gives more time for the oxygen gas to dissolve in the water, which lowers the amount of oxygen gas collected. In this lab, an electrical current will be passed through water, and this will cause the hydrogen and oxygen that make up water to dissociate. Since water doesn't transport electrical charge very well, it is common to add an electrolyte (an acid, base, or a salt) to speed the flow of the electrical current through the water. The most common acid electrolyte for this experiment is sulfuric acid (H2SO4), the most common base is sodium hydroxide (NaOH), and the most common salt is sodium sulfate (Na2SO4). None of the electrolytes are involved in the reactions at the electrodes. We will use table salt (NaCl) as our electrolyte because it doesn't cost much and is readily in every home. "Low sodium salt" or "Lite Salt" will work just as well in this experiment because they contain KCl, which is also an electrolyte.
V. Procedure 1. Make sure that the stockpot or saucepan is clean. Fill it with tap water until it is about 6 cm deep. The exact amount is not critical except that there has to be enough to completely immerse a 9 V battery when it is standing in the water. Dissolve 1 tablespoon of salt in the water. 2. You will be using both of your hands continuously for the next steps. You may want to practice first and/or do the experiments several times and then take the average measurement. First, label one small test tube as "positive" and the other as "negative". These refer to the electrodes on the battery. Fill the test tubes with tap water (not the salt water). Use the fingers of the same hand that holds the tubes to place your fingers over their ends of the test tubes. Invert the tubes and submerge them into the salt water but don't remove your fingers. If there is any air in either of them, use a Sharpie to mark the starting volume of gas in the tube. 3. Inspect the battery and note which electrode is positive and which is negative. Place the 9volt battery in the water. You should see gas bubbles emerging from each electrode on the battery. Carefully remove your fingers from the test tubes and simultaneously place the "positive" test tube over the positive electrode and the "negative" test tube over the negative electrode. Do not to lift the test tubes above the surface of the water. Because the battery is not waterproof, some salt water may enter the battery and cause bubbles to emerge from the lip of the battery on the positive side rather than directly from the electrode. If this happens, capture all of the bubbles no matter where they emerge. 4. Wait ten minutes, place your fingers over the test tube ends, and mark the final water level. Place the tubes with their contents into the small beaker. 5. Use a ruler to measure how many millimeters of water have been displaced. This value is proportional to the gas volume generated. "Positive" Tube: _________ mm "Negative" Tube: _________ mm
6. Use strips of narrow range pH paper to estimate the pH of the water in each of the test tubes. The pH of the two tubes should differ but possibly not by much. "Positive" Tube: pH = _________ "Negative" Tube: pH = _________
VI. Report Questions 1. What was your ratio of hydrogen gas volume to oxygen gas volume? How does your ratio compare to the theoretical value? How do explain the difference between your ratio and the theoretical ratio. 2. If you did the experiment more than twice, did your ratio improve every time or did it vary relative randomly relative to your average value? Even if you did not do the experiment more than twice, explain how those two outcomes could be used to determine whether the difference from theoretical was due to error from the experimenter or due to the experimental design. How would you test your hypothesis? 3. What evidence do you have that a chemical change took place during the experiment? 4. Explain the action of the pH indicator at the electrodes. 5. Are the gases lighter or heavier than air? 6. Explain the safety mistakes in the following situations: A. John was late for his lab class and didn't have time to eat his lunch so he thought he would snack on it during lab. After John raced to the lab, he was hot and decided that, since the experiment didn't involve anything especially toxic, he didn't need to wear his goggles and that his eyeglasses would be sufficient. After lab was over, he washed his used beaker with sulfuric acid to make sure it was really clean. B. Melissa has long flowing blonde hair that she insists on always wearing down. She reports to lab on a sunny fall day wearing her seasonal flip-flops and finds that she is the first one there. Since Melissa has read the lab ahead of time in preparation, she decides get started before anyone else has arrived to lab. She thinks she'll complete the lab early enough to go outside and enjoy the rest of the sunny day.
2. Qualitative Analysis of Inorganic Ions
I. Objectives To determine the ions present in tap water. To do a precipitation reaction. To learn about iron oxidation states. II. Safety Considerations Wear eye protection at all times. Iron (II) Chloride is slightly toxic if ingested and is a body tissue irritant. Iron (III) Chloride is also a skin and tissue irritant, corrosive and toxic by ingestion. Sodium Hydroxide is a corrosive liquid is very dangerous to your eyes and skin. It is recommended that you wear gloves while handling. III. Materials From Your Home Distilled water, 500 mL Tap water Provided for you Sharpie 1 Dropping pipet, plastic 4 Erlenmeyer Flasks, 125 mL 1 Graduated cylinder, 50 mL 2 Beakers, 250 mL 2 rubber stoppers, for 125 mL flask 0.5 grams of FeCl2 0.5 grams of FeCl3 1 molar NaOH, 1 mL
IV. Facts to Know It is very important to know what ions are present in the water you drink. Some ions are harmful at high concentrations. For instance, there are 80 communities in Nebraska that have wells containing arsenic and/or uranium at levels that slightly exceed the Federal Government standards. Some ionic water contaminants, such as nitrates from fertilizer run-off, can even change the pH of the water. Qualitative analysis is the identification of a sample's components. This contrasts with quantitiative analysis, which seeks to determine the amounts of each component. The best qualitative analysis method should be able to identify a single type of ion even in the presence of other similar ions. For instance, the ideal analysis would be able to establish the presence of a small amount of sodium ion in the presence of a high concentration of potassium ions. Unfortunately, this is not always possible because sodium and potassium have similar chemical properties. It is only possible to distinguish sodium and potassium by conducting a series of analytical experiments. After that, it is necessary to confirm the identity of each ion by isolating them from the solution. In this experiment, you will focus on iron ions. You will carry out the same tests on tap water in a preliminary assessment of its iron content. The two most common reactions used to confirm the presence of ions in solutions are called precipitation and complexion. Precipitation reactions occur when certain ions are added to a test solution and a solid is formed. Since it is already well known which cations form insoluble complexes with which anions, it is a simple matter to select test solutions that can distinguish
between many salt solutions. You won't perform any complexation reactions today but they are when cations form covalent bonds with one or more ligands and remain soluble. Most of these have distinct colors and are very soluble. Many transition metal ions form color complexes.
V. Procedure Part A: Precipitation Tests and Testing Water for Hardness Ions 1. You will need three solutions to complete this lab. The NaOH is already a solution at 1 molar concentration (probably labeled "1 M NaOH"). You will need to make the other two iron stock solutions. Label one of your small beakers as "FeCl2 stock" and other as "FeCl3 stock". Take the 0.5 grams of each of the iron compounds and place them in the correctly label beaker. Add 50 mL of distilled water to each for a final concentration of about 80 mM each. Swirl or stir to dissolve but don't cross-contaminate your stir rod. Note the colors of the solutions, if any, in your notebook. 2. Label the Erlenmeyer flasks as "FeCl2", "FeCl3", "tap water", and "distilled water". The "distilled water" sample is your "negative control". Place about 20 ml of each solution in the correctly labeled flask. Clean your graduated cylinder with distilled water between uses. 3. Place 3 drops of 1 molar NaOH solution into each the flask and record your observations. This will give you a final NaOH concentration of about 10 mM, which is less than the concentration of the two iron stock solutions that you made. Add more drops of NaOH to each solution to see whether more reaction occurs. Record your observations and the number of drops at which nothing further seems to happen. 4. Typically, a precipitate is separated from solution by filtering it for closer inspection. You will do a crude home filtering process. Label three paper towels as "FeCl2 precipitate", "FeCl3 precipitate", and "tap water precipitate". Place one paper towel over the drain and pour its solution onto the portion of the paper towel that is directly over the drain. If you pour the solution onto your stainless steel sink, you may be scrubbing stain out of your sink when the experiment is over. After it has drained, remove the towel and allow it to dry. Repeat the process with the other two precipitates. Part B: Iron Oxidation 1. Stopper the flask labeled "FeCl2 stock" solution. Shake vigorously for 30 seconds. A change may or may not occur. If a change has not occurred, then un-stopper the flask for a moment to allow more atmospheric gases (including oxygen) to enter the flask. Swirl for a few seconds. re-stopper the flask and shake vigorously until a change occurs. Record your observations. Repeat with the "FeCl3 stock" solution and record your observations.
VI. Report Questions 1. Summarize all of your observations from Part A. 2. 3. What conclusions can you draw about the iron ions in your tap water? Use all of your data to prove your conclusion. Was it helpful to collect the precipitates? When soluble ions react to form insoluble compounds, they do so by forming neutral compounds. Write the chemical reaction that describes the precipitation reaction for both the FeCl2 and FeCl3 solutions. Indicate the colors of the reactants and products, keeping in mind that chloride ion is colorless. Summarize your conclusions for Part B. chemical reaction for the reaction. Which iron solution changed color? Write the
3. Models of Covalent Structures
I. Objectives To visualize atomic bonds in covalent molecules To predict and classify bond order in covalent molecules To draw structural formulas and/or Lewis dot structures for selected models To draw all of the isomers of selected formulas II. Safety Considerations Wear eye protection at all times. The various parts of the model kit are a choking hazard for small children III. Materials From Your Home Provided for you Organic Chem Model Set
IV. Facts to Know 1. Covalent Bonds A covalent bond is formed between two atoms when the electron clouds of the two atoms overlap. The electrons become localized in the area of cloud overlap and we call this a covalent bond. The covalent bond energy results from the electrostatic attraction between the positively charged nucleus of each atom with the negatively charged electron cloud of the other atom. Quantum mechanical theory developed in 1926 allows us to predict the energy and properties of the covalent bond.
Electron Cloud Overlap between two atoms
The strength of the covalent bond is equal to or stronger than an ionic bond, which is also electrostatic in nature. Covalent bonds form between two nonmetal atoms whereas ionic bonds form between a metal ion and a nonmetal ion. In 1903, Gilbert Newton Lewis realized that the number of bonds that an atom can make is determined by noting how many electrons are in its outermost electron shell. He realized this while summarizing the information in preparation for his introductory chemistry students at the University of California-Berkeley. He worked with his ideas for 13 more years before he published them in 1916. Lewis' electron dot structures were adopted immediately by chemists everywhere as the best way to explain covalent and ionic bonding. They strongly influenced Linus Pauling, who later won a Nobel Prize in Chemistry for his research on the nature of the chemical bond. As you will learn in today's experiment, the number of bonds that each atom makes also influences the overall structure and properties of the molecule. 2. Bond Order and Rigidity The most common type of covalent bond is the single bond. Single bonds form between two atoms that share two electrons. Double bonds form when two atoms share two pairs of electrons. Triple bonds form when two atoms share three pairs of electrons. Each single bond is rigid in length but the atoms on either side of a single bond can rotate about the bond's axis. These rotated structures are called conformations because they differ only in the angles between various parts of the molecule. Double and triple bonds do not rotate very easily and add rigidity to the molecule. The molecules that you will make using the model kits demonstrate
these principles very well. 3. Bond Resonance There are molecules for which it is possible to draw multiple equivalent Lewis dot structures. These are called resonance structures as though the bonds are able to shift rapidly between the two possibilities. When you use your model kit, you will find that you can make the two structures but not the hybrid of them. Benzene is a beautiful example of bond resonance. Your kit will only allow you to build benzene with alternately single and double bonds. In reality, benzene adopts a structure that is a hybrid and each C-C bond is halfway between being a single and a double bond. Molecular orbital theory can describe these hybrid structures effectively but the model kits cannot. 4. Organic Model Building Kits and Drawing Structural Formulas Every kit includes a variety of atoms and connectors. The connectors represent the bonds and fit into the holes of the atoms. The atoms are different colors, which are assigned traditionally. The atom with four holes represents carbon and is black because charcoal is made of carbon. The atom with four holes (but you will only use three) represents nitrogen and is blue because the sky is blue. Nitrogen makes up most of the atmosphere but isn't very reactive. Cool blue nitrogen. The atom with two holes represents oxygen and is red because oxygen is so reactive. Red hot oxygen. The atom with one hole represents the halogens and is green or yellow because diatomic halogens are colored and because green and yellow haven't been used yet. Finally, the small white atoms represent hydrogens and they only one hole for bonding. The four holes in the black atoms are oriented 109.4 from one another. This ensures that all four atoms that are attached to to it are as far from one another as possible. It also means that the four atoms attached to the central atom form a tetrahedron if you were to connect them together with imaginary lines. Note the relationship between the tetrahedron and the methane molecule. When you draw the "structural formula" for molecules, you are representing three dimensions in two dimensions. Note how the bonds from the methane carbon are drawn to its hydrogens. Two are sticks, one is a solid wedge, and one is a dashed wedge. The stick bonds mean that those bonds are in the plane of the paper, the solid wedge bond is coming forward from the paper, and the dashed wedge is going back below the paper.
H C H H H
H When you connect multiple carbons together, H H H H you will have many solid and dashed wedges. Here C C is the structural formula for ethane as an example. C C The final thing to remember about drawing H H H H structural formulae is that you need to show the H Ethene geometry of the atoms that have multiple bonds. For (ethylene) Ethane instance, ethane is drawn such that the three atoms attached to each carbon are 120 from one another. You will see this in the models that you build.
V. Procedure 1. Simple Compounds: For each of the following chemical formulas in Table 1, use the periodic table to determine whether the compound is ionic or covalent. If it is a covalent compound, use your model kit to build its structure. Then draw its structural formula and its Lewis Dot Structure. The first compound has been completed for you. Table 1: Simple Compounds Formula Bond Type Covalent H2 F2 H2O H2O2 H3N KCl KOH CH4 CBr4 C2H6 C3H8 C2H6O C2H6O (isomer) Structural Formula H--H Lewis Dot Structure H:H
2. More Complex Compounds: Use your model kit to build each molecule in Table 2. Then complete the table. Table 2: More Complex Compounds Formula IUPAC Name Structural Formula CO2 C2H4 C2H3Cl chloroethene 1,1-dichloroethene cis-1,2-dichloroethene trans-1,2-dichloroethene CH2O C3H4 C4H10 C4H10 (isomer) n-pentane C5H12 C5H12 2-methylbutane C5H12 2,2-dimethylpropane propene C6H6 [draw both resonance structures] cyclopropane cyclobutane cyclobutene cubane [use the internet to find its formula & structure] adamantane [use the internet to find its formula & structure]
VI. Report Questions 1. How do you predict whether a bond will be ionic or covalent? 2. What is an isomer? 3. If you flip and invert your cis-1,2-dichloroethene molecule, you will find several ways to draw it. Explain how you know they are all the same molecule even though they can be drawn differently. 4. Include your results from Table 1. Be sure to leave enough room to draw each structure. 5. Include your results from Table 2. Be sure to leave enough room to draw each structure.
4. Virtual Radiation Safety
I. Objectives To teach the principles of radiation safety, which are shielding, distance, and time. To determine the relationship between a radioactive source and the distance from that source. To use a set of standards to determine an unknown reading. To demonstrate that everyday household things are radioactive but not necessarily hazardous. II. Safety Considerations (Duke University has an excellent website that describes radiation safety procedures for nurses (http://www.safety.duke.edu/RadSafety/nurses/default.asp) to use as your resource for the safety questions) III. Materials From Your Home Computer etc. Provided for you The urls
IV. Facts to Know Unfortunately, the least expensive Geiger counters cost about $250. We decided not to include one in your lab kit to keep its cost as low as possible. You will have to make due with a virtual lab that demonstrates most of the features of radiation safety. Radioactive materials emit invisible rays with energy that are strong enough to kill. But just as with any dangerous material, it is possible to use and work with radioactive materials without bringing harm to yourself. In fact, biochemists and radiology health care professionals who use radioactivity in their work all the time treat radioactive materials with respect rather than fear. They have been trained to handle the materials so that their exposure levels are almost undetectable. The three principles of radiation safety are shielding, distance, and time. These principles apply to all potentially toxic substances. The time factor is the most obvious: limit the length of time and the frequency that you are exposed to a radioactive material. The overall dose of radiation that your body receives will be minimized. Many studies have been performed on the people who work with radioactive materials and have shown that the human body can handle much higher levels than those found in the background radiation. Shielding is also easy to understand once you consider the three types of emission: alpha rays, beta rays, and gamma rays. Alpha rays are composed of two neutrons and two protons and are the heaviest of the three rays so they don't travel very far. Alpha rays can be stopped by particles in the air, a layer of dead skin, or a thin sheet of paper. Beta rays are composed of electrons. Beta rays are stopped by a sheet of plastic. Gamma rays are bundles of very high energy and can travel the furthest. They require lead shielding. Distance is probably your best protection because of the 1/r2 dependence. You will prove this principle in the lab today. This relationship indicates that for every distance that you move away from a source, the intensity of the rays decrease by the square of that distance. This effect is the same as the light intensity as you move away from a bulb. It occurs because the energy is emitted in all directions from the source but you are only located at one point away from that source. The power of radioactive rays is also their most useful medical property. For instance,
radioactive iodine-131 emits both beta and gamma rays and is used to cure thyroid cancer. Likewise, cobalt-60 emits very high-energy gamma rays that are used in radiation therapy to cure brain tumors. Finally, technetium-99m (m is for metastable) is the most widely used radioisotope in medical diagnosis. It accumulates around soft organs such as the heart, kidneys, and liver and the emitted gamma rays pass through the body's tissues so they can be recorded on film. Background radiation is unavoidable. We are continuously bombarded by background radiation because it is in the rocks below us (accounting for 15% of your annual radiation exposure) and in the radon gas that rises from those rocks (accounting for 55% of your annual radiation). About 11% of your annual exposure is from the natural radioactive isotopes such as potassium that are in your body. About 3% of your annual exposure comes from consumer and food products such as smoke detectors (americium) and bananas (potassium). In your virtual lab, you will be able to test some consumer products that are no longer on the market such as the beautiful red-orange glaze found on some ceramics that gets its color from the 3% uranium oxide in the glaze. Other famous products that are no longer on the market are the glow-in-the-dark, radium-painted dials of the 1920s to 1960s and the fluorescent uranium-containing "Vaseline" glassware.
Watch with glow-in-the-dark, radium-painted hands, popular from 1920s-1960s, Yellow green "Vaseline" fluorescent glass tumbler popular from 1890s-1930s, Red orange glazed ceramic plate, popular from 1936-1972 V. Procedure When reading the counts per minute (cpm) on a real Geiger counter, you should be aware that it reports a new reading every minute. We have sped up the virtual Geiger counter it reports a new "average cpm" every 10 seconds. Nevertheless, there is inherent noise in the signal and there is a delay between each reading. When you move your Geiger counter, you will have to wait for a new number to flash two different times. The first number won't be accurate because it is an average of time spent moving. Part A: Everyday Sources & Shielding 1. Sequentially place each object in front of the Geiger counter at a distance of 5 cm for 60 seconds and record the readings. Do this three times for each material and average the value. 2. Place "nothing" in front of the Geiger counter and record the background signal three times. 3. Perform a shield test to determine the type of ray emitted from the most intense sample. Choose the material with the highest reading _________. Measure its radioactivity by placing a different shielding material (paper, plastic or lead) in front of the Geiger counter and take three readings. Source Reading 1 Reading 2 Reading 3 Average Reading Avg - Bkgd
Plate Watch Background Shielding Material Paper Plastic Lead Reading 1 Reading 2 Reading 3 Average Reading Avg - Bkgd
Part B: The Inverse Square Law 1. Place the strongest emitting material in front of the Geiger counter. Vary the distance between the material and the Geiger counter (some suggestions are provided below) and record the counts per minute. Repeat three times and calculate the average. 2. Place "nothing' in front of the Geiger counter and record the background radiation three times. Distance (cm) Reading 1 Reading 2 Reading 3 Average Avg - Bkgd Reading 2 cm 3 cm 5 cm 10 cm 20 cm Background 1/r2 0.2500 0.1111 0.0400 0.0100 0.0025
VI. Report Questions 1. Background Radiation: Why is it necessary to subtract the background from all of your measurements before you use them in your analysis? Your Annual Average Radiation Exposure. This is for your interest only. Calculate your annual radiation dose (http://dwb.unl.edu/Teacher/NSF/C03/C03Mats/ARDC/ARDC.html). Keep in mind that the average annual exposure of 250 mrem. 2. Natural Materials and Shielding: Which of the household materials had the greatest radiation intensity? Which has the least radiation intensity? It may interest you to know that in 1981 the U.S. Food and Drug Administration ruled that the radioactivity levels in the red orange uraniumglazed plates was not a health hazard. The levels are well below acceptable limits of occasional exposure. 3. Natural Materials and Shielding: The simplest way to determine the type of ray emitted by a radioactive sample is to know the type of shielding that reduces its intensity. What type of ray is emitted by the household material with greatest radiation intensity? 6. The Inverse Square Law: There are several ways to prove that your data conforms to the inverse square law. The simplest but slightly inaccurate way is to plot your data (average minus background) versus "1/r2" and then draw the best straight line through your data. According to this method, you are treating 1/r2 as a constant. Mathematically, that means your line conforms to the equation: I = Io (1/r2), where I is the intensity of your sample at distance r and Io is a constant equal to the intensity of the sample when the detector is right next to the sample. Your data should give a straight line that passes through the origin (0,0) if it conforms to the inverse square law. Does it? You can determine Io from the slope of the line (slope = (y2 y1)/(x2 x1), where x1, y1 and x2, y2 are any two points on the line. 7. The Inverse Square Law: What happens to the intensity of the beta activity when the distance between the Geiger counter and the beta source is four times as great as the initial distance?
Inverse Square Dependence 400 300 200 100 0 0.00
5. "Me & Isaac Newton"
I. Objectives To learn what motivates scientists To learn who developed the best chemotherapy for childhood leukemia II. Safety Considerations Do not sit too closely to the television screen because it emits a low level of x-rays III. Materials From Your Home Provided for you DVD "Me & Isaac Newton" (1999)
IV. Facts to Know The documentary "Me & Isaac Newton" was released in 1999 and was directed by Michael Apted. Apted is a gifted director of both movies and documentaries. His movies include "Coal Miner's Daughter" in 1980, "Gorillas in the Mist" in 1988, "Nell" 1994, in and most recently "Enigma" in 2002. He began his "7 Up" documentary series in 1964 for British television. Every seven years, he revisits the same dozen or so people that were 7 years old in 1964 to discover how they've fared since the last visit. The premise is that you can predict how someone will turn out by asking them at age 7 "What are you going to be when you grow up?" In 1997, Apted directed a documentary titled "Inspirations" in which he interviewed 7 artists from 7 fields to find out what inspires them to do art. He followed this in 1999 with "Me & Isaac Newton". The point of this documentary is to find out what motivates 7 scientists to do what they do. This documentary includes Gertrude Elion, who won the Nobel prize for her work with anti-viral medicines. In fact, the class of drugs she developed interfere with DNA synthesis and have also proven useful for curing many cancers, including most childhood leukemia. Her drugs are used against herpes, AIDS, and malaria. She died in 1999 just before the film was released and the movie is dedicated to her. The seven scientists in "Me & Isaac Newton" are: Gertrude Elion, pharmaceutical chemistry Ashok Gadgil, water purification Michio Kaku, string theory Maja Mataric, robotics Steven Pinker, language disorders Karol Sikora, gene therapy Patricia Wright, lemurs V. Procedure Read the report questions. Watch the film while taking notes and with the report questions in front of you.
VI. Report Questions 1. How did Gertrude Elion become interested in science? 2. What academic credentials did Elion have when she started her first scientific job? 3. What is the name of the pharmaceutical company that Elion worked for? 4. What was the name of Elion's collaborator and co-winner of the Nobel prize? 5. Why was Elion so eager to develop drugs against cancer? 6. What was Elion's first Eureka moment? 7. What does Elion have to say about taking risks as a scientist? 8. Why is Elion optimistic about the future of science? 9. What is the chemical structure of the anti-leukemia drug that Drs. Elion & her collaborator developed? It is not mentioned in the movie. The internet has several excellent sites that describe Elion's work. Once you know the official name of a molecule, you can obtain its structure in electronic format in several ways including a visit to www.chemfinder.com. Place the electronic structure in your report. Be sure to cite the source of your structure in your report. The proper way to cite a webpage: "Title of Webpage" by Author_Names (Date page written) Page Host Name, date page accessed If you cannot find this information for your webpage, it may not be a reliable source. The proper way to cite an entire book is: Title of Book by Author_Names (Year) page xx, Publisher_Name, Publisher_location 10. What is the LD50, "dose that kills 50% of the population", for the anti-leukemic compound? The Merck Index is a good source for LD50s and so are the Material Safety Data Sheets that are prepared for every compound (www.asu.edu/lib/noble/chem/msda.htm).
You can learn about the symptoms, diagnoses, and treatments for childhood leukemia, which often strikes before age three, from the American Cancer Society at www.cancer.org.
6. Qualitative Analysis of Organic Compounds
I. Objectives To determine unknown molecules using a battery of qualitative organic analytical tests To collect a data set and interpret it to deduce the identity of several compounds To learn the chemical reactivity of some functional groups II. Safety Considerations Always wear safety goggles. NaOH is corrosive and can cause damage if it comes into contact with your skin; rinse immediately. III. Materials From Your Home Provided for you 4 Test tubes 1 Scoopula 50% ethanol in water Dimethylsulfoxide (DMSO) Phenolphthalein solution, 5 mL (0.1% in 50% ethanol) 1 M NaOH, 1 mL 2.5% FeCl3 in water, 5 mL The four unknowns are labeled a, b, c, and d and will be aspirin, acetaminophen, caffeine, and ibuprofen
IV. Facts to Know Aspirin has been the most used drug since it was first marketed by the Bayer company in 1899. Its phenomenal success launched the modern pharmaceutical era. Today, it is considered to be a member of the class of drugs called Non-Steroidal Anti-Inflammatory Drugs. NSAIDs are among the least toxic and most used members of the Painkillers. Painkillers make up the largest part of the pharmaceutical market today just as they have since apothecarists first mixed their potions.
O OH O O CH3 Aspirin
HO H N O CH3
The chemical name for aspirin is Acetylsalicylic acid. This name tells you that the molecule has a carboxylic acid functional group and that it is an ester. Salicylic acid is the common name for a benzene ring with both a carboxylic acid and an alcohol. In the case of aspirin, the alcohol of salicylic acid forms an ester with an acetyl group. Ibuprofen is marketed as Advil, Midol, Motrin, and Nuprin. It is similar to aspirin in that it has a benzene ring, a carboxylic acid. In addition, it has hydrocarbon groups. It was first marketed in the 1960s and many people prefer it to aspirin because it is less upsetting to the stomach. Otherwise, its effects are similar to those of aspirin.
Acetaminophen is marketed as Datril, Panadol, and Tylenol. It has a benzene ring, an alcohol, and an amide. When an alcohol is attached to a benzene ring, it has very different properties from an alcohol attached to an alkyl chain. Therefore, it is more accurate to say acetaminophen has a phenol functional group rather than an alcohol. It was first marketed in 1956. It is less irritating to the stomach than Ibuprofen. It is absorbed fastest of the three so that it gives the fastest pain relief. Unfortunately, it is also the most toxic of the three and prolonged use can cause liver damage Caffeine is the stimulant in No-Doz and Exedrin. The natural sources of caffeine are coffee beans, tea leaves, coca nuts, and the cola nuts. Caffeine gives the stimulating effect when you consume products made from extracts of those natural products. Caffeine has a different structure from the three NSAIDs. It has four nitrogens, which means that it has amine functional groups. Since each of the nitrogens has three different hydrocarbons attached to them, they are called tertiary amines. The two-ringed structure with four nitrogens is actually an important family of compounds called purines that are important for DNA structure.
O H3C N N Caffeine CH3 CH3 N N
The functional group theory says that the different functional groups of an organic compound provide that compound with a "fingerprint" identity. Each functional group reacts differently with different reagents to give it a unique chemical identity. In some cases, though, the difference in reactivity is too slight to be determined in a simple qualitative test. We have chosen the four compounds above because three of them have similar biological effects that are very different from the fourth. Can the qualitative tests readily distinguish these four molecules from one another? Table I: Summary of Test Results Compound Water soluble? Aspirin Ibuprofen insoluble insoluble
DMSO Soluble? pale pink soln
yellow soln that becomes yellow ppt
Carboxylic Acid? yes yes no no
Phenol? no no yes no
soluble Acetaminophen soluble Caffeine soln = solution; ppt = precipitate
V. Procedure Test 1: Water solubility This is a test to determine whether a 1 % solution is soluble in water. The more oxygen and nitrogen atoms that a compound has compared to its alkyl carbons and halides, the more soluble the compound will be in water. The oxygen and nitrogen atoms are able to form hydrogen bonds with water and that is why they add water solubility. A very crude rule of thumb is that a compound will be water soluble if there are fewer than 4 carbons for every oxygen plus nitrogen. A complicating factor for this rule is that when carbons are in a benzene ring; they make the molecule even less soluble than when they are in an alkyl chain. Label your test tubes so you don't get them mixed up. Add 30 mg of the compound to a test tube. Add 3 mL water to make a 1% solution and agitate the tube for 1 min at room temperature. If some solid remains undissolved after this length of time, it is considered insoluble.
Being dissolved for 1 min does not mean that the solution will remain dissolved forever. Let the tube sit for several minutes to see if anything settles to the bottom of the test tube. Record whether or not anything happens in your notebook. Test 2: DMSO solubility This is a test to determine whether a 1 % solution is soluble in dimethylsulfoxide (DMSO). A wide variety of compounds are soluble in DMSO such as carbohydrates, polymers, peptides, and even many inorganic salts. Most of DMSO's ability to dissolve other compounds is attributed to: 1) its polarity, which nearly matches that of water and 2) its two methyl groups help dissolve hydrocarbons. DMSO is the most powerful of the low-cost organic solvents, is itself fully soluble in water, is compatible with nearly all instrumentation used by chemists, has very low toxicity, and has very low environmental toxicity. For these reasons, it is very common to try to dissolve waterinsoluble compounds in DMSO.
O S H3C CH3 Dimethylsulfoxide
Label your test tubes so you don't get them mixed up. Add 30 mg of the compound to a test tube. Add 3 mL DMSO to make a 1% solution and agitate the tube for 1 min at room temperature. If some solid remains undissolved after this length of time, it is considered insoluble. Being dissolved for 1 min does not mean that the solution will remain dissolved forever. Let the tube sit for several minutes to see if anything settles to the bottom of the test tube. Record whether or not anything happens in your notebook. Test 3: Base Titration Test for Acid functionality This test will determine whether your compound has an acidic functional group. The idea is to dissolve your compound in some solvent (we chose 50% ethanol because all four compounds are soluble in that), add a pH indicator, and then add a drop of base (NaOH). If your compound has a carboxylic acid, it will neutralize the base and the pH won't change much. If your compound does not have a carboxylic acid, the base will change the pH of the solution enough to cause a change in the indicator color. Dissolve about 10 mg of your unknown in 5 mL of 50% ethanol. Add 1 drop of phenolphthalein. Phenolphthalein turns pink in a basic solution. Gently shake the solution. Your solution should be clear at this point. Add one drop 1 M NaOH to your solution. If your compound lacks a carboxylic acid, the solution will be a light shade of pink that is stable for at least one minute. If your compound has an acid group, then it will take many drops of NaOH before the color turns pink. Record how many drops it took for the compound to turn pink.
Test 4: Ferric Chloride test for Phenol functionality This test will determine whether your compound has a phenol. Dissolve about 50 mg compound in 10 mL 50% ethanol solution in a test tube. Gently shake the solution. Add 5 drops of 2.5% FeCl3. If your compound has a phenol, it will react with the Fe(III) ion to create an intensely colored complex. The color is determined by the other functional groups attached to the phenol.
X + 3 HCl HO X
O Fe O O
Unknown A B C D
Test for Acid
Test for Phenol
Identity of Unknown
VI. Report Questions 1. What are the identities of the 4 compounds? State the evidence and reasoning that proves your identification for each of the compounds. 2. How does a phenol differ from an alcohol? 3. What is an ester? How does an ester differ from an amide? 4. If a student's compound tested negative for having an acidic functional group and then tested negative for a phenol, what is the compound? 5. Create a flow chart of these confirmatory tests that would allow someone to easily deduce which of these compounds was their unknown. See the flow chart in Experiment 2 for an example.
7. Vapor Density of Butane
I. Objectives To determine the vapor density of butane To determine the accuracy of a measurement by calculating its percent error
II. Safety Considerations Always wear safety goggles. Liquid butane and butane gas is flammable; keep away from open flames. III. Materials From Your Home Kitchen Sink (or Small Washing Tub) Butane lighter Provided for you Test tube, large Graduated cylinder, 50 mL Sharpie Thermometer
IV. Facts to Know In 1911, Dr. O. Walter Snelling at the US Bureau of Mines discovered how to compress propane, butane, and other hydrocarbons so that they were liquid at room temperature. This started the Liquid Propane industry that remains strong even today. In 1933, many manufacturers began producing liquid butane cigarette lighters after someone realized that the liquid butane was under enough pressure that its release could be easily controlled with a well-sealed valve. Butane is also so volatile that a flame can be started with just a few sparks. These lighters remain popular because they don't require a flint or a wick or oil.
Savinelli's "Old Boy" Butane Lighter, 1933
Butane is the name for any four-carbon hydrocarbon. There are two butane isomers: n-butane, and iso-butane. Isobutane has a lower boiling point than n-butane (-11.5 C versus 0.5 C), which makes it more suitable for use at room temperature. The typical composition of the liquid in a butane lighter is 95% isobutane, 4% n-butane, and 1% of propane. The ideal gas law, PV = nRT, relates pressure and volume with temperature and moles for most gases that aren't subjected to the extremes of pressure or temperature. In this experiment, the ideal gas law will be used to determine the moles of isobutane gas released from a butane lighter. You will also measure the mass of that released butane gas. The mass will be divided by the moles to yield the isobutane molar mass (or molecular weight). This will be compared to the molar mass that you calculate using its molecular formula of C4H10. The volume of the gas will be measured by gas displacement just as you did in experiment 1 except on a larger scale in a flask. The pressure in the flask is a combination of both the pressure from the water vapor and the butane gas. Since the water pressure is determined by the atmospheric pressure, which changes with the weather, you will be given that day's barometric pressure to use in your calculation. The temperature of the gas can be assumed to be that of the water. The mass can be found by the difference in the mass of the lighter before the experiment and after the experiment.
1. Prepare to collect the gas. Fill a clean sink about 2/3 full with water. 2. Gas temperature. Measure and record the temperature of the water bath, which we'll assume is the same as the gas __________ C. You can convert the temperature to absolute temperature units of Kelvin by adding 273.15 to it: T(K) = T(C) + 273.15. 3. Set the butane lighter flame to high. 4. Fill a large test tube with water. Place your finger or thumb on the opening to prevent the water from coming out, invert the test tube and submerge it almost completely water. If some air bubbles are in your test tube, you should try again. Once it is in the water with no bubbles in it, it is OK to remove your finger or thumb. 5. Collect some gas from the lighter. Holding the test tube nearly vertically, lower the butane lighter into the water and put the opening of the lighter beneath the mouth of the test tube. Press the lever on the lighter and collect some gas into the test tube. The experiment must be restarted if butane bubbles escape. Continue to hold down the lever until the test tube is nearly full but not overflowing. One of the biggest sources of error in this experiment is that butane is quite soluble water. This makes it impossible to collect 100% of the butane gas. 6. Measure the volume of butane gas in the inverted test tube, part 1. Keep the test tube submerged but adjust it so that the water level inside is the same level as the outside water. Under these conditions, the water pressure is the same inside and outside the test tube. Use your Sharpie to mark a line on the tube to record the water level. 7. Measure the volume of butane gas in the inverted test tube, part 2. Remove the test tube from the water and release the butane gas to the atmosphere. Fill the test tube with water to the pen mark. To determine the volume of water as accurately as possible, decant the water into a 50 mL graduated cylinder. _____________ mL 8. Pressure. Obtain the barometric pressure for your address by visiting www.weather.com and entering your zip code in the "Local Weather" area. The pressure will be given in inches of Mercury. You will convert this to different units of pressure as described in the Report Questions section. 9. Repeat the experiment two additional times.
VI. Report Questions 1. Partial Pressure of the Butane Gas. According the Law of Partial Pressures, the total pressure of gas in the flask is the sum of the vapor pressure of water in the flask and the pressure of butane. To determine the pressure of the butane in the flask, you need to do the following calculation for each trial:
P(butane) = P(atmosphere) P(water) = ____________ atm (the units should be atm),
where P(atmosphere) is that day's atmospheric pressure and P(water) is the known quantity taken from the table below. Like all liquids, the vapor pressure of water is a function of its temperature. Because water is so important, its pressure as a function of temperature has been measured to high precision. You will have to convert into atm before you can use it in the equation above. Temperature (C) 16.00 17.00 18.00 19.00 20.00 21.00 22.00 Water Vapor Pressure as Function of Temperature Pressure (kPa) Temperature (C) Pressure (kPa) 1.818 23.00 2.810 1.938 24.00 2.985 2.064 25.00 3.169 2.198 26.00 3.363 2.339 27.00 3.567 2.488 28.00 3.782 2.645 29.00 4.001
There are many units of pressure. The most common are atm (atmospheres), bar, mm Hg (millimeters of mercury), psi (pound-force per square inch), Pa (pascal), and torr. In chemistry, you usually want atm or Pa so here are the conversion factors into Pa. Pressure in atm x 101,325 = pressure in pascal (Pa) Pressure in bar x 100,000 = pressure in pascal (Pa) Pressure in mmHg x 133.32 = pressure in pascal (Pa) Pressure in psi x 6894.8 = pressure in pascal (Pa) Pressure in torr x 133.32 = pressure in pascal (Pa) To convert from inches to millimeters, you need to know that there are 2.54 mm per inch. 2. Calculate the density of butane gas for each of your three trials and then report the average. A. For each trial, use your data to calculate the moles of butane. To do this, you will solve the PV=nRT equation for n. This will give you n = PV/RT, where n is in moles. You determined the pressure P in atm above. Now, you need the volume V of the butane gas in liters, the temperature T of the butane gas (actually the water) in Kelvin, and the gas constant R. The gas constant is 0.08206 atmL/molK. B. For each trial, divide the moles butane by the molar mass of butane to get the mass of butane that you collected. Convert it to mg. C. For each trial, divide the mg of butane you collected by the volume in L to get the vapor density of butane gas. D. Calculate the average vapor density from your three measurements in units of mg/L
3. % Error. Use the known isobutene vapor density of 2.46 mg/L at 21 C to calculate the percent error of your measurement.
Density (theoretical) -- Density (measured) Density (theoretical)
x 100% = _______% Error
4. Sources of Error. What are the major sources of error in the experimental design? If you were to do the experiment again, how would you change the experiment to reduce those errors?
8. Molar Volume of Hydrogen
I. Objectives To determine the standard molar volume of hydrogen gas To determine the effect of a limiting reagent on a reaction To learn about stoichiometry and molar conversions. II. Safety Considerations Always wear safety goggles. Hydrogen gas is very flammable; keep away from flames. Don't use any scented candles while doing this lab. HCl may cause severe burns; the vapors are extremely irritating to the eyes and respiratory system. Magnesium metal burns with a hot bright light that can permanently damage your eyesight. When magnesium metal is burning, it cannot be put out with water or a standard household fire extinguisher! Do not under any circumstances ignite the magnesium metal. III. Materials From Your Home Kitchen Sink (or Small Washing Tub) Provided for you Gloves Erlenmeyer flask, 250 mL Erlenmeyer flask, 125 mL Rubber stopper with hole for 125-mL flask that has glass tubing in hole and 1.5 feet of rubber tubing attached to glass tubing Magnesium metal, 0.4 g 1.0 M HCl, 18 mL Graduated cylinder, 50 mL Ruler Sharpie
IV. Facts to Know On August 27, 1783, Jacques Charles poured 250 kg of sulfuric acid over 500 kg of iron to generate enough hydrogen gas to fill a balloon that carried 9 kg of weight aloft for 45 minutes. The balloon was filled at the Paris Champs de Mars (location of the Eiffel Tower today) and traveled to the village of Gonesse (near the location of Charles de Gaulle airport today). Local villagers destroyed the balloon after it landed because they were so frightened by it. In today's experiment, you will repeat Charles' experiment except that you will substitute magnesium metal and hydrochloric acid for the iron and sulfuric acid. You will also do the reaction on a much smaller scale. The balanced equation is: Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g) Today we know that when any metal reacts with an acid, a salt MX and hydrogen gas are produced. This is actually a problem for many reasons. One of the most important reasons in industrial labs and medical hospitals is that you always want to keep strong acids away from electronic equipment because the acid will corrode the equipment. M(s) + n HX(aq) MXn(aq) + n/2 H2(g)
This is fourth experiment in which you have examined the properties of gases (see Expt 1, 3, and 7). One reason for this is that it is relatively easy to measure the volumes of gases. In today's experiment, you will determine the "standard molar volume" of hydrogen gas and compare it to the known value. One mole of any gas has a volume of 22.4141 L at 0 C (273.15K) and 1 atm. This volume is called the standard molar volume. It is independent of the gas's molar mass or density. It is amazing that every gas has the same volume under the standard condition of 0 C and 1 atm. V. Procedure (very similar to experiments 3 & 7) 1. The procedure and analysis for this experiment is very similar to Experiment 7 but is on a slightly larger scale. It is possible to do this experiment alone but if there is someone who could provide a second set of hands, it may save you some frustration to have them help you. 2. Fill your sink 2/3 full of tap water. Gather all of the materials and place them next to the sink. 3. Use your graduated cylinder to place exactly 18.0 mL 1.0 HCl into the 125-mL flask. Set it near the sink. You will need three piles of Mg metal. Divide or cut the magnesium in half (0.2 g) and then divide one half into two quarters (0.1 g each). 4. Fill the 250-mL flask with tap water and then invert it in the sink so that there is no air in flask. You can hold the flask out of the water just make sure the mouth stays below the water line so that no air gets in. If any air gets in before you are ready, you'll have to start over. 5. While holding the 250-mL flask with one hand, use the other hand: a) to place the largest pile of Mg (0.2 g) into the flask with the acid; b) to place the stopper with rubber tubing firmly onto that flask; c) to place the other end of the rubber tubing into the water so that the emerging hydrogen gas can be collected inside the inverted 250-mL flask in your other hand. 6. Collect the gas until all of the acid is used up, which is the same as saying that the Mg has stopped reacting but that there is still some Mg remaining. If all the Mg is consumed, you will need to place more Mg in the 125-mL flask containing the acid. Do this quickly. You are done when there is some Mg left in the acid flask and it is no longer reacting. 7. Adjust the 250-mL flask so that the water level inside the flask is level with the water in the sink. Mark this level with a marker. 8. You may now disassemble the system and allow the hydrogen gas to emerge. Rinse the 250-mL flask twice with tap water. Fill the 250-mL flask with tap water up to your marked level. Use your graduated cylinder as many times as necessary to determine the exact volume of the gas. 9. All liquids can be poured down the drain. The excess magnesium can safely be discarded in the trash.
VI. Report Questions 1. Moles of HCl consumed. You can calculate the moles of HCl consumed by multiplying its concentration by the volume you added to the zinc. Every mole of HCl added is equal to one mole of H+ that reacts. 2. Partial Pressure of the Hydrogen Gas. See Experiment 7 for the full explanation
P(hydrogen gas) = P(atmosphere) P(water) = ____________ atm
3. Moles of hydrogen gas produced. Use your data to calculate the moles of hydrogen gas produced. To do this, you will solve the PV=nRT equation for n. This will give you n = PV/RT, where n is in moles. You determined the pressure P in atm above. Now, you need the volume V of the hydrogen gas in liters, the temperature T of the hydrogen gas (actually the water) in Kelvin, and the gas constant R. The gas constant is 0.08206 atmL/molK. According to the reaction equation, there are two moles H+ and one mole of Mg metal consumed for every mole of H2 gas produced. 4. Standard Molar Volume of Hydrogen Gas. Once again, you will use your data and the Ideal Gas Law (PV=nRT), except that this time you will adjust your V to the standard conditions. We can use the Ideal Gas Law to adjust P, V, and T because PV/T = nR, where both n and R are constants. R is always constant. This is done by equating PV/T under one set of conditions (we'll use subscript "measured") to PV/nRT under a different set of conditions (we'll use subscript "STP" for "standard temperature and pressure"), like this:
PmeasuredVmeasured = PSTPVSTP nmeasuredTmeasured nSTPTSTP
Since we want to know the VSTP of hydrogen gas based on the data you collected, it is convenient to solve the above equation for VSTP, like this:
VSTP = (PmeasuredVmeasurednSTPTSTP)/ (PSTPnmeasuredTmeasured)
Calculate the standard volume of hydrogen gas using this equation. Remember that the subscript "measured" values are from your data and the subscript "STP" values are 273.15 K and 1 atm. 5. Compared Measured with Theoretical. How does the value that you measured for standard volume compare to the theoretical value of 22.4141 L? What is your percent error? How do you explain the difference between your value and theoretical?
9. Cheese Puff Calorimetry
I. Objectives To determine the amount of energy is food and different sources. To learn about energy transfer II. Safety Considerations Always wear safety goggles. The wire holding the food item will become warm when the food is heated in the flame Have a home fire extinguisher available (or a pitcher of water) to extinguish the flaming food item. This lab needs to be performed somewhere with excellent ventilation, such as a porch or room with a large open window. The food gives off a large flame and some sooty smoke that could set off your fire alarm. III. Materials From Your Home Matches Cold water from the refrigerator Two or three samples to burn. Examples include: cheese puff, peanut, marshmallow, cashew, kernel of buttered popped corn, kernel of unbuttered popped corn, potato chip, M&M, chocolate chip Nutritional label for the items you will burn. Needle or wire or nail or paperclip to impale the food Block of wood or cardboard to place beneath burning food Provided for you Test tube, large Test tube holder Thermometer Graduated cylinder, 50 mL
IV. Facts to Know The empirical formula for carbohydrates is CH2O. This is because sugars such as glucose and fructose have the molecular formula of C6H12O6. Since starch and complex carbohydrates make up the mass bulk of many vegetables, it is common to use CH2O as their empirical formula. The empirical formula for fats and oils is CH2. Highly unsaturated fats are somewhat different but CH2 is a good back-of-the-envelope approximation for most fats and oils. We'll ignore protein-rich foods in this lab because they tend to have a high water content, have nitrogen, and their empirical formula doesn't work out to nice round numbers. Their empirical formula is C3H7NO2 When you burn food with a flame or when you digest foods, the food molecules react with oxygen gas to create carbon dioxide and water. Food that is rich in carbohydrates will consume less oxygen gas than one that is rich in fats or oils. The energy in food comes from the net energy released for every C--C and C--H bonds that are broken to create the C=O and O--H. That is why food molecules with oxygen in them such as carbohydrates, generate less energy production per gram of food (about 16 kJ/gram = 4 Cal/g) than fatty foods (37 kJ/gram = 9 Cal/g). Here are the reaction equations using the empirical formulas. CH2O + O2 CO2 + H2O CH2 + 1.5 O2 CO2 + H2O
Today, you are going to measure the number of calories in various junk foods by burning them and compare that number against the value on the nutritional label of the food itself. The gist of the experiment is to set the junk food item on fire and to capture the heat released by warming up some water. The amount of heat absorbed by the water is calculated by knowing three things: the mass of the water, the change in temperature of the water, and the specific heat capacity of the water. The specific heat capacity of water at room temperature is 4.184 J/gK. This value changes somewhat with temperature but we try to get around that problem by starting the cooling the reation to below room temperature and heating it to above room temperature. Now, you just need to measure the other two things very carefully. The heat equation is written like this: q = mcT, where q is the heat absorbed or released in Joules, m is the mass of water in grams, c is the specific heat capacity of water, and T is the change in temperature in Kelvin or C. Food calories are still reported in Calories with a capital C. This is equal to 1000 calories of energy. These are related to Joules of energy by the following equation: Energy (J) = 4.184 J/cal x 1000 cal/Cal x Energy (Calories). V. Procedure 1. Construct your food-holding device. There are at least three considerations to building your food-impaling device. One is that the paperclip or nail will get hot while the food burns, making it hard to hold. The second is that nearly all of the foods will drip at least one drop of hot grease, possibly leaving a stain. The third consideration is that the device needs to be stable so that it doesn't tip over while you are burning the food item. A long thin nail hammered through a thin piece of wood or wallboard will work. Or, you could push an unfolded paper clip through some cardboard. Or, you could choose to hold the unfolded paperclip with a pot holder and place a paper towel beneath it to catch the grease. If you choose to burn a potato chip, you will have to construct some device that will support the chip while allowing its flames to burn upward in as small an area as possible. 2. Prepare the Heat-Measuring Device. Use your graduate cylinder to fill a test tube with 10 ml of cold water. The experiment is most accurate if you start at 4 or 5 C and then rise above room temperature. Hold the test tube with the test tube holder. Record the water's initial temperature and leave the thermometer in the test tube. Initial Water Temperature _______________ C 3. Move to a well-ventilated area. Wear your goggles. Have the extinguisher ready just in case. 4. Choose and Burn the Food Item. Choose a food item of average size. It is important that it is average because you will use the information from the nutritional label to estimate its mass. Impale your first food item onto the device you constructed. Light the food on fire with a match or lighter. Immediately place the burning food about 1 cm below the test tube. If the flame extinguishes before the food item has completely burned, relight it as quickly as possible. Some food items may need to be relit several times. Each food item will burn for a different length of time. After the food has been completely burned, slightly stir the water with the thermometer and record its final temperature. Record how many times you lit the food item, the size and color of its flame, whether you were able to capture all of the heat from the flame with your test tube of water, whether there was any grease emitted, whether there is any ash remaining, etc. Final Water Temperature _______________ C 5. Normally, you would have recorded the actual mass of the food item at the beginning of the experiment and the burned ash at the end of the experiment. Since you do not have an accurate enough balance, you will have to rely on the nutritional label to estimate the mass of your food
item without accounting for the ash. 5. Burn two different foods twice. Use a fresh water source each time. Measurement Food 1, 1st trial Food 1, 2nd Trial Food item 1 ________ Mass of piece of food (g) Mass of water in test tube (g) Initial temperature of water (C) Final temperature of water (C) Change in temperature (C) Measurement Food 2, 1st trial Food 2, 2nd Trial Food item 2 ________ Mass of piece of food (g) Mass of ash after burning (g) Mass of water in test tube (g) Initial temperature of water (C) Final temperature of water (C) Change in temperature (C) Remember that the density of water is 1.0 g/mL. Use the "serving size" and its mass to estimate the mass of a single food item.
VI. Report Questions 1. What are the average calories [not Calories] per gram for each of your food items?
2. Use the Nutritional information on the food bag to calculate the calories per gram? How does this information compare to your information? How do you explain the discrepancy? How could you do the experiment differently to get a more accurate value?
3. What is in the ash that remained after burning?
4. When a 35.6 mg food source was burned, the heat absorbed by water was 113.4 kJ/g water and the change in temperature was 26.2C. What was the mass of the water?
10. Cost of an Aluminum Atom
I. Objectives To work with Avogadro's number and molar mass To calculate the price per atom of Aluminum II. Safety Considerations No real safety considerations and no real need to wear the safety goggles. You will be handling a box of aluminum foil. Be wary of cutting yourself on the box. III. Materials From Your Home Box of aluminum foil, and its cost Calculator Ruler Provided for you
IV. Facts to Know Aluminum is a lightweight metal that is used for many commercial products including aluminum foil. In the home, aluminum foil is used for wrapping items prior to storage and for covering items during cooking. Because it is a metal, it forms a strong vapor barrier that prevents moisture loss during freezing or cooking. On the other hand, its metallic nature makes in unsuitable for other uses in the kitchen. For instance, like any metal, it will react with acidic foods to release hydrogen gas and form aluminum salts. So, you should avoid wrapping vinegar solutions, highly salted foods, or highly acidic foods such as tomatoes and lemons. You should also avoid wrapping other metallic objects such as silverware or flatware. In combination with another metal, it will produce a voltaic reaction, which would tarnish the other metallic object. Aluminum is used a lot because it is relatively cheap, doesn't tarnish very easily, and is highly recyclable. Its relatively low melting temperature makes it easy to recycle. It takes 95% less energy to recycle aluminum metal than it does to transform bauxite ore (Al2O3) into aluminum metal. Even so, you may be amazed to learn that only a small fraction of aluminum metal is recycled! Most consumers toss it in the trash to be buried in a landfill. This is the worst choice; since aluminum doesn't readily degrade and will last longer than 100 years in the landfill. The density of aluminum is 2.7 g/cm3. The formula for calculating density is mass divided by volume (D = m/v.) The formula for calculating volume is length times width times height (v = lwh.) The approximate thickness (height) of regular aluminum foil is 1.6 x 10-3 cm.
V. Procedure 1. Collect the following information and use it to determine the cost for each aluminum atom in the roll of aluminum foil. If you don't have any aluminum foil, you can gather all of this information from a box in the store. While you are there, compare several brands of aluminum foil. You will need to convert all length measurements to centimeters and the mass to grams. Brand name Cost of roll Length of roll Width of roll Total Area of roll Thickness of roll Mass of aluminum foil roll _____________________ $____________________ _____________________ cm _____________________ cm _____________________ cm2 _____________________ cm _____________________ g
(from the introduction or if given on the box)
2. Use the information above and that you have learned from your textbook to calculate the following. Moles of Al atoms/roll Number of Al atoms/roll Cost/atom of Al _____________________ _____________________
$/atom = _____________________
After you complete today's short experiment, you may wish to build your own "Aluminum Foil Deflector Beanie"(AFDB) to protect yourself from electromagnetic psychotronic mind control (EPMC). To learn how, visit: http://zapatopi.net/afdb/ If that doesn't keep you busy enough, learn why it hurts some people but not others to bite aluminum foil at http://science.howstuffworks.com/question564.htm.
VI. Report Questions 1. What is the cost per atom for the brand of aluminum foil that you chose? Show your work, use scientific notation, and calculate all numbers to three significant figures. The most points for this experiment will be awarded to those students who show the shortest route to calculating the cost per atom and who compare it with one longer route to the same answer. 2. A metal sample has a mass of 52.0 grams and a volume of 17.1 mL. Is the sample pure aluminum or an alloy of aluminum with some other metal? If it is an alloy, is the other metal heavier or lighter than aluminum? Show your work and explain your answer.
11. Household Acids, Bases, and pH
I. Objectives To classify common household items as acids, bases or neutral To become familiar with pH indicators II. Safety Considerations Always wear safety goggles. At high concentration, acids and bases will destroy tissue. Handle with care. III. Materials From Your Home Scissors Saucepan Wax paper Vinegar Household ammonia Window cleaner Red Cabbage Lemon juice Red Beet juice (not pickled) Drain declogging crystals Garden lime Alum Corn starch Table salt Baking soda Table Sugar Aspirin (If you don't have all of these materials, then substitute them with others that are different) Provided for you Red Litmus paper Blue Litmus paper 4 Dropping pipets Stirring Rod
IV. Facts to Know In this lab you will be use some common pH indicators to test various acids and bases. There are many natural pH indicators. For instance, litmus in the famous "litmus paper" is harvested from a species of lichen (a symbiont of an algae and a fungus). It comes in one of two colors, red or blue, depending on the pH.
H3C O CH3 HO HO H3C OH OH H3C O H3C N O OH H3C H3C HO N H3C O O OH H O HO H3C CHHO 3 OH N O OH
The structure of the litmus molecule is a bit daunting so you will also use the natural pH
indicator from Red Cabbage leaves. The name of its pH sensitive compound is anthocyanidin. It is green at pH 10, it is purple or blue at pH 7 and it is pink at pH 2. It also has characteristic intermediate colors at other pH's.. This allows you to determine the pH of those solutions. When all three hydroxyls on anthocyanidin have hydrogens on them (are protonated), the solution is pink. When only two are protonated the color is blue, when two are deprotonated it is green, and when all three are deprotonated it is yellow.
OH OH O O
The following is a chart that summarizes the anthocyanidin color changes in response to pH change.
V. Procedure A. Determination of pH of household materials 1. To conserve litmus paper, have a scissors handy to cut off the wet ends so that you can use the dry remaining part. Use the wax paper as your surface to protect your counter top but to also allow the various solutions to remain separate from one another. It will help to place a piece of white paper beneath the wax paper to help improve your ability to observe color changes. 2. Your standard acid will be vinegar, the standard base will be ammonia, and the neutral solution will be either tap water or distilled water. Place each of these three reagents in their own container and label a dropping pipette for each so that you don't mix them up. 3. Place a drop of vinegar (acid), a drop of water (neutral) and a drop of ammonia (base) on your wax paper. Test each with red litmus and blue litmus. Remember to cut off the wet edges so that you don't contaminate each sample. Record your observations. These are your control reactions. 4. On your wax paper place a baby pea size amount of alum, cornstarch, baking soda, cleanser, salt and aspirin (you may need to crush the aspirin with a spoon). Add a few drops of water to each and stir to dissolve. Clean the stir rod between substances to avoid contamination. Test each with both types of litmus paper and record your observations. 5. Find other common materials in your house that you'd like to test. Record your observations and add them to your report.
B. Big Red Cabbage pH Indicator 6. Take a small saucepan, place about one fourth of a shredded red cabbage into the pot, and fill with enough water to cover the leaves. Boil for 10-15 min until the water is dark purple blue. Collect the cabbage water extract and discard the leaves. You can either decant or pour it through a coffee filter. Allow the solution to cool. This is a homemade pH indicator. Keep in mind that cabbage water has a pH of around 5.5 while you are interpreting your results. It is best to keep the volume of the cabbage water drops small and the volume of the sample large. This minimizes the problem of adding acid with your pH indicator. 7. Place a few drops of vinegar, water, and ammonia on the waxed paper with a white background. Place a few drops of cabbage water on each to determine how red cabbage water responds to change in pH. You may want to experiment to find the best volume of sample to test. 8. Use the cabbage water to test garden lime, sugar, and lemon juice. Dissolve in a little water first if necessary 9. Use your cabbage water to test the baking soda, crushed aspirin, cleanser, garden lime, sugar, and lemon juice. Dissolve in water first if necessary. C. Red Beet pH Indicator 10. Open the can of non-pickled red beets. Use its juice as an acid/base indicator. The name of the molecule is betanin and it has a different structure from anthocyanidin. The colors are not as vibrant as the red cabbage so you will need to be very descriptive.
glucose O H N+ COO--
H HOOC N H COOH
11. Develop an experiment to determine the pH values that correspond to beet juice colors.
VI. Report Questions 1. How do the two types of litmus paper respond to acid, bases and neutral substances? 2. 3. 4. 5. 6. 7. Classify the alum, baking soda, cornstarch, cleanser, salt and aspirin as an acidic, basic, or neutral substance. If you have made a substitution, be sure to note that you did so. How does the cabbage water respond to change in pH? Classify garden lime, sugar and lemon juice as acidic, neutral or basic. Cabbage water will spoil after a few days. How might you make a cabbage indicator with a longer shelf life? According to your cabbage water, what is the pH of baking soda, crushed aspirin, cleanser, garden lime, sugar and lemon juice? Outline your experiment to determine the pH's for the colors of the beet juice pH indicator. Present your pH-dependence in a color chart like the one in the introduction.
12. Peroxide Content in Teeth Whiteners
I. Objectives To determine the concentration of peroxide in a teeth whitener To learn the concept of equivalence To learn how to do a titration without a buret II. Safety Considerations Always wear safety goggles. Potassium permanganate is very corrosive; it also stains everything including skin to leave a dark brown stain Sulfuric acid is a corrosive poison; at 3 molar it is also a fairly high concentration; breathing its fumes will destroy lung tissue. BE CAREFUL!!! III. Materials From Your Home Teeth Whitening gel Teaspoon measuring device Distilled water Provided for you Beaker, 250 mL Graduated cylinder, 50 mL 1 Dropping pipet, plastic 3 M sulfuric acid, 5 mL 0.1 M KMnO4, 30 mL
IV. Facts to Know The tannins from coffee, tea, and red wine leave a brown stain on teeth. There are many types of tannins but they are all polyphenols and are responsible for the antioxidant properties of all fruits, nuts, and vegetables. The tannins are brown because they have a number of conjugated double bonds that absorb light but also because they stick to each other to occlude the light. When organic material such as tannins reacts with a strong oxidant such as hydrogen peroxide, many of the bonds are broken with the result that the oxidized tannin no longer absorbs visible light and no longer appears brown.
The tannin from apple cider has a number of hydroxylated phenols connected to a glucose
The use of hydrogen peroxide as a tooth whitener was discovered in the 1970s when peroxide was being tested as the active ingredient in a mouthwash. Users reported that they thought their teeth were getting whiter. Within a few years, teeth whiteners were on the market. The fad is now exploding and the sales of teeth whiteners tripled between 2001 and 2004. Oxygen is commonly found in the 0 or -2 oxidation state. But in peroxide it is bonded differently than most and is in an oxidation state of -1, meaning that it can be oxidized or reduced. In today's experiment, you will quantify the amount of hydrogen peroxide in a teeth whitener. You will do this by using permanganate as an oxidant of the peroxide. Since permanganate is purplecolored, it provides its own indicator. When permanganate reacts with the peroxide, it transfers its
oxidizing power to the peroxide and the permanganate is transformed into Manganese ion. The Mn2+ doesn't have any color. The reaction equation is: 2MnO4- (aq) + 5 H2O2 (aq) + 6H+ 2Mn 2+ (aq) + 5O2 (g) + 8 H2O (l)
Potassium permanganate was discovered in 1659 by a German alchemist named Johann R. Glauber. Its use as an oxidizing agent dates to the early 1800s when it was used as a disinfectant. By 1910, it was used to treat the London water system. Since then, it has been used to treat water systems around the world. When you add the first few drops of permanganate to the peroxide-containing solution, the purple drop clears up as it gets mixed into the solution. This shows that the permanganate is reacting with the peroxide. At some point, though, you will add just a little bit more permanganate than there is peroxide. When that happens, the color of the solution will remain pink, or a very light purple. This is called the "equivalence point". If you record the volume of the permanganate at the equivalence point and multiply it by the concentration of the permanganate, you will know the moles of permanganate consumed. From the reaction equation, we know that this is equal to 2/5 of the moles of peroxide consumed in the reaction. The number of moles of peroxide consumed tell us the concentration of peroxide in the teeth whitener product. V. Procedure 1. Measure out 1 teaspoon of teeth whitening gel. This volume has a mass of approximately 0.5 g. Brand name______________________ 2. Completely dissolve the gel or whitening strips in 50 ml distilled water in a 250 mL beaker. Slowly add -12 drops of 3 M sulfuric acid to the solution to help the gel dissolve. As you stir the solution with a rod, the gel will slowly dissolve. This will free the hydrogen peroxide so that it is even distributed in solution.Record your observations. 3. Carefully fill the graduated cylinder with some potassium permanganate. You will use a transfer pipet to add the permanganate to your gel solution. Initial Volume of KMnO4--_______________ mL 4. Use the pipet to add some KMnO4- to the solution. . It should take about 4 mL, so you may wish to add 2 mL fairly quickly and then slow down. On the other hand, it is best to do it drop by drop or a few drops at a time. The first drops will change rapidly from purple and pink to brown to clear as it is consumed. The closer you get to the end, the longer the brownish yellow color persists. The end point is when the solution remains pink or light purple. If you add too much, it will be purplish brown. Once your titration is done place the unused permanganate back into the cylinder and record how much permanganate you used. Final Volume of KMnO4-- _______________ mL 5. Repeat the procedure two more times so that you can calculate an average value.
VI. Report Questions 1. How many moles of MnO4- reacted in each trial? 2. How many moles of H2O2 reacted in each trial? 3. How many grams of H2O2 reacted in each trial? 4. What is the average percentage of H2O2 in the bleaching gel on a gram/gram basis? 5. What is being oxidized in the titration? 6. Balance the following redox reaction:
___FeO42- + ___Cr2O7H3-7
___Fe(OH)2+ + ___CrO4-5