1 2 College Preparatory High School Chemistry Laboratory Manual and Course Curriculum 9th Edition 2014 “A day without chemistry is a day without sunshine” - Vicki Collins, Warren Wilson College, Asheville, NC “The world is not only stranger than we imagine, but it is stranger than we can imagine.” - Haldane, J.S. as per Dean Kahl, Warren Wilson College, Asheville, NC Compiled and developed by Brian Wright M.Ed. First edition 2004 9th edition 2014 For the students and faculty of Olympia High School Copyright 2014 Brian P. Wright All rights reserved Olympia High School 1302 North Street Olympia WA 98501 3 4 Table of Contents: Number 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 Section Course Policies and Procedures Safety Contract How to be Successful How to Read a Science Textbook Course Syllabus Lab Notebook Set-up How to Write a Lab Report How to Write a Discussion Experiments Distillation of Wood Techniques and Measurement Emission Spectroscopy Rapid Oxidation of Metallic Fuel Inorganic Nomenclature Determination of an Empirical Formula Reaction Rate: Marble Lab Flame Test Qualitative Analysis Chemical Reactions Stoichiometry Beer’s Law Determination of Sodium Hydroxide Concentration Comparisons of Geometry and Shape Vapor Pressure of Liquids Calorimetry Boyle’s Law Quantitative Reaction of HCl and Mg Heat of Neutralization Model Project Appendix: Reference Material Page 6 9 11 12 12 13 14 15 16 21 28 36 40 47 52 56 60 67 77 83 87 96 106 111 114 117 124 130 134 5 Mr. Wright Chemistry bWright@OSD.WedNet.Edu Grading Your grade will be determined as follows: Approximate value Exams 30% Quizzes 30% Problem sets and in class assignments 5% Laboratory experiment reports 30% Notebook 5% A AB+ B BC+ 93.0% 90.0% 87.0% 82.0% 79.0% 76.0% C CD+ D 66.0% 64.0% 60.0% 52.0% Required Course Materials: • Approved safety goggles • Non-graphing Scientific calculator (TI 30 is recommended) • Notebook sized periodic table (Student Store) • Black or blue pen for all work, and a red or green pen for corrections only • Straight edge (clear plastic is best) • Mini stapler that uses standard staples. • Notebook paper • Three-ring notebook with dividers • Duplicating laboratory notebook (not spiral bound) • Interactive notebook (spiral notebook) • Calendar/Planner to record assignments and test dates • Lab manual A more detailed description of these items will be provided during the first days of class. Students are expected to bring a required materials everyday to class. Not having needed materials or not being prepared and ready to work when the bell rings is counted as a tardy. I will not loan you a calculator for use during a test or quiz. You will not be allowed to use your phone as a calculator. Laboratory work: The only person you may work with on a lab is your lab partner. You will work with this person all year long. If your lab partner drop the course or is absent you will work alone. Your lab group is composed of you and your partner and the pair working across from you. Please discuss your lab with your lab group. Each partnership must complete the lab independently from the lab group. You will not be allowed to work in the laboratory until you successfully complete the laboratory safety test. Safety rules will be enforced at all times during laboratory experiments. You will receive one correction (warning) for laboratory safety violations. On the second correction you will be required stop working on the experiment. If you are not prepared for laboratory work or not dressed appropriately for laboratory work, you will not be allowed to participate. During laboratory activities you are required to wear closed-toe shoes and clothing that completely covers your legs. Attendance: This is a college preparatory course. Ninety percent of your future success is just SHOWING UP. You will have several legitimate reasons to be away from class. Please realize when you are not here you miss out. You are expected to take responsibility for your learning. If you miss a class for any reason it is your responsibility to learn what you missed. I want you to succeed and I care about your learning. Your success in this class is dependent on your consistent attendance. When the bell 6 rings to begin class I expect you to be in your seat and ready to work. If you are not in your seat you will be counted tardy. My attendance policy is consistent with the policy found in your student handbook. If you have any questions regarding this I am happy to discuss it with you. Policy for making up missed work and turning in late assignments: If you are going to miss class please inform me with an email. Your opportunity to make up work is dependent on written approval. In general most missed work is due on the next Friday. For any planned absence work that is due during the absence must be turned in before your departure. Laboratories, quizzes, and exams cannot be made up unless prior arrangements in writing have been made. If you make prior arrangements and miss laboratory work you have one week or until the materials are put away to complete your work. In many cases it is not possible to make up laboratory work. The instructor may be able to provide you data so that you can complete the laboratory report. If you miss an exam you have one week to make it up or until corrected exam are returned. You can drop up to three quiz grades from your final grade. Quiz make-ups are not generally given therefore, if you miss a quiz it will count as one of your three dropped quizzes. Making arrangements to take a quiz before your absence is possible. Make-up quizzes are not generally given. (If there are fewer then 10 quizzes then only 2 will be dropped) Classroom Behavior: Each student is expected to behave in a manner that enhances the learning process. Inappropriate behavior will be dealt with following the Olympia High School behavioral management program, as found in your student handbook. During the first week of class we will discuss specific classroom expectations. Each student is expected to be in class on time and prepared to learn every day. Academic/Scientific Honesty: You will often work with other students in this class. However the work you turn in must be completely yours. No part of your work can be copied from another student. Your work will be compared to other students work, including your lab partner. If any portion of your work does not appear to be original it will be deemed to be a violation of the schools policy regarding cheating and plagiarism. If you are working in a “lab group” each student’s participation must be acknowledged. You are only allowed to conduct laboratory experiments with your assigned lab partner. If your partner is absent then you will work alone, you may not work with someone else without permission in writing. If your partner is not present to conduct the lab you may not share your data with them without your instructor’s permission. Classroom Procedures: At the start of the period You are required to be in your seat and ready to work when the bell rings. If you are not, you will be marked tardy. Upon entering the classroom, collect graded work from your period’s box. Unless otherwise stated, all work (laboratory reports, homework, etc.) is due at the beginning of class. Backpacks and bags are not allowed to sit on the floor. Hooks are conveniently placed on desks to hang your backpacks and bags. During the period Do not ask to leave the class for personal reasons. Please remain in your seat during class time unless instructed otherwise. Raise your hand if you wish to ask a question or add a comment. No 7 use of personal electronic devices such as MP3 players or cell phones is allowed in the classroom and any electronics must remain in your backpack or bag. Permission must be explicitly granted to text, record sound or images. Calculators are the only approved electronic devices for classroom use. Do not consume gum, food, or drink nor apply cosmetics in the classroom, it is not safe because this is also a science laboratory. At the end of the period Students will stay in their assigned seats until the class is dismissed. Your instructor will dismiss the class. The bell does not dismiss the class. Explanation of Evaluated Course Work: Your assignments are a reflection of you, your commitment to quality and your interest in the class. All assignments will be turned in on flat, smooth paper with no tears. Notebook paper will not have spiral notebook fuzz. All assignments are to be done in ink, blue or black only. Assignments should have your name, your class ID number, and a heading which includes the date, title of the assignment, and your period. Exams: will cover material presented in the current unit. Unless otherwise stated the only materials students may use to complete midterm exams are a calculator and a provided periodic table. Most exams will have a practical laboratory component. Quizzes You will be allowed to drop your two worst quizzes (unless there is 10 or more quizzes then worst three will be dropped). Some quizzes will be given without prior planning. Problem sets will be assigned frequently. They will be checked in class on the following day. They may or may not be collected for credit. Laboratory reports are due the day after the lab is completed. The only exception is if the following day is a test or quiz day. Then the lab will be due on the day after the test or quiz. A complete description of how to write a lab report will be handed out to each student. Notebooks: You will keep a notebook, which will contain all handouts and returned work. Periodically and without warning you will be asked to produce three items from your notebook. Your notebook is for your own use. I would suggest that you organize it in a manner that allows you to find useful information quickly. A possible scheme for organizing your notebook follows. § Section 1: Course syllabus and laboratory report guidelines at the front of the notebook. § Section 2: Graded exams, quizzes, and problem sets in chronological order. § Section 3: Class notes and handouts in chronological order. § Section 4: Laboratory instruction handouts in chronological order. Interactive Notebook: your Interactive notebook (INB) will be kept in a spiral notebook. This will be a place to keep your class notes, and other work from the class. The INB has a fairly strict format that will be described in detail. Final exams will be cumulative. All guidelines that applied to the midterm exams will also apply to the final exam. This constitutes one of your first assignments. Please read, understand, and sign this syllabus, then have Mr. Wright check it off. Thank you, Brian P. Wright NBCT, M.Ed Student Signature_____________________________________ 8 Parent or Guardian Signature_____________________________ Student Laboratory Safety Contract Purpose Science is a hands-on laboratory class. You will be doing many laboratory activities which require the use of hazardous chemicals. Safety in the classroom is the #1 priority for students, teachers and parents. To ensure a safe science classroom, a list of rules has been developed and provided to you in this student safety contract. They must be followed at all times. Two copies of the contract are provided. One copy is to be signed by both you and a parent/guardian. The other is to be kept in your science notebook as a constant reminder of safety rules. 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. Labels and equipment instructions must be read carefully before use. Set up and use the prescribed apparatus as directed in the laboratory instructions or by your instructor. 15. Keep hands away from face, eyes, mouth, and body while using chemicals or preserved specimen. Wash your hands with soap and water after performing all experiments. Clean (with detergent), rinse, and wipe dry all work surfaces (including the sink) and apparatuses at the end of the experiment. Return all equipment clean and in working order to the proper storage area. 16. Experiments must be personally monitored at all times. You will be assigned a laboratory station at which to work. Do not wander around the room, distract other students, or interfere with the laboratory experiments of others. 17. Students are never permitted in the science storage rooms or preparation areas unless given specific permission by their instructor. 18. Know what to do if there is a fire drill during a laboratory period; containers must be closed, gas valves turned off, fume hoods turned off, and any electrical equipment turned off. 19. Handle all living organisms used in a laboratory activity in a humane manner. Preserved biological materials are to be treated with respect and disposed of properly. 20. When using knives and other sharp instruments, always carry the points and tips pointing down. Always cut away from your body. Never try to catch falling sharp instruments. Grasp sharp instruments only by the handles. General Guidelines: Conduct yourself in a responsible manner at all times in the laboratory. Follow all written and verbal instructions carefully. If you do not understand a direction or part of the procedure, ask the instructor before proceeding. Never work alone. No student may work in the laboratory without an instructor present. When first entering a science room, do not touch any equipment, chemicals, or materials in the laboratory until you are instructed to do so. Do not eat, drink or chew gum. Do not use laboratory glassware as containers for food or beverage. Perform only those experiments authorized by your instructor. Never do anything in the laboratory that is not called for in the laboratory procedures or by your instructor. Carefully follow all directions, both written and oral. Unauthorized experiments are prohibited. Be prepared for your work in the laboratory. Read all procedures thoroughly before entering the laboratory. Never fool around in the laboratory. Horseplay, practical jokes, and pranks are dangerous and prohibited. Observe good housekeeping practices. Work areas should be kept clean and tidy at all times. Bring only your laboratory instructions, worksheets, and/or reports to the lab area. Other materials (books, purses, backpacks) should be stored in the classroom area. Keep aisles clear and your chair pushed under your desk. Know the locations and operating procedures of all the safety equipment including the first aid kit, the eye wash station, safety shower, fire extinguisher, and fire blanket. Know where the fire alarm and exits are located. Always work in a well ventilated area. Use the fume hood when working with volatile substances or poisonous vapors. Never place your head into the fume hood. Be alert and proceed with caution at all times in the laboratory. Notify the instructor immediately of any unsafe conditions you observe. Dispose of all chemical waste properly. Never mix chemicals in sink drains. Sinks are to be used only for water and solutions designated by the instructor. Solid chemicals, metals, matches, filter paper, and all other insoluble materials are to be disposed of in proper waste containers, not in the sink. Check the label of all waste containers twice before adding your chemical waste to the container. Clothing 21. Any time chemicals, heat or glassware are used, students will wear laboratory goggles. There will be no exceptions! 22. Contact lenses should not be worn in the laboratory unless you have permission by your instructor. 23. Dress properly during laboratory activities. Long hair, dangling jewelry, and loose/baggy clothing are a hazard in the laboratory. Long hair must be tied back, and dangling jewelry and loose/baggy clothing secured. Shoes must completely cover the foot; no sandals are allowed. 24. Long pants are required for lab work and natural fibers are suggested. Accidents and Injuries 25. Report any accidents (spill, breakage etc.) or injury (cut, burn, etc.) to the instructor immediately, no matter how trivial it may appear. 26. If you or your lab partner are hurt immediately yell out “code one, code one” to get the instructor’s attention. 27. If a chemical should splash in your eyes or on your skin, immediately flush with running water from the eye wash station or safety shower for at least 20 minutes. Notify the instructor immediately. When mercury thermometers are broken, mercury must not be touched. Notify the instructor. Handling Chemicals 28. All chemicals in the laboratory are to be considered dangerous. Do not touch, taste, or smell any chemicals 9 32. 33. 34. 35. unless instructed to do so. The proper technique to smell 49. chemicals will be demonstrated. 50. 29. Check the label on chemical bottles twice before removing any of the contents. Take only as much as you need. 51. 30. Never return unused chemicals to their original containers. 31. Never use mouth suction to fill a pipette. Use a rubber bulb or pipette bulb. When transferring reagents from one container to another, hold the containers away from your body. Acids must be handled with extreme care. You will be shown the proper method for diluting strong acids. Always add acid to water, swirl or stir the solution, and be careful of the heat produced, particularly with sulfuric acid. 1. Handle flammable hazardous liquids over a pan to contain spills. 1. Never dispense flammable liquids anywhere near a source of flame 2. or heat. Take great care when transferring acids and other chemicals from one part of the laboratory to another. Hold them securely and walk carefully. Handling Glassware and Equipment 36. Carry glass tubing, especially long pieces, in a vertical position to minimize the likelihood of breakage or injury. 37. Never handle broken glass with your bare hands. Use a brush and a dustpan to clean up broken glass. Place broken or waste glass in the designated broken glass container. 38. Inserting and removing glass tubing from rubber stoppers can be dangerous. Always lubricate glassware (tubing, thistle tubes, thermometers, etc.) before attempting to insert it into a stopper. Always protect your hands with towels or cotton gloves when inserting glass tubing, or removing it from, a rubber stopper. If a piece of glassware becomes “frozen” in a stopper, take it to your instructor for removal. 39. Fill the wash bottles only with distilled water and use only as intended, ex. rinsing glassware, or adding water to a container. 40. When removing an electrical plug from its socket, grip the plug, not the electrical cord. Hands must be completely dry before touching an electrical switch, plug, or outlet. 41. Examine the glassware before each use. Never use chipped or cracked glassware. Never use dirty glassware. 42. Report damaged electrical equipment immediately. Look for things such as frayed cords, exposed wires, and loose connections. Do not use damaged electrical equipment. 43. If you do not understand how to use a piece of equipment, ask the instructor for help. 44. Do not immerse hot glassware in cold water, it may shatter. Never look into a container that is being heated. Do not place hot apparatuses directly on the laboratory desk. Always use an insulating pad. Allow plenty of time for hot apparatuses to cool before touching. When bending glass, allow time for the glass to cool before handling. Hot and cold glass have the same visual appearance. Determine if an object is hot by bringing the back of your hand close to it prior to grasping it. Questions Do you wear contact lenses? yes____ no_____ Are you color blind? yes_____ no_____ Do you have any allergies? yes____ no_____ If yes, please list: __________________________ ____________________________________________ Agreement I, __________________(student’s name) have read and agree to follow all of the safety rules set forth in this contract. I realize that I must obey these rules to ensure my own safety, and that of my fellow students and instructors. I will cooperate to the fullest extent with my instructor and fellow students to maintain a safe lab environment. I will always closely follow the oral and written instructions provided by the instructor. I am aware that any violation of this safety contract that results in unsafe conduct in the laboratory or misbehaviour on my part, may result in being removed from the laboratory, study hall, receiving a failing grade, and/or dismissal from the course. _______________________________ student name (print) _______________________________ student signature _______________________________ date Dear Parent or Guardian: We feel that you should be informed regarding the school’s effort to create and maintain a safe science classroom/laboratory environment. With the cooperation of the instructors, parents, and students, a safety instruction program can eliminate, prevent, and correct possible hazards. You should be aware of the safety instructions your son/daughter will receive before engaging in any laboratory work. Please read the list of safety rules above. No student will be permitted to perform laboratory activities unless this contract is signed by both student and parent/guardian and is on file with the teacher. Your signature on this contract indicates that you have read this Student Safety Contract, are aware of the measures taken to insure the safety of your student in the science laboratory, and will instruct your son/daughter to uphold his/her agreement to follow these rules and procedures in the laboratory. Heating Substances 45. Exercise extreme caution when using a gas burner. Take care that hair, clothing, and hands are a safe distance from the flame at all times. Do not put any substance into the flame unless specifically instructed to do so. Never reach over an exposed flame. Light gas (or alcohol) burners only as instructed by the teacher. 46. Never leave a lit burner unattended. Never leave anything that is being heated or is visibly reacting unattended. Always turn the burner or hot plate off when not in use. 47. You will be instructed in the proper method of heating and boiling liquids in test tubes. Do not point the open end of a test tube being heated at yourself or any one else. 48. Heated metals and glass remain hot for a long time. They should be set aside to cool and only be picked up with caution. Use tongs or heat protective gloves if necessary. _______________________________ parent/guardian _______________________________ date 10 How to Get an A in High School Chemistry 1. Come to class every day with the proper clothing, materials (calculator, class notebook, lab notebook, homework, lab manual, etc.) and a positive attitude. 2. Keep an organized planner and notebook 3. Memorize all polyatomic ions and elements that you are asked to memorize. (See labs on nomenclature and reactions) 4. Learn your naming rules. (See labs on nomenclature and reactions.) 5. Learn how to use significant figures. 6. Know how to solve for moles. 7. Know how to use dimensional analysis. 8. Keep up with your Interactive Notebook and do Left Hand Activities. 9. Make a note card for each quiz. 10. Ask questions regarding topics you do not understand. 11. Develop a study plan for each exam. 12. Turn in all assignments on time. 13. Have fun. 14. To maximize your time in the laboratory complete as much of the lab report as possible before arriving on a lab day. 15. Read the text book effectively. To do this: a. Skim the text, look at pictures, diagrams and tables b. Read side bars and text for pictures, diagrams and tables c. Read the text. If you do not understand a section, take note of this and keep on reading. d. Read the section you did not understand and reread it until you do. e. Take notes while reading. This includes paying special attention to any words that are bolded. This may include making an outline. f. Do the problems. g. Do the reading over many short sessions. h. Take reading notes. 11 ilovechem@gmail.com Topics to be covered History of science* Measurement* Physical vs chemical* Properties of phases* Precision, accuracy* Significance** Safety** Dimensional analysis** Know elements 1-38, 46-50, 53-56, 78-83, 86 87** Atomic structure* Sub particles Theories of the atom Experiments to Determine structure* Octet rule Exceptions Subatomic particle* Size location Charge Symbol Mole* Molecular weight AMU Empirical formula* Nomenclature** Acid/Bases Two non metals Metal / non metal Polyatomic Periodic table* Development Trends Atomic size Electronegativity Ionization energy Ion size Titration* Molarity* M1V1=M2V2* Solution making (Gramsà)* Limiting reactants* Reaction types* Gas forming Oxidation reduction Combustion Combination Decomposition Neutralization Acid Base* pH Electron configuration* Orbitals Geometry Noble gas configuration Trends VSEPR Electron dot config Bonding* Polar Nonpolar Ionic Lewis Structure** Resonance Octet Hybridized orbitals Formal Charge* Gas Laws* Combined Charles’ Boyle’s Kinetic theory Ideal Graham’s Dalton’s Diffusion Effusion 12 Electrochemistry Organic chemistry Functional group IUPAC naming Substitution Mechanism Stereochem Polymers Inter- vs. intra-molecular forces* Balancing redox reactions Final Project* Laboratory activities Measurement lab* Reaction Rate* Density lab* Qualitative analysis* Determination of an empirical formula* Making borate glass Copper Brass (Alloy of copper) Spectroscopy* Modeling and Lewis structure model lab* Titration* Stoichiometry* Ice (testing Hypothesis)* Heat or Reaction* Heat of fusion* Heat of neutralization* End of year project* Hydrogen rockets* *Topics covered last year **Fundamental topic that your success depends on. How to Prepare Your Notebook and Write a Lab Report (Front cover) (Title Page) Chemistry 1 Olympia High School 2012-2013 Chemistry 1 Olympia High School 2012-2013 Name Address Email Phone # Name Address Email Phone # Table of contents Date Title Description: pg: Write a Short title Describe the purpose and results of experiment. This may take several sentences. Using various glassware and laboratory techniques, the density of water will be determined. Copper coins will be zinc plated and then heated to make bronze. Using NaOH of known molarity, the concentration, in moles/liter, of an unknown diprotic acid will be determined. 4 Determination of the density of water Copper to Gold Titration of an unknown acid 8 13 15 Skip one page before the first lab entry Every page in your lab notebook should have the following information: Name_____________________ Seat no.___ Period/Section_________ Page No. _____ Partner’s Name ______________ Experiment Title________________ Date____ ______________________________________________________________________________ 13 Brian Wright Seat no. 32 Per 1 page 05 Dean Kahl Determination of the density of water 2/22/02 The above sample heading, which includes you and your partner’s names, the date, page number and the experiment being reported, must appear on every page of your lab notebook. Read and understand the background information before writing your Pre-Lab. Each lab will need the following sections. Title: This is what your experiment is called. This will often be provided, however, if another appropriate title makes more sense, you are welcome to use it. part of Pre Lab ***The “Pre-Lab”*** Purpose: This is a short description of why you are doing the lab. What theory are you testing? What do you hope to learn from the experiment? Hypothesis: What are the likely results of this experiment? Why is this the likely result? Pre-Lab Questions: These will be assigned questions. They typically will help prepare you to succeed in the lab. Write out the question and then answer it. Use sig figs, units, and labels as appropriate. Show your work and circle the answer. Materials: This is a list of all equipment and supplies that are needed to conduct the experiment. This is an excellent place to list chemical formulas and concentrations, as well as appropriate molecular weights. Trust me on the concentrations and molecular weights. You will someday wish you had written them down. Planned Procedure: YOU WILL LOSE POINTS IF YOU ATTEMPT TO COPY THE WRITTEN PROCEDURE. Instead what I would like you to do is to paraphrase the written procedure. Read and understand the experimental procedure. Then write in your own words what it is you plan to do. This is a way for me to gauge your level of understanding before you do the experiment. Please indicate the source of the actual procedure that you plan to follow. This may be written in the future tense. Please paraphrase and keep it short, about 4-5 sentences. It may need to be longer if the lab has more then one part. Procedure: Write down what you actually did. Please note any variations from the planned procedure. An acceptable procedure may state: “Followed procedure as per page ## in Laboratory Manual (Wright 2013) with the following revisions… Data: This is where you organize your data. Data can take the form of observations, which are qualitative, and measurements, which are quantitative. In all labs I expect a qualitative observation section and in many labs a quantitative data section. Any time you make a table you are required to use a straight edge. All labels and units as well as significant figures must be used when making observations. ****The only place that data is recorded is directly into your lab notebook.*** Calculations: Any time you do any mathematical operations with data you must include it in your calculations section, even if it is a simple adding or subtracting problem. It is very important to show every step. Show the correct units and make sure to use significant figures. Remember: do not round until the end, just keep track of the sig figs as you go. You must show all work that you do and must keep track of units, labels and sig figs. Label each calculation so that it is very clear as to which part of the lab the calculation is relevant. Please use dimensional analysis and or the “4 step method” as appropriate. 14 Results: This is where you can report the objective consequences of your experiment. You can report what happened. This is not where you interpret what anything means. Your results should somehow reflect the intentions stated in your purpose. Any time you manipulate data using mathematical operations you get an answer. These are your results. Report what your calculations equal. This is not a section for you to report what the data means or what the results imply. It is just a place for you to present the facts. Discussion: In every lab report I expect the following issues to be discussed: • Make a statement regarding how the data, observations and results relate to the stated purpose of the lab. And whether the results supported or refuted the stated hypothesis (do not write “it was proved…”) • What went wrong and what effect might those errors have in your results? I am looking for a well thought out treatment of how the data may be flawed, and what specific effect this would have on the results. • Compare your results with known values or theoretical values. Then determine the % error between the known value and the experimental value. • What could be done differently to improve the validity of the experimental results? • What are appropriate follow up experiments? Post Lab Questions: These will be direct questions that I may ask you. Remember to always write out the question and then answer in complete sentences. If it is a calculation, remember to use sig figs, units and labels. Guidelines to Scientific Writing The general purpose of any scientific paper is to report what happened in enough detail that other people can read your work, follow what you did, and understand your results. It needs to be clear and concise. When writing for a science class (at least Mr. Wright’s) please use the following style conventions. • Write in the past tense. • Leave out any affective comments regarding how you feel about the experiment or your work. • People do not do things in science, things happen. Except in an extraordinary circumstance leave out I, you, we, he, she, it. • Be clear and concise. • Make as few assumptions as possible. • Write using the objective voice. Avoid using colloquialisms. • Your grade depends on presentation. Be very neat, use rulers, use pens, and if you make a mistake put a single line through the mistake instead of scribbling all over it. • Never destroy your data or work; never tear out pages. • Take your time; create something you can be proud of. Do it right the first time. • Make statements that you can support with experimental evidence. • Use a tentative voice in your writing. 15 Experiment 1 Distillation of Wood Purpose: Develop skills related to writing la reports, the scientific method and lab work. This laboratory activity also pertains to the law of Conservation of Mass. Hypothesis: Before beginning the experiment, predict the forms of matter you might produce. What changes to the wood do you predict will occur? Background: Ahmed Zewail was fascinated by the world around him. His curiosity and pursuit of the answers to his questions of the behavior of everyday things ultimately led to the greatest recognition of achievement, the Nobel Prize. As a young boy he wondered how atoms and molecules behave. He asked, "What caused matter to change state; go from solid to liquid, liquid to gas or gas to liquid? Is it true that a flammable gas is present in wood?" Ahmed knew the answers lay in experimentation. Why not take wood apart, he asked. In his first experiment, he heated pieces of wood in a test tube and observed what happened. What does it feel like to be a scientist? You and your group will perform the same experiment that started Ahmed Zewail's career in science. Before beginning the experiment, predict the forms of matter you might produce. What changes to the wood do you predict will occur? As your group proceeds through the activity, you will have many questions about the changes in matter that are occurring. Record these questions for discussion at the end of the activity. In this lab you will also be comparing the mass of the material before the reaction (reactants) to the mass of the material produced in the reaction (products). To make these comparisons it is necessary to carefully mass the reactants before the reaction and then very carefully mass all of the products. To mathematically compare these you will use the equation for percent yield. Experimental Mass Theoretical Mass * 100 = Percent yield (Equation 1) Detailed description of each part of your lab report: The following laboratory report format is based on my personal work in both industrial and academic laboratories. It is also based on what will be expected from you when you take college science classes. I think the one primary criticism of the following instructions is that it leaves out any personal reflection. Many scientists use their lab notebook as a journal as well as a place to record their experiments. Reading the Background: Please make sure to read and understand all the information found in the Background. The goal is that all the information necessary to do the lab should be found in the Background. This includes formulas, constants and how to do the calculations. • Writing the Pre-Lab: 16 It is essential that you know what do in the lab before you arrive to class. To ensure this you will write the Pre-Lab. Please carefully follow the directions on how to write a lab report. If your PreLab is not completed you will not be allowed to participate in the laboratory activity. Please construct all data tables prior to coming to class. • Data Table: The data to be collected must be based on the purpose of the experiment, the hypothesis, and procedure. Make sure to have it well organized, labeled and use a straight edge. In many labs it is necessary to collect numeric data or quantitative data. However make sure that you also collect qualitative data for every experiment. • Quantitative Observations: Make sure to record all data to the correct number of significant figures. Be careful about remembering to estimate the last digit for all non-digital measuring devices. Be sure to list units and include all labels. • Qualitative Observations: When making observations, please consider the following *Sight* *Smell* *Hearing* *Time* *Manner* • Calculations: Your calculation section must be very carefully labeled. It needs to be very clear where the numbers have come from (what part of the experiment). Make sure to use sig figs, labels and units. You will be expected to use dimensional analysis and or the 4-step method for any appropriate calculations. • Results: Your results are the summarized presentation of the data you have collected. This may take the form of a graph. You do not need to explain the results. Just present them. Often times the results represent an answer to the hypothesis and/or purpose. If the question is, “What color is Amy Pond’s hair?” the answer is “Red.” You do not need to discuss red hair, or discuss if Amy is better than Rose. The results are just the answer to the question. The results of an error analysis are often presented here. • Discussion: The discussion is where you will write about what happened during the lab. There is a pretty good description of how to write the discussion in “How to Write a Lab Report.” Please remember to be succinct and to write in the past tense. Do not use any personal pronouns. • Post-Lab Questions: Most of these questions attempt to measure what you learned during the experiment. Often times the Post-Lab questions will reflect the kind of work done in the calculations section. Materials: Chemical splash goggles Chemical resistant apron wood splints 1-hole stopper 2 glass bends beaker Burner striker glycerin ring stand 3 test tubes (labeled: wood, distillate, purified distillate) Pre-Lab Questions: None in this lab. 17 boiling chips 2-hole stopper ice water rubber tubing test tube clamp Procedure: Heating the Wood 1) Wear chemical splash goggles and a chemical resistant apron. 2) Find the mass of all three test tubes and record values. 3) Place enough wood splints into the test tube (labeled wood) to fill it. 4) Break the splints in half so that they fit, and leave room for a stopper to fit securely on the test tube. 5) Connect the test tube (labeled wood) containing the splints to the stopper assembly and condensing tube (labeled distillate) as shown in Figure 1. 6) Heat the wood-filled test tube using a burner with a medium flame. Keep the flame away from the rubber stopper and test tube clamp 7) Carefully move the burner back and forth along the length of the test tube to make sure all the wood is heated. 8) Observe what happens and record your observations in the table on page 3. 9) Try lighting the gas coming from the second glass elbow. Record the observations and questions your group has generated. 10) Continue heating the wood and re-light the gas coming from the 2nd glass elbow until there are several mL of liquid in the condensing test tube. Let the gas burn as long as possible, relighting as needed. 11) When all the wood has been heated, turn off the burner. Examine the liquid (distillate) in the upright test tube in the beaker of water. Record the observations and questions your group has generated. 12) Examine the remains of the wood splints. Try burning some of this wood. Record theobservations and questions your group has gener Distillation of the Liquid 1) Place a few boiling chips in the test tube (distillate) with the liquid you have collected. 2) Use the apparatus setup shown in Figure 2 to boil off about half of the liquid. Make sure the rubber tube is all the way into the collecting test tube (purified distillate) in the beaker of ice water. 3) Record the observations and questions your group has generated. 4) In this order: remove the rubber tube from the test tube in the beaker of water, then turn off the burner. 5) Compare the liquid you collected in the test tube in the beaker to the original liquid. Do they look the same? Record the observations and questions your group has generated. 6) Try mixing the liquids in the two test tubes. Record what happens. 7) Clean up. If you do not do a good job of cleaning your test tubes now, they will not come clean later and you will be required to purchase new test tubes. 18 Data: You must always write your observations, in ink, into your laboratory notebook. This is a permanent record of the laboratory work that you do. Quantitative Data Mass of test tube (wood) full of wood splints Mass of test tube (wood) empty Mass of test tube (wood) after heating Mass of test tube (distillate) with liquid Mass of test tube (distillate) empty Data: Data# (1) (2) (3) (4) (5) Qualitative Observation When you make the data table in the laboratory notebook it may be a good idea to leave more space between lines so as to not limit the quantity of observations made. Procedure Observations Questions you may wonder Procedure Step #7 Procedure Step #8 Procedure Step #10 Procedure Step #11 Part 2 Observations Questions you may wonder Procedure Step #3 Procedure Step #5 Procedure Step #6 Calculations: Make sure to show all work and use labels. You will also be required to use sig figs (you may have to learn how to use sig figs first. It is fun so you have this to look forward to). It is also critically important that the calculations section is organized and it is completely clear and easy to understand. Each time you do any math, even simple adding and subtracting of data, it is essential that the work is shown and it is clear where the numbers (data) came from and what calculation was used to process this data. The calculations for this lab are straightforward and easy to organize. The data table has been arranged and labeled in such a manner to make labeling the calculations easier to follow. This is something that you should repeat in the future when you are making your own data tables. Please note how the equations are listed and what is being solved for is very easy to understand. 19 Mass of Wood Reactants (Data 1) - (Data 2)= (Calculation #1) Mass of Distillate Product (Data 4) - (Data 5) = (Calculation #2) Mass of Unreacted Reactant (charcoal) (Data 3) - (Data 2) = (Calculation #3) Mass of products from distillation (Calculation #2) + (Calculation #3) = (Calculation #4) Percent yield of distillation (Calculation #4) / (Calculation #1) * 100 = Results: • • • Loss of mass of reactants from the distillation reaction Percent yield for distillation Purified distillate Discussion: The Law of Conservation of Mass is a fundamental law of chemistry. Matter cannot be created or destroyed due to a chemical reaction. This means that the mass of the products of a reaction must equal the mass of the reactants. Discuss if the data in this lab supported or did not support the Law of Conservation of Mass. Discuss why it did or did not. Make sure to use the data and results to support any statements. In addition to the above please make sure to discuss all other topics in the section of your lab manual called “How to write a lab report.” Post-Lab Questions: 1. Sometimes the best way to understand how something works, or what it is made of, is to take it apart. Did you take the wood apart? 2. How many different states of matter were produced in the process? 3. What happened to the wood? Did it melt? 20 Experiment 2 Techniques and Measurements Purpose: The methods and equipment used to measure volume and mass will be studied and used to find density. Hypothesis: State a hypothesis regarding which glassware will produce the most accurate result. Background: Volume The SI unit for volume is 1 cubic meter, 1m3. This, however, is too large a volume for use in most laboratory settings. Volume may also be measured in cubic centimeters. 1cm3=milliliter. Volume is usually measured in units of liters (L) or milliliters (mL) (1000mL=1L). Volume may be measured with various degrees of accuracy depending upon the equipment used. Pipette A pipette is a glass tube, usually with a bulb in the middle and an index mark on the upper stem. It is designed to accurately deliver a specific volume of liquid. Pipettes are available in various sizes, 1.00, 2.00, 5.00, 10.00, 25.00mL. Pipettes are always to be used with a rubber bulb or pipette filler. Never use your mouth. They are accurate to the 0.01mL. ↑ Fill line is here Graduated cylinder A glass tube standing on the flared base with uniform calibration marks along the tube. The TC accuracy of a 10.00mL and 25.00 mL +/- 0.02. If the Cylinder is used to deliver a solution the accuracy drops to +/- 5% 21 Burette A burette is a calibrated glass tube with a stopcock and tip on its lower end. A burette is designed to measure how much liquid has been delivered. The amount delivered is read from the difference in the graduations on the tube. When used and read correctly, a burette is accurate to the 0.01mL. Be careful when using a burette, please note that for instance the water level may start at 1.00mL when water is drained out of the bottom it may stop at 4.50mL. In this case 3.50 mL of water would have been delivered. Volumetric Flask A flask with a narrow neck designed to contain a given volume of liquid up to an index mark. Typical capacities are 100.00mL, 200.00mL, 500.00mL, and 1000.00mL. Accuracy is 0.1%. The correct glassware for the job Burettes can be used when the volume needed is not a specific quantity, but when the actual volume delivered must be known. An initial volume reading is measured using the bottom of the meniscus (the curved surface of the liquid) against the lines on the burette. Some liquid is delivered through the stopcock. A final volume reading is measured in the same manner as the initial volume. The difference between the initial and final volumes is the exact quantity delivered. A pipette or a volumetric flask is used when a specific quantity of liquid must be accurately measured. The vessel is filled until the bottom of the meniscus is level with the calibration mark. ß24.52mL A measuring cylinder is used when an approximate volume of liquid is needed. Some beakers and Erlenmeyer flasks have graduations marked on them. These graduations are accurate to about 20% and are only useful for crude estimates. TC vs. TD Glassware that is marked TC is only accurate when measuring materials in the container. Think of TC as “To Contain.” While glassware marked TD only accurate measures volumes delivered. Think of TD as “To Deliver.” Mass 22 There are two primary kinds of tools to determine mass. First is a mechanical balance which determines the mass of an object by “balancing” against the mass of known objects. An example of this is called the triple beam balance. The other method of determining mass uses an electronic device referred to as a top loading balance, digital balance, or analytical balance. These balances use electricity and magnetism to determine the mass of an object. Triple Beam Balance It is important that before loading that it is correctly zeroed. Read the needle at eye level to avoid the effects of parallax. This balance will mass objects to approximately 600g and reads to the 0.01g. It is necessary to estimate the final digit. Electronic balance The mass of samples can be determined quickly and accurately using an electronic balance. The balance can be read to the nearest 0.01g or 0.001 g depending on the model. The last digit of a digital reading is an estimate. Before loading the balance it is essential that it reads zero. If it does not, press the “tare” button gently. Procedure for finding mass When the mass of a solid object is to be determined, the mass of an empty weighing bottle is first measured and recorded. The object is then placed into the weighing bottle and the mass of the bottle and the contents are remeasured. Again the mass is recorded. The difference in the two masses is the mass of the object. This is called weighing by difference. When the mass of a chemical substance, a reagent, is needed for an experiment, the technique of “weighing by difference” is used. The reagent (liquid, solid, or solution) is placed into a container. The total mass of the container and its contents is measured. The mass is recorded into a notebook. Some material is tipped out into a beaker or the glassware where it will be used. The mass of the container and its contents are remeasured. Again, the mass is recorded into a notebook. The difference in the two masses is the amount of material transferred to the beaker. Laboratory instructions generally say that the approximate mass of a sample is to be placed into a container and its mass determined. This means that it is not necessary to measure out exactly the specific amount stated. The amount used should be within about 0.05g of the amount stated, but the actual mass of the amount used should be measured to the nearest 0.01 or 0.001g. The mass of any object should be measured with the object at room temperature, otherwise convection currents make it impossible to obtain a stable balance reading. The sample should not be wet or damp. Evaporation makes it difficult to obtain a stable reading. Furthermore, moisture increases the apparent mass of the sample and leads to error when the degree of moisture is not known. Moisture can also lead to unwanted corrosion of the balance pan. The tolerances on the masses in the balance are extremely small; therefore a mass obtained on one balance will be very close to that found on another balance. Nevertheless, it is best to use the same balance for all measurements during the course of an experiment. 23 Density Density is defined as mass per unit volume, d=m/v. It is an important physical property of a substance. The SI unit of density is the kg/m3. These units are somewhat hard to work with in the classroom setting, thus the usual units quoted are g/mL or g/cm3. Density was one of the earliest methods of sample identification. The density of a material can give information about its identity. Significant figures The accuracy of the results is limited by the accuracy of the equipment used. The significant figures show the accuracy of the results of any experiment. The last figure in each measurement is an estimate thus it is “in doubt.” Each measurement has a range or reliability. Thus, an object that has a mass of 3.672g when measured on an electronic balance with an accuracy of 1mg has a range of 3.672g +/- .001g, which is 3.671g to 3.673g When more than one measurement is used in a calculation, the ranges of all measurements must be considered and the rules of significant figures must be used in the calculations. For example, in determination of the density of an object, the mass is found to be 6.45 +/-.01g and volume is 1.43ml +/-0.1mL. The density range is therefore 4.47g/mL to 4.55 g/mL. (6.44g/1.44mL) to (6.46g/1.42mL). To obtain the smallest possible density take the smallest mass over the largest volume. To obtain the largest possible density take the largest mass over the smallest volume. Range of values for calculated measurements When finding density there is both uncertainties in the measurement of mass and in volume. To find the range of volume it is necessary to consider what the greatest possible volume is and what the least possible volume could be. Vá = Vfinal á - Vintial â and Vâ = Vfinal â - Vintial á . Thus it is necessary to factor in the uncertainty of the denominator and the numerator. To do this it is necessary to determine what is the largest and smallest possible densities that could result. Large mass ámass Largest density = áD= Small vol âvol Small mass âmass âD= Large vol ávol Thus if the measured mass is 47.58 grams +/- 0.01 grams the largest possible mass is 47.59 and the smallest mass is 47.57 grams. If the measured volume is 28.5 mL +/- 0.5mL then the largest volume is 29.0 mL and the smallest is 28.0 mL. 47.58 g ámass 47.59 g Density= =1.67g/mL áD= = 1.70 g/mL áD = âvol 28.5 mL 28.0 mL Smallest density= âD = âmass 47.57 g âD= ávol 29.0 mL Thus the density is------------- 1.67 g/ml with a range of 1.64g/mL to 1.70 g/mL or it could be reported as ----- 1.67g/ml +/- 0.03 g/mL. Error analysis 24 = 1.64 g/mL It is possible to compare the experimental results to that of the literature value (theoretical value). This is an important calculation for it helps illustrates the accuracy of the experiment. Percent error is calculated as follows | theoretical – experimental | theoretical * 100 = Percent error Water density The density of water is the basis of many measurements. Water is often considered to have a density of 1g/ml, thus 1000g = 1000mL or 1kg=1liter. Rarely does water have a density of 1g/mL. The density of water at 25.0 oC is 0.997 g/mL. Only at 3.98 oC is water 1.00 g/mL. Pre-Lab Questions: 1. How many significant figures are in each of the following measurements? Write both the measurement and answer. a. 42.78mL f. 6,000mL b. 4.0767g g. 1010.0g c. 50.00mL h. 350mL d. 0.005g i. 100.000g e. 0.0610L j. 300ml 2. The mass of an object is measured on an electronic balance (accurate to 0.001g). Its mass is recorded as 8.432g. The volume of the object is measured using the displacement of water in a 25mL measuring cylinder. The initial volume of water was 10.10 mL (before the object was submerged). The volume after the object was placed in the cylinder was 22.40mL. a. What is the range of possible masses for the object? b. What is the range of possible values for the volume of the object? The cylinder is only accurate to +/- .05mL. c. Calculate the density of the object (to the correct number of significant figures) based upon the reported mass and volume (ignore ranges) d. Calculate the density of the object at the lowest possible mass and the highest possible volume e. Calculate the density of the object at the highest possible mass and the lowest possible volume. f. What is the range of densities based on parts d and e? g. State the density of the object and its range. 3. An iron object was determined experimentally to have a density of 7.78 g/mL. What is the percent error? Please use the Sargent Welch periodic table to find the theoretical value. Procedure: Density of water 1. Determine the mass of a clean, dry weighing container (the smallest beaker in your drawer). Use a burette to measure between 10 and 15 mL of water into the weighing container. Record the initial and final readings from the burette to the nearest 0.01 mL Calculate the actual volume of the water delivered into the weighing container to the highest number of significant figures possible. Do not attempt to start the burette at exactly 0.00mL. Also, do not attempt to measure an exact number of mL, it is far more important to know exactly what volume was 25 delivered. Determine the mass of the weighing bottle with the water in it (remember to use the same balance throughout the experiment). Use this data to calculate the density and range of water. 2. Repeat part 1 using a 10.00 mL pipette instead of a burette. Compare your results with part 1, making note of the number of significant figures obtained by measuring the volume with a pipette. There are no initial and final readings on a pipette; it can only deliver one volume. Use this data to calculate the density and range of water. 3. Repeat part 1 using a 10mL measuring cylinder instead of the burette. Compare your results with part 1. Make note of the number of significant figures obtained by measuring the volume with a measuring cylinder. When using a measuring cylinder to deliver a volume it is filled to a certain level and this level is recorded. The cylinder is then completely emptied. Do not attempt to obtain initial and final readings. Use this data to calculate the density and range of water. Do not attempt to fill cylinder to 10.00mL. Density of unknown liquid 4. Obtain 20mL of the unknown liquid in your Erlenmeyer flask. Write down the liquids ID #. Make observations regarding the liquid. Determine the mass of a clean dry weighing container. Use a pipette to place 10.00mL of the unknown liquid into the weighing container. Find the mass of the container and liquid. 5. From data, established in step 4, determine the density and range of the unknown liquid. Density of unknown object 6. Obtain a “density unknown” from your instructor. Write down the solids ID #. Make observations regarding the solid. Determine the mass for the object. Using a measuring cylinder, determine the volume of the object by measuring the volume of water that the object displaces. To do this, partially fill the cylinder with water. Record the volume. By tilting the cylinder slowly slide the density unknown into the water. BE CAREFUL NOT TO BREAK THE BOTTOM OF THE CYLINDER WITH THE OBJECT. TILT THE CYLINDER! Record the volume again. Use this data to calculate the density and range of unknown object. Data and Observations: Your data should show all masses and volume measurements with the correct units and to the correct number of significant figures. Do not forget to also include any qualitative observations. You will need data tables for all sample of water, unknown liquid and solid. Example data tables follow. Water density determined using a burette Mass of beaker 87.01g Mass of beaker and water 97.03g Mass of water 10.02g Initial burette reading 0.15mL Final burette reading 10.18mL Volume of water 10.03mL Water density determined using a pipette Mass of beaker Mass of beaker and water Mass of water Volume of water Calculations: Calculate the following. Use the information about the accuracies of the equipment to calculate the ranges. Express the results of each calculation to the appropriate number of significant figures and show the units in each case. 1. The density of water from part 1 (with range and percent error) 2. The density of water from part 2 (with range and percent error) 3. The density of water from part 3 (with range and percent error) 4. The density of liquid from part 4 (with range) 26 5. 6. 7. 8. The mass of the object from part 6 (with range) The volume of the object from part 6 (with range) The density of the object from part 6 (with range) From the possible metals and their densities, calculate a percent error for your metal Results: Make clear statements regarding the results of each density calculation. Show the possible range in each case and the % error. Glassware used Density Range % error Volumetric pipette Etc. Possible element Unknown liquid Unknown object Discussion: In addition to the standard discussion topics discuss the following: The density of water was measured three times. A different piece of glassware was used each time to measure the volume. Based on your percent errors and range, which measuring tool is likely to provide the best accuracy and precision? If one result is different from the other two, explain how the glassware used affected that result. Explain your answer using the accuracy of the equipment used. The density of water is 1.000 g/mL at 3.98oC. Explain why your results were not exactly the same as the literature value. Provide a possible identity of your unknown object based both qualitative and quantitative evidence Post-Lab Questions: 1. What is the density of a sample with a mass of 30.425g and a volume of 16.8mL? Report the answer to the correct number of significant figures. 2. The mass of an empty weighing bottle is 22.814 g when measured on a digital balance. The mass of the bottle containing a solid object is 71.115 g. The object’s volume is determined by the displacement of water in a 100mL (+/-0.1mL) measuring cylinder. The volume reading before the object is immersed is 49.5mL, and after immersion the reading is 86.0mL. a. What is the density of the object? b. How many significant figures should the answer have? c. What is the range of the answer? 3. A piece of household aluminum foil has a mass of 0.257g and measures 3.96 cm*11.2cm. If the density of aluminum is 2.70g/cm3, how thick is the aluminum foil? (Hint: the mass and density are given, so think carefully about how to calculate the volume of the foil.) References Finnegan, M. Place, H. Weissbart, B. (2000) Washington state university chemistry 101-102 laboratory manual. Star Publishing Company: Belmont, California 27 Experiment 3 Emission Spectroscopy Purpose: To compare predicted emission spectra with experimental emission spectra. Background Information: Electrons control the chemistry of the elements. In order to understand the role of the electron it is necessary to know something about the arrangement of the electrons in the atom. This information can be obtained by studying the way in which atoms absorb and emit light. It is first necessary to know something about the properties of light. Light is a form of energy known as electromagnetic radiation. One of the ways to describe light is to say that it has the properties of a wave. Light can be thought of as a wave traveling in a straight line (Figure 1). Figure 1 Direction of travel Light in a vacuum travels at 2.998x108 meters per second (or about 185,000 miles per second). The frequency is the number of waves that pass a given point during a given unit of time (usually one second). Light waves are also characterized by wavelength. Wavelength is the distance between crests of the waves (Figure 2). Direction of travel Figure 2 If the wavelength is short (λ1) then the number of waves which pass a given point per unit of time is large. This creates a high frequency wave (ν 1). If the wavelength is long (λ2) then the number of waves which pass a given point per unit of time is small. This creates a low frequency wave (ν 2). Since the speed of light is independent of wavelength or frequency, it is possible to summarize the behavior of light by the following equation: 28 λ* ν = c (Equation 1) Where λ = wavelength ν = frequency c = speed of light (3.00 x 108 m/s) Equation 1 indicates that when the wavelength is short (λ1), the frequency is high (ν 1). When the wavelength is long (λ2), the frequency is low (λ2). Equation 1 can be used to calculate the frequency of light if the wavelength is known. The frequency of light having a wavelength of 651 nm is calculated below. First we solve Equation 1 for frequency (see Equation 2). (Equation 2) ν= c λ Convert the wavelength from nanometers to meters. Note that 1.00 nm = 1.00 x 10-9 m. 651 nm 1.00 x 10-9m (Equation 3) = 6.5100 x 10-7 m 1.00 nm Then substitute in Equation 2 and calculate frequency: 3.00 x 108 m/s (Equation 4) ν= c = = 4.61 x1014 sec-1 -7 6.5100 x 10 m λ Note 1 = sec-1 sec This means that frequency is measured as wave crests per second. In the above example, a crest of the light wave passes a given point 461,000,000,000,000 times per second (461 trillion times per second). It is interesting to note that the only difference between radio waves, visible light and x-rays is the wavelength of the electromagnetic radiation. This is illustrated in Figure 3. Radio, TV 106 104 Visible light Ultraviolet Infrared X-rays Radar 102 Gamma rays Cosmic rays 100 10-2 10-4 10-6 10-8 10-10 10-12 10-14 Low energy High energy Wave length (m) Figure 3 If a narrow beam of white light is directed at a prism or diffraction grating, the white light is separated into its component parts. A continuous spectrum appears which shows that the white light is composed of light of all wavelengths. Low energy V 380nm B 450nm G 495nm Y 570nm O 590nm R 620nm 750nm Visible Light Spectrum High energy 29 Schematic diagram of emission spectroscopy apparatus Figure 4 If a narrow beam of light from a hydrogen discharge tube is directed at a prism or diffraction grating, a continuous spectrum is not obtained. Instead, a series of lines appear. These lines correspond to light of a certain wavelength. If a sodium lamp is used instead of a hydrogen discharge tube, a different series of lines appears. The lines of emission spectra are unique for each element. The study of emission spectra is called emission spectroscopy (Figure 4). A very simple experiment illustrates that emission spectra are caused by electrons. The spectra of He, Li, Be2+ and B3+ are quite similar in spite of the fact that their nuclei are different. These four species all have two electrons. The fact that the emission spectra will be different when the number of electrons changes (regardless of the nuclear composition) indicates that the electrons are responsible for the emission spectra. In 1913, Niels Bohr proposed a theory which yielded a fairly good explanation for the emission spectra. In time, this theory was shown to be partially incorrect, but some of Bohr’s original ideas have been retained in the modern atomic theory. The current theory of electronic structure assumes that an atom consists of a positively charged nucleus with one or more negatively charged electrons at a distance from the nucleus. This theory assumes that the electrons can have only certain energy levels (or specific distances from the nucleus). We say that the energy of an electron is “quantized” or that electrons can exist only in certain energy levels. The only way an electron can go to a higher energy level is to absorb energy. This concept is illustrated in Figure 5. When an electron absorbs energy, it is excited and is propelled into a higher energy level. Higher energy levels are generally farther away from the nucleus than lower energy levels. An exact amount of energy is needed to move an electron from the 2nd to 3rd energy level. A lesser amount of energy can not raise the electron to the 2½ energy level because there is no energy level between the 2nd and 3rd. Electron absorption of energy Figure 5 30 When an electron drops from a higher to a lower energy level, energy is released from the atom (Figure 6). Electron emission of energy Figure 6 In this case an excited electron returns to its ground state (the energy level normally occupied by the electron). Energy is released. Because the energy between the 2nd and 3rd energy level is always the same, an electron which drops from the 2nd to the 3rd level always releases a predictable quantity of energy in every atom. If an electron in energy level 3 drops to energy level 2, it will emit energy which is exactly equal to the difference in energy between levels 2 and 3 (see Figure 7). Energy emitted = ΔE = Ef - Ei = E2 - E3 (Equation 5) Ef = energy of the final level Ei = energy of the initial level Electronic energy levels and emission of energy Figure 7 Please note that in this case ΔE will have a negative value since the energy of the e- at E3 is greater than the energy of the e- at E2. This indicates that the electron has lost energy, confirming that the energy is emitted or given off. The electron could have fallen to energy level 1 and emitted a greater amount of energy (see Fig. 7) Where ΔE’ > ΔE ΔE’ = E1 - E3 (Equation 6) According to this theory it is not possible for the electron to fall to some intermediate level between E2 and E1. The energy of the electron is quantized. The electron can therefore occupy only certain levels as dictated by the principal energy level, n. An electron can also absorb energy and be excited to a higher energy level (see Figure 8). 31 It is interesting to note that the energy absorbed must exactly numerically equal the difference in energy of levels 3 and 4. Energy absorbed = ΔE’’ = Ef - Ei = E4 - E3 (Equation 7) ΔE’’ is positive since the energy is absorbed by the electron. Energy level n=4 is greater than energy level n=3. Electronic energy levels and absorption of energy Figure 8 If the exact amount of energy is not absorbed the electron will not make the transition. It is not possible for an electron to absorb enough energy to be excited to an intermediate level because there are no intermediate levels—the energy levels of the electrons are quantized. With these concepts it is fairly easy to explain why the emission spectrum of hydrogen is a series of lines. The electron of the hydrogen atom in the discharge tube is subjected to bombardment of energy when the tube is powered on. When this occurs some of the electrons absorb energy and these electrons are forced to leave their ground state to occupy a higher energy level (excited state). When the excited electron falls to a lower level, a certain amount of energy (ΔE) is emitted. Where ΔE = Ef - Ei Ef = energy of the e- at the final level Ei = energy of the e- at the initial level (Equation 8) The energy is emitted by each atom as a quantum of light. In 1900, Max Planck described the relationship between the energy and frequency of a quantum of light (and the energy and frequency of any electromagnetic energy). ΔE = h * ν (Equation 9) Where ΔE = Energy of emitted light (J) h = Planck’s constant (6.626x10-34 J*sec) ν = Frequency of light (sec-1) If Equation 8 and Equation 9 are combined, Equation 10 is obtained. Ef - Ei = ΔE = h * ν thus Ef - Ei = h * ν (Equation 10) This mathematical expression means that when an electron drops from a high energy level to a low energy level, a quantum of light, a flash with a fixed frequency and wavelength, is emitted. Since only certain energy levels are allowed, only certain electronic transitions can occur, and the emission spectrum of hydrogen consists only of certain wavelengths. Thus we should see an emission spectrum which consists of light flashes of a certain wavelengths. Since a tremendous number of atoms emit flashes simultaneously, we see instead, lines having wavelengths the same as those of the flashes. 32 Bohr’s great achievement consisted of constructing a theoretical model of the hydrogen atom and calculating its properties which fitted the experimental data exceedingly well. His model was found to be excellent for the hydrogen atom. Unfortunately, it failed for any other atom. Bohr’s model began with fundamental constants of nature such as the speed of light and the charge and mass of the electron, and the assumptions outlined above. He arrived at the Bohr-Rydberg equation for the energies associated with electronic transition. ΔE = Ef - Ei = -RH ( ΔE = -RH ( Where 1 Nf 2 1 Nf 2 ) - 1 =h*ν Ni 2 or in a more simplified version - 1 Ni 2 ) (Equation 11) nf = the principal energy level of the final level ni = the principal energy level of the initial level RH = Bohr-Rydberg constant, 2.178x10-18 J (calculated from fundamental constants) Pre-Lab Questions: 1. Find the frequency for the following wavelengths. 1a. 6.55 x10-7 m 1b. 4.84 x10-7 m 1c. 4.32 x 10-7 m 2. Find energy in joules for the following frequencies of light. 2a. 4.58 x 1014 sec-1 2b. 6.20 x 1014 sec-1 2c. 6.94 x 1014 sec-1 3. Find energy emitted or absorbed in Joules for the following electron energy level transfers. Was energy absorbed or emitted? 3a. ni = 1 to nf = 2 3b. ni = 3 to nf = 4 4. Find the frequency of light given off for the following electron energy level transfers. 4a. ni = 3 to nf = 1 4b. ni = 5 to nf = 3 Materials: Spectroscopy apparatus, discharge tube and power supply 33 Procedure: In this laboratory activity you will replicate Bohr’s work and evaluate the Bohr model of the hydrogen atom. Use the Bohr-Rydberg equation to calculate the energies associated with the transition for… ni = 3 to nf = 2 ni = 5 to nf = 2 ni = 4 to nf = 2 ni = 6 to nf = 2 You will make observations of hydrogen’s line spectra. Then from the observed wavelengths you will calculate the energy associated with each spectral line. You will then compare the line energy data from the experiment with theoretical values. To complete this laboratory activity, view into the spectroscope and make careful observations. Record the data and complete the calculations. Then compare the results of your experimental calculations with theoretical values. Data: Line Color of line spectra and other qualitative data Observed Wavelength of spectral line (nm) Red Green Blue Violet Calculations: Calculations Table 1: (this is an example of how you may wish to report your calculation, you will still need to show all of your work) Line Convert wavelength (nm) Observed frequency Observed energy of spectral line to wavelength (m)* of spectral line (sec-1)** (Joules)# Red Green Blue Violet * consult Equation 3 ** consult Equation 4 # consult Equation 9 The next step is to calculate the theoretical energies associated with the electronic transitions. To do this, use the Bohr-Rydberg equation (Equation 11). Calculations Table 2: (Violet has the highest energy followed by blue, green, yellow, orange, and red has the lowest energy of light) Electron energy level transition ni = 3 to nf = 2 ni = 4 to nf = 2 ni = 5 to nf = 2 ni = 6 to nf = 2 Theoretical energy in Joules 34 Likely color The last step is to match up the observed energies and the theoretical energies and to calculate the absolute and relative error. Determine the accuracy by calculating the absolute errors and relative errors for the red and violet lines only. The measured values will be the observed energies from calculation table 1. The theoretical values will be the theoretical energies from calculation table 2. Absolute error = E = | m – t | m = experimental measured value t = theoretical value Relative error (accuracy) E/t *100 E = | m – t | (see above) t = theoretical value Calculations Table 3: Keep in mind that red light has least energy and violet has the most. You might want to be careful how you match up your data for error analysis Line color Observed energy Theoretical energy Observed energy Relative error Absolute error Relative error Red Violet Results: Line color Red Violet Discussion: Discuss how this lab met the stated purpose. Discuss how the observed and theoretical energies matched. If the relative error is less than 5%, the data validates the Bohr model. Discuss one area for potential follow up experiments. Discuss the relationship between light and energy, frequency and wavelength. How does this lab support the observation regarding the quantized nature of electron energy levels? Post-Lab Questions: 1. Do some research and find the color spectrograph of a fluorescent light. What is the difference between the light the fluorescent bulb emitted and light emitted by electrons of excited atoms in a hydrogen discharge bulb? 2. Calculate the wavelength of radio waves transmitted by KMTT. The frequency of the these waves is 103.7 MHz (103.7 x106 sec-1). 3. Discuss why the orange sidewalk lights are called “Sodium Lamps.” Provide a citation for your source of information. References: Collins, V. Kahl, D. Perry, F. (1996) Good stuff from the chemistry laboratory. Warren Wilson College Press: Swannanoa, NC. Silberberg, Martin. (1996) Chemistry: The molecular nature of matter and change. Mosby: New York, NY 35 Experiment 4 Rapid oxidation of metallic fuels with color additives “Sparklers” Purpose: To determine effective procedures to overcome the difficulties in the creation of sparklers, especially in how to get a good mixture that keeps burning and a composition that binds easily and effectively to the wire or thin wood stick. Hypothesis: Predict what your sparkler should look like before it is ignited. Also propose a recipe for the unknown sparkler based on your observations of the other sparklers. Materials: Supplies Popsicle stick Rubber cement Wax paper Chemicals Oxidizer (provides oxygen for rapid oxidation) Potassium perchlorate Potassium nitrate Burner (other nitrates may also serves as an oxidizer) Fuel (rapid oxidation; reacts with oxygen and produces sparks) Magnesium powder Aluminum powder Iron Filings Magnalium (an alloy of magnesium and aluminum) Coloring agent (because of the electron configuration these metals burn with very specific colors) Copper II oxide Barium nitrate Strontium nitrate Sodium chloride Other (other materials that may be used in various sparkler recipes) Shellac Parlon Charcoal Sulfur Manganese IV oxide Background: In the fireworks industry, sparks are probably the oldest pyrotechnically produced effect and can be generated from many different materials. What makes a spark glow? It’s incandescence— yes, just like an incandescent bulb. When a solid is heated to a high temperature, it incandesces or glows. Sparks that you see in fireworks are solid or liquid particles that have been ejected from a burning surface. This surface has been heated to a high temperature (hotter than 1700 degrees C!) by the chemical reactions and flame. One of the factors involved in making sparklers is controlling the color. Manufacturers of pyrotechnics can produce sparklers with only a limited range of colors: they vary from dim red to orange to yellow to bright white. No other colors are possible from incandescing objects. You can see this range of colors if you watch the filament of a clear incandescent bulb as you vary the voltage delivered to it with a dimmer switch. Another factor important in sparkler production is the longevity of the spark. We don’t want the sparks to cool too fast or they’ll be duds: they’ll glow brightly at first, but then cool quickly, grow dim, and be invisible. We want the sparks to stay bright for a longer time, and to do this, they must continue to react and produce heat energy as they fall through the air. This spark reaction is 36 oxidation, with the spark as the fuel and the air oxygen as the oxidizer. In the sparkler recipes you’ll be testing, magnesium and aluminum are the fuels: 4Al (high temp) + 3O2 à 2Al2O3 + heat 2Mg (high temp) + O2 à 2MgO + heat Pre-Lab Questions: 1. Many metal compounds can be identified by the color that is produced when they are heated. Do some research and determine what colors will be produced when the following materials are heated (burned)? A. Copper salts B. Strontium salts C. Barium salts 2. Many metals can be identified by what there sparks look like. Do a little research and then draw and describe what the following sparks will look like. A. Iron Filings B. Magnesium C. Aluminum Safety: • • • • • Do not grind potassium perchlorate, KClO4 or the potassium nitrate, KNO3! They could detonate. Do not grind any of the ingredients once they have been mixed together! They could detonate. Do not make a recipe that makes more than 2.0 grams of powder. If you have powder left over after your sparkler has been made, please return it to the waste beaker in the fume hood. Do not place any of the powder in the trash cans. Wash your hands thoroughly before leaving the lab. Potassium Perchlorate recipes: KClO4 Mg powder Sr(NO3)2 30% 50% 20% KClO4 Magnalium Copper II Oxide 50% 39% 11% KClO4 Al Powder Copper II Oxide 50% 39% 11% KClO4 Al powder Copper II oxide 60% 30% 10% KClO4 Al powder Magnalium 60% 30% 10% KClO4 Magnalium Sulfur Copper oxide Barium nitrate Sodium chloride 20% 42% 12% 12% 12% 2% KClO4 Ba(NO3)2 Magnalium Titanium Mg Powder 20% 14% 16.5% 33% 16.5% 20% 45% 17% 13% 5% KClO4 KNO3 Sulfur Charcoal Al powder Other Perchlorate recipes: Red KClO4 Sr(NO3)2 Magnalium Sulfur 21% 46% 20% 13% Green KClO4 Ba(NO3)2 Magnalium Sulfur Copper II oxide Lilac 37 7% 64% 18% 5% 6% Potassium nitrate recipes: KNO3 CuO Magnalium Carbon 12.5% 25% 37.5% 25% KNO3 Mg Powder Charcoal Sulfur 70% 15% 10% 5% KNO3 Al powder Charcoal Sulfur 47% 18% 24% 11% KNO3 Charcoal Sulfur 75% 15% 10% KNO3 Mg Powder Charcoal Sulfur Iron Filings 33.8% 45% 6.7% 4.5% 10% Ba(NO3)2 Al powder Fe filings Boric acid 45% 19% 34% 2% Sr(NO3)2 Shellac dry 86% 14% Other Nitrate based recipes Sr(NO3)2 Mg powder Fe filings 40% 17% 43% Ba(NO3)2 Iron Filings Al flakes 50% 30% 20% Sr(NO3)2 KNO3 Magnalium Sulfur Charcoal 25% 30% 37.5% 2.5% 2.5% Procedure: 1. Decide which ingredients and what amounts you will use to create no more than 2 grams of sparkler powder. Obtain your instructor’s initials on the recipe, indicating that you have permission to proceed. 2. Make sure each of the ingredients is a fine powder. Lumps decrease the surface area for reaction, slowing the process (it’ll be a dud). Use a mortar and pestle if necessary, but DO NOT grind the potassium perchlorate or the potassium nitrate!! 3. Mix the powders VERY WELL together. Make sure you have a uniform texture. 4. Obtain a piece of wax paper and a Popsicle stick from the materials table. 5. Write your name on one end of the stick. 6. Place your powder onto your wax paper. 7. Gently hold up opposite sides of the wax paper to form a thick line of powder down the center, then place it down on the lab counter. 8. Holding onto the end marked with your initials, dip the stick into the rubber cement so that one half of the stick has a medium coat of the glue. Use a clean stick to return the excess glue to the rubber cement container. 9. Immediately press the gooey end of the stick into the pile of powder. Work the stick, pressing down and pulling up, coating both sides and the edges with powder. You should have a nick thick coating on the stick. If you have no powder left on the waxed paper then you are finished. 10. If you have given your sparkler as many coats as you can and still have leftover powder, return the powder to the proper waste beaker. 11. Carefully carry your sparklers over to the fume hood. Place them down on your wax paper and leave them to dry. Your sparkler must dry completely. 38 Data: Do not write on the following data table make your own. Also remember you will need to have a data table for at least 2 sparklers and a table for the unknown. Recipe 1 Name Mass % Mass (g) Observations before ignition Oxidizer Metal fuel Color additive Total Predicted Charactersitics Color Spark Observations of burning Color sparkler Spark Other Unknown ID # Oxidizer Metal fuel Color additive Observations Name Results: Connect the color of sparkler to the chemical used. Connect the type of spark to the metal used. Copper produced a _____________ colored sparkler. The aluminum produced a ___________ colored spark that had a _____________tail. For the unknown determine the • Metal fuel • Oxidizer • Color additive Discussion: None at this time. Post-Lab Questions: None at this time. References: Beth Eddy 39 Experiment 5 Inorganic Nomenclature Purpose: The names and formulae of common anions and cations will be learned. The formulae of compounds will be written from their names. The name of compounds will be written from their formulae. Background: In order to communicate with other scientists around the world, chemists use a set of standard rules for naming compounds. The system used is called the IUPAC (International Union of Pure and Applied Chemists) system. The older way of naming chemicals, which is still in use, utilized a mixture of various systems, and such names are now called common or trivial names. IUPAC names will be used as much as possible in this class, but some very common trivial names will also be used so you should become familiar with them as well. The set of rules should be learned along with a list of names and formulae for common ions. Predicting Charge Column 1A 2A 3A 4A 5A 6A 7A 8A 1B-8B Family Name Alkali metals Alkaline earth metal Boron family Carbon family Nitrogen family Oxygen family Halogens Noble gas Transition metals Charge of Simple Ions 1+ 2+ B and Al 3+, others vary Varies N can be 3-, others vary 21None Varies however, 2B typically 2+, 3B typically 3+ Binary Ionic Compounds Nearly all ions are made of a metal and a nonmetal. The easiest compounds to name are binary ionic compounds. These contain only metal ions (cations) and nonmetal ions (anions) whose charges can be predicted from the periodic table. If the cation is a metal with a known charge, it is only necessary to know the names of the elements in order to name the compound. The metal (cation) is named first and is given its elemental name. The nonmetal (anion) is named next. To name the nonmetal, the root of the element name is used but given the suffix –IDE. In these simple cases it is not necessary to use prefixes or Roman numerals as the charges on all ions are known. Examples: NaCl is called sodium chloride K2O is called potassium oxide MgF2 is called magnesium fluoride The formula for a binary ionic compound is called a formula unit. The formula unit describes the lowest whole number ratio of the atoms in the compound. Ionic compounds are crystals. In the crystal, regardless of size, there is a specific ratio of metal ions to nonmetal ions. There is not a discrete molecule. In the picture to the right there is 1 salt crystal. It is made up of sodium ions and chloride ions. The formula unit (lowest whole number ratio) is NaCl. 40 Compounds Made with Variable Charged Metals A transition metal or a post-transition metal can have more than one possible positive charge as a cation. It is not sufficient to simply name the metal and the nonmetal. The charge of the metal ion must be included in the name. It is included as a Roman numeral immediately after the name of the metal. It is not possible to predict this charge, but it can be deduced if the number and charge of the nonmetal anions is known. The anions again use the suffix –IDE. Examples: FeCl2 is called iron II chloride. The iron has a 2+ charge FeCl3 is called iron III chloride. The iron has a 3+ charge Binary Molecular Compounds When two nonmetals form a compound, they must each be named using prefixes. The more metallic (to the left or further down the periodic table) element is named first. It is simply the name of the element. It is preceded by a prefix which tells exactly how many atoms of that type are present. In the case where there is only one atom of the first element, the prefix mono- is dropped and presumed redundant. The second element is then named. It is again preceded by a prefix which shows the number of atoms of that element present. Then it is given the –IDE ending. In the case where the prefix ends in a vowel and the element name begins with a vowel, the vowel is dropped from the prefix. A list of common prefixes follows. Prefix Number of atoms A molecular formula describes the quantity and type of each mono1 atom in the molecule. Below are 4 water molecules. Each di2 discrete molecule has the formula of H2O. tri3 tetra4 penta5 hexa6 hepta7 octa8 Examples: N2Cl4 is called dinitrogen tetrachloride SF6 is called sulfur hexafluoride (mono- is not used before sulfur) P2O5 diphosphourous pentoxide (pent- instead of penta- before oxygen) Polyatomic Ions Many ionic compounds contain polyatomic ions. It is necessary to learn the names of all the polyatomic ions. It is also necessary to learn the charges of the polyatomic ions. When the cation is a metal with a known charge, the metal is named first, followed by the name of the polyatomic ion. When the cation is a metal with a variable charge, Roman numerals are used to show the charge on the metal ion. The charge of this cation can be deduced by knowing the charge on the polyatomic anion. Compounds which contain polyatomic ions are usually named without prefixes. Examples: • K2SO4 : potassium sulfate. • Cu2SO4 : Sulfate has a 2- charge. Therefore, the two copper ions must have a net charge of 2+, since there are two coppers each copper must have a 1+ charge. This is called copper I sulfate. • CuSO4 : Sulfate has a 2- charge. Thus, the copper ion must have a 2+ charge. This is called copper II sulfate. • Pb(SO4)2 : Sulfate has a 2- charge. Therefore, the sulfate ions must have a net charge of -4, so the lead ion must have a 4+ charge. This is called lead IV sulfate. 41 Writing Formulae from Names All of the compounds in this laboratory exercise will have a neutral charge. To write the formula for a compound from its name, you must use enough ions to balance the charges. Use subscripts to indicate the number of times an ion occurs in a compound. When polyatomic ions are used, the entire ion is placed in parentheses and subscripts are added as needed. Examples: calcium chloride Calcium has a 2+ charge and chloride has a 1- charge, therefore each calcium ion requires two chloride ions to achieve a net charge of zero on the compound. The formula will be CaCl2. The subscript 2 shows that there are two chlorides. This CANNOT be written as Ca2Cl. aluminum oxide Aluminum has a charge of 3+, oxygen in oxide ion is 2-. The compound between aluminum and oxygen, Al2O3, has two aluminum ions and three oxide ions. The net charge of 6+ on the two aluminum ions is balanced by 6on the three oxide ions. barium nitrate Ba(NO3)2: Barium has a 2+ charge, nitrate has a 1- charge. It takes two nitrates to combine with each barium. barium Ba3(PO4)2: The three 2+ barium ions combine with two 3- phosphate ions. phosphate Note that you must know the charges on the ions before you can write the formula for the compound. Simple Anions and Polyatomic Anions Most of the polyatomic ions in this course will be anions and will consist of one or more elements combined with oxygen. These polyatomic anions are negatively charged. You will need to memorize all of these polyatomic anions and their charges. Note: when the charge on an ion is 1+ or 1- it is customary to omit the 1 and simply write + or -. Anions with a 1- charge Hypochlorite ClO Nitrite NO2Chlorite ClO2Nitrate NO3Chlorate ClO3 Cyanide CNPerchlorate ClO4Cyanate OCNPermanganate MnO4 Thiocyanate SCNAcetate CH3COOHydroxide OHHydrogen carbonate HCO3 Hydrogen sulfate (bisulfate) HSO41(bicarbonate) All nonmetals from 7A F-,Cl-, Br-, etc. 2- Oxide Sulfide Sulfite Sulfate Thiosulfate O S2SO32SO42S2O32- Phosphate Nitride PO43N3- Anions with a 2- charge Carbonate Oxalate Chromate Dichromate Hydrogen phosphate Anions with a 3- charge Phosphide 42 CO32C2O42CrO42Cr2O72HPO42P3- Cations Some common metal cations you will be expected to know are given below. When a Roman numeral is used it indicates that a metal is present which can have more than one possible charge. Cations which can have more than one possible charge must include the charge along with the name of the metal. This is not an all-inclusive list. Cations with a 1+ charge Copper I Cu+ Ammonium NH4+ + Silver ion Ag All metals ions from 1A Li+, Na+, K+, etc. Proton H+ Chromium II Cadmium ion Copper II Iron II Lead II Manganese II Cr Cd2+ Cu2+ Fe2+ Pb2+ Mn2+ Cations with a 2+ charge Nickel II Tin II Zinc ion Mercury I Mercury II All metals from 2A Chromium III Iron III Cr3+ Fe3+ Cations with a 3+ charge Aluminum ion Cobalt III Lead IV 2+ 4+ Pb Cations with a 4+ charge Tin IV Ni2+ Sn2+ Zn2+ Hg22+ Hg2+ Be2+, Mg2+, etc Al3+ Co3+ Sn4+ Acids Many acids are compounds which donate H+ ions. When a H+ is attached to any of the polyatomic anions on the previous page an acid is formed. To name an acid, first look to see if any oxygen atoms are present. If no oxygen atoms are present the prefix is HYDRO- and the suffix –IC is added to the root of the element name. Finally the word acid is added. Examples: HCl is called Hydrochloric acid (the acid from the chloride ion) HCN is called Hydrocyanic acid (the acid from cyanide ion) If any oxygen atoms are present, no prefix is added. The root of the polyatomic anion name is used. a. If the anion name ends in –ATE, the suffix –IC is added to form the acid name. b. If the anion name ends in –ITE, the suffix –OUS is added to form the acid name. Examples: H2SO4 is called sulfuric acid (the acid from the sulfate ion) H2SO3 is called sulfurous acid (the acid from sulfite ion) Hydro is never used to name an acid when oxygen is present in the anion. 43 Acid Salts Acid salts are formed when a metal ion replaces some, but not all, of the hydrogen ions in a polyprotic acid (an acid with more than one proton i.e. H2SO4). The word hydrogen then appears in the name of the salt. Examples: KHSO4 is called potassium hydrogen sulfate. When there is more than one metal ion or hydrogen ion, prefixes are used to indicate the number of these ions Examples: Na2HPO4 is called disodium hydrogen phosphate NaH2PO4 is called sodium dihydrogen phopshate Bases The names of many common bases follow the normal rules for naming salts but the anion involved is the hydroxide ion, OH-. Examples: NaOH is sodium hydroxide Ca(OH)2 is called calcium hydroxide CuOH is called copper I hydroxide Pre-Lab Questions: 1. Look at the ingredients on a tube of toothpaste. Does the toothpaste contain fluorine or fluoride? What is the difference between fluorine and fluoride? State the brand of toothpaste that you used for this assignment. 2. Look at a soda pop container. Determine if it is a low sodium beverage (less than 35mg of sodium in 240mL). What is inaccurate about this statement? What kind of beverage did you use? 3. People refer to the chemical used in swimming pools as chlorine. Why is this inaccurate? What chemical is in fact used in swimming pools? 4. Define the following trivial names of ions a. Ferric b. Ferrous c. Stannic d. Stannous e. Plumbic f. Plumbous 5. From the above list state a rule that governs the assignment of the suffix –IC and the suffix – OUS regarding the trivial naming of variably charged metal ions. 44 Procedure: Read the data and observation section to know how to arrange data table. For each question write out the question and its answer in your laboratory notebook in the data section. Provide a brief explanation which shows your reasoning where appropriate. Data and Observations: Arrange data table as follows. 1. To which family does the metallic element in each of the following compounds belong? What is the charge on the metallic element in each of the following compounds? Question Answer Explanation Alkali metals lose e- to be more like the nearest noble gas. a. Na3N Alkali metals, Na+ b. MnO2 a. Na3N c. MgS e. Al2(SO4)3 b. MnO2 d. PbCl4 f. Ni(NO3)2 1. To which family does the non-metallic element in each of the following compounds belong? What is the charge of the non-metallic element in each of the following compounds? Example Question 1. Na2O 2. Ca3P2 Answer O is a member of the oxygen family Phosphorus is a member of the nitrogen family. a. Na2S b. CrCl3 It has a 2- charge. Explanation Calcium has a 2+ charge. There are 3 calcium ions, each with a 2+ charge for a net charge of 6+. Each of the two phosphorous ions must therefore be 3-. c. Li3N d. BaO e. CaF2 f. KBr 3. Name the following compounds Example Question Answer Explanation nickel III SO42- is called sulfate. The sulfates have a net charge of 6-, Ni2(SO4)3 sulfate. a. b. c. d. NaI CuSO4 KHCO3 KBr e. f. g. h. therefore the two Ni ions must each have a charge of 3+. Nickel is a transition metal and therefore the name must include the charge. Na2SO4 N2 O4 SrCO3 Ca3(PO4)2 i. NaOCN j. LiNO3 k. SF6 l. K2C2O4 m. SCl2 n. Pb(CH3COO)2 o. (NH4)2Cr2O7 p. Sn(ClO4)4 4. Write the formula for the following compounds. Question Answer Explanation 2+ HgI2 Mercury II is Hg and iodide is I-. It will take two iodide ions to Example balance the charge on each mercury ion. Mercury II iodide a. Iron II sulfate g. Tin IV oxalate m. Mercury II nitrate b. Barium phosphate h. Chromium III carbonate n. Calcium perchlorate c. Sulfur dioxide i. Dipotassium biphosphate o. Strontium hydrogen carbonate d. Sodium oxalate j. Magnesium nitrite p. potassium cyanide e. Potassium permanganate k. Chromium II bromide q. mercury I bromide f. Sodium hydrogen sulfate l. Silver chromate r. Nickel II nitrate 45 5. Name the following acids or bases. Question Answer Explanation Example chlorous acid this acid comes from a chlorite ion and thus the ending is _ous HClO2 a. HBr b. HCN c. HNO3 d. H2S e. H2SO4 f. Mg(OH)2 g. H3PO4 h. KOH i. HNO2 j. HClO3 4. Write the formula for the following compounds. Question Answer Explanation Example Ba(OH)2 Barium ion is always Ba2+, Hydroxide is OH-. Two Barium hydroxide hydroxides are needed for each barium. a. b. c. d. Hydrofluoric acid Perchloric acid Aluminum hydroxide Copper II hydroxide e. Oxalic acid f. Hypochlorous acid g. Ammonia h. Carbonic acid i. Potassium hydroxide j. Acetic acid k. Sulfurous acid l. Ammonium hydroxide Calculations: Not applicable Results: Not applicable Discussion: Not applicable Post-Lab Questions: Not applicable References Collins, V. Kahl, D. (1995) Good stuff from the chemistry laboratory. Warren Wilson College Press: Asheville, North Carolina Finnegan, M. Place, H. Weissbart, B. (2000) Washington state university chemistry 101-102 laboratory manual. Star Publishing Company: Belmont, California Finnegan, M. Place, H. Weissbart, B. (1997) Washington state university chemistry 105-106 laboratory manual. Star Publishing Company: Belmont, California 46 Experiment 6 Determination of an Empirical Formula Purpose: To determine the simplest (empirical) formula of a chemical compound through experimentation. Develop a hypothesis regarding the mass of the material at start and the mass of the material at the end. (Will the burnt magnesium have more or less mass than the unburned magnesium?) Background: Chemical formulae and moles The subscripts in a simple formula provide the relative number of atoms in a substance. For example, in the compound CO2 the subscripts show that each carbon dioxide molecule contains one carbon atom and two oxygen atoms. A simple formula provides two pieces of information. The subscripts in the simple formula give both the number of atoms and the number of moles in a substance. • For 1 molecule of CO2 there is one atom of carbon and two atoms of oxygen • For 10 molecules of CO2 there are ten atoms of carbon and twenty atoms of oxygen • For 1 dozen molecules of CO2 there are 1 dozen atoms of carbon and 2 dozen atoms of Oxygen • For every 6.02 x1023 molecules of CO2 there are 6.02 x1023 atoms of carbon and 2x6.02 x1023 atoms of oxygen • For 1 mole of CO2 there is 1 mole of carbon atoms and 2 moles of oxygen atoms • For 0.25 mole of CO2 there are 0.25 moles of carbon and 0.50 moles of oxygen atoms In the laboratory we can not count individual numbers of atoms as we did in the example above. Thus, we use the mole concept to determine the formulas because 1.00 mole of any element has 6.02 x1023 atoms and a mass which is equal to the atomic mass. This concept allows us to count atoms indirectly. If the mass of an element is known, the moles and atoms of that element can be calculated. In practical terms the problem is to determine how many grams of each element are present in the compound. In this experiment we will start with pure magnesium, and then determine the mass of magnesium. The magnesium will react with oxygen to produce magnesium oxide. Because magnesium and magnesium oxide have very high boiling temperatures the mass of magnesium will not change during the experiment. When we produce magnesium oxide, oxygen will combine with the magnesium; the increase in mass will be due to oxygen that has reacted with the magnesium. We can find the mass of oxygen in the magnesium oxide by subtracting the mass of magnesium from the total mass of the compound. 47 When the masses of magnesium and oxygen in the sample are known we can calculate the moles of magnesium and oxygen. Once the moles of magnesium and the moles of oxygen are known the simple whole number ratio of atoms can be calculated. This information will provide the formula. In this experiment we will carefully measure the mass of magnesium to three decimal places using the analytic balance. We will then combust (burn with oxygen) the magnesium in the presence of oxygen to form magnesium oxide. When we produce magnesium oxide, oxygen will combine with magnesium. The oxygen content will be determined by subtracting the mass of magnesium from the final mass of the magnesium oxide. Once we know the mass of magnesium and the mass of oxygen we can calculate the moles of magnesium and the moles of oxygen and the formula of the magnesium oxide. Error analysis In this lab you will form magnesium oxide from a given quantity of magnesium metal. The mass of the magnesium oxide is affected by experimental errors. It is useful to know how much magnesium oxide should have been formed theoretically and compare that to the quantity that was formed experimentally. If the values are close it may indicate that experimental error did not significantly impact your results. If on the other hand the difference is large between the mass produced experimentally and the mass that you should have obtained theoretically then experimental error did have an impact on your results. To calculate how much product should be produced it is necessary to write the balanced equation then use dimensional analysis to convert grams of reactants to grams of products that should be produced. For instance: If 1.035 grams of oxygen reacts with sulfur to form sulfur monoxide how much sulfur monoxide could be formed? 2S(g) + O2 (g) à 2SO (g) Step 1: Find moles of oxygen 1.035 g O2 1 mol O2 32.00 g O2 = 0.03234 mol molecular oxygen From the periodic table 1O=16.00g/mol thus O2=32.00g/mol Step 2: Convert moles of molecular oxygen to moles of sulfur monoxide 0.03234 mol O2 2 mol SO 1 mol O2 = 0.06469 mol SO From the balanced equation Step 3: Convert back to grams 0.06469 mol SO 48.07 g SO 1 mol SO From the periodic table 1 mole of oxygen is 16.00g/mol 1 mole of sulfur is 32.07g/mol = 3.110 g sulfur monoxide Or you could combine all the steps 1.035 g O2 1 mol O2 2 mol SO 32.00 g O2 1 mol O2 48.07 g SO 1 mol SO Once the theoretical mass of the product is = 3.110 g sulfur monoxide | experimental-theoretical | 48 *100 = % error calculated this can then be used in the percent error formula to determine the accuracy of your data. theoretical Pre-Lab Questions: 1. From the periodic table find the atomic weights for: a. Oxygen c. Carbon b. Magnesium d. Nitrogen 2. Find moles in: a. 0.101grams of magnesium c. 1.25 grams elemental hydrogen (H2) b. 0.102 grams of elemental oxygen (O2) d. 12.35 grams of water (H2O) 3. Find grams in: a. 12.1 mole nitrogen c. 1.087 mole carbon monoxide (CO) b. 0.00345 mole chlorine d. 0.00891 mole of magnesium oxide 4. 12.25 grams of hydrogen are combined with excess oxygen to form water. 2H2(g) + O2(g) à 2H2O(g) a. How many grams of water should be formed? b. Find the percent error if 99.75 grams of water are formed. 5. Find the mole ratio (divide both materials by the smallest coefficient) Materials: 0.100 - 0.200 g Mg ribbon Crucible and lid Clay triangle, ring and ring stand Steel wool Bunsen burner Wire gauze Distilled water Tongs Procedure: 1. Obtain a crucible and lid. Wipe out crucible with paper towels, do not get the crucible wet. Use only crucible tongs to handle the crucible and lid for the remainder of the experiment. Support the crucible on a clay triangle, heat gently at first and then heat with an intense blue flame for 5 minutes. (A yellow flame will deposit carbon on you crucible which negatively effecting your results). Using tongs (the crucible is HOT!), place them on a wire gauze and allow them to cool to room temperature. The crucible may appear dirty; in most cases they have permanent stains. If you are uncertain about the cleanliness of the crucible, check with your instructor. 2. While the crucible is cooling, obtain a sample of magnesium ribbon from your instructor and polish it with steel wool. 3. Use an analytical balance to determine the mass of the cooled crucible and lid; record the mass directly into your lab notebook data table. Make certain the balance reading is zero prior to obtaining your mass. 4. Curl the Mg ribbon to lie in the crucible and determine the combined mass of the crucible, lid and magnesium, then record this mass directly into your lab notebook. Verify that the mass of your magnesium is between 0.100 grams and 0.200 grams. 49 5. Place the crucible containing the Mg ribbon and lid on the clay triangle. Heat slowly, occasionally lifting the lid to allow oxygen to reach the Mg ribbon. If too much air comes in contact the Mg ribbon, rapid oxidation will occur and it burns brightly. This will negatively impact your results. You do not want this to happen. Immediately return the lid to the crucible. 6. a. Continue heating the crucible until no visible change is apparent in the magnesium ash at the bottom of the crucible when the lid is lifted b. Remove the lid. Continue to heat the open crucible and ash for 30 seconds. Remove from the heat, replace lid and allow the crucible to Experimental Setup cool to room temperature. c. Add 3 drops of water to decompose the magnesium nitride formed during combustion. Caution: cool water added to a hot crucible will cause it to break. Make sure it has cooled before adding water. d. Dry the ash with a low blue flame for a minute or two and then allow it to cool. Allow the steam to escape by setting the lid slightly to one side leaving a gap. e. Measure the mass of the crucible, lid and ash on the same balance that was used earlier and record. 7. a. Reheat the sample for 1 minute b. Again measure the mass of the crucible, lid and ash. If this second reading is greater than + or – 1% from that recorded for the 1st mass repeat procedure 7. Repeat as required until the change in mass is less than + or – 1%. 8. Clean up your crucible with water and invert on the drying rack. Data and Observations: 1. 2. Mass of crucible and lid Mass crucible, lid and magnesium 1st mass of crucible lid and magnesium oxide 2nd mass of crucible lid and magnesium oxide 3. Final mass crucible lid and magnesium oxide Observations 50 Calculations: 4. 5. 6. 7. 8. Mass of magnesium (data item 2. –data item 1.) Mass of oxygen (3. – 2.) Moles of magnesium Moles of oxygen Experimental ratio between moles of magnesium and moles of oxygen (ratio of 6.:7.). Use significant figures. 9. Divide both the moles of O and the moles of Mg by the smaller coefficient. 10. If the ratio in step 9 is within + or - 0.15 units of an integer, you can round off to the closest integer. If this is not the case, multiply both terms in the ratio by the same number until a whole number for each is achieved. Error analysis 1. Determine how much MgO was produced (3.-1.). (This is the Experimental mass.) 2. 3. Use the balanced equation and stoichiometry to find out how much magnesium oxide should have been produced. (This is the Theoretical mass.) Use the % error to evaluate this experiment. (|Theoretical-Experimental|/Theoretical)*100 = % error Results: 1. Empirical formula of oxide formed 2. Name of oxide formed 3. % error Discussion: Summarize results, comment on the obvious sources of error, explaining how they would make your results too high or low, use the percent error as evidence. A percent error of less than 10% is acceptable, less than 5% is good, less than 1% is excellent. Address the discussion questions provided in “How to Write a Lab Report.” Post-Lab Questions: 1. In a laboratory experiment to determine the formula of the compound formed between Cu and S, a student obtained the following data. Mass of crucible 19.732 g Mass of crucible and Cu 27.304 g Mass of crucible and compound of Cu and S 29.214 g What is the formula of the compound? 2. A compound containing nitrogen and oxygen is analyzed. A sample is found to contain 0.483 g N and 1.104 g O. What is the simplest formula of the compound? Show all calculations, use dimensional analysis, use sig figs. 3. A compound containing carbon and hydrogen is analyzed. A sample is found to contain 4.804 g C and 1.210 g H. What is the empirical formula of this compound? Show all calculations, use dimensional analysis, use sig figs. References: Collins, V. Kahl, D. Perry, F. (1996) Good stuff from the chemistry laboratory. Warren Wilson College Press: Asheville, North Carolina 51 Experiment 7 Reaction Rate Marble Lab Purpose: Graph the reaction rate of Calcium carbonate and Hydrochloric acid. Then determine the changes in reaction rate due to changes in temperature. Hypothesis: State a hypothesis regarding reaction rate and temperature (how will temperature effect reaction rate? Is the correlation negative or positive?). Materials: Marble (calcium carbonate) chunks Stop watch Ice Thermometer Balance 200 mL 1.0M HCl Styrofoam cup Burette stand & clamp Background: Reaction rate Reaction rate is a measure of the rate at which a chemical reaction proceeds from reactants to products. The units used are moles/sec. A useful model for understanding reaction rate is the Collision Theory. The collision theory proposes that atoms, ions and molecules can form a chemical bond when they collide, provided the particles have enough kinetic energy. Particles lacking the necessary kinetic energy may collide but simply bounce apart. The minimum energy that colliding particles must have in order to react is the activation energy. During a reaction, when the products initially collide, particles that are neither reactants nor products momentarily form. An activation complex is the arrangement of atoms at the peak of activation energy. Activation complexes are very unstable and last on the order of 10-13 seconds. Activation complexes are sometimes referred to as the transition state. Collision theory explains why some spontaneous reactions are imperceptibly slow at room temperature. The reaction of carbon and oxygen is spontaneous, but it has a high activation energy. At room temperature the collision of oxygen and carbon molecules are not energetic enough to overcome the activation energy required to break the O-O and the C-C bonds. These bonds must first be broken to form the activation complex. Thus the reaction rate of carbon with oxygen at room temperature is essentially zero. Factors effecting Reaction Rate Each of these factors which effect reaction rate can be considered using the collision theory of chemical reaction. When considering the collision theory it may be useful to think about these chemicals as clay balls rolling around on a billiards table. For a reaction to take place these clay balls have to collide with one another with enough energy and the correct orientation to overcome the activation energy to form a transition state and ultimately chemically recombine. Temperature Temperate is a measure of the average kinetic energy of a sample. Thus raising the temperature in general makes the molecules move faster, increasing their kinetic energy. This has two impacts on reaction rate. First the particles are moving faster and thus have a higher kinetic energy and thus are 52 more likely to have the energy required to overcome the activation energy. Secondly because the particles are moving faster this also increases the number of collisions. Thus by increasing temperature, the particles heat each other with more energy and more often. These two things increase the reaction rate. As temperature goes up reaction rate often goes up. Concentration As more particles are placed in a fixed volume container the concentration increases. If there are more particles in a given volume the chances of a collision also increase. Imagine trying to randomly collide two particles on a billiards table. The chance of a collision is less than if there are 100 particles on the billiards table. Surface area One of the universal answers to all questions in science is surface area. If you need a vague answer that is probably correct at some level, surface area is it. In fact surface area effects reaction rate. As particle size decreases surface area increases dramatically. For instance it is much easier to ignite some dry leaves than it is to ignite a whole tree. Why? Surface area. One of the ways chemists often increase surface area is to dissolve materials into aqueous solutions. This increases surface area. Two very concrete examples of surface area are coal dust and grain silos. Grain itself is not particularly flammable nor would you consider a chunk of coal as an explosive. Certainly there is a lot of potential chemical energy stored in these two items. However to access their explosive characteristics it is only necessary to increase their surface area. When pouring grain into the top of a silo a very fine powder of cellulose is suspended in the air. If this is ignited the silo explodes creating a bomb that will certainly snuff out the cigarettes of any nearby smoking farmers. Coal dust in the mines is notorious for huge levels of destruction. In fact a safety lamp was created that utilizes a wire mesh to reduce the possibility of igniting the coal dust. This of course leaves all the coal dust in the air for our ancestral miners to breathe. Coal dust explosions kill. Catalysts A catalyst is a substance that increases the rate of a reaction without being used up in the reaction. Catalysts reduce the activation energy necessary for a chemical reaction. Think of it as making the products stickier so that when things collide they are more likely to react. It is probably more correct to say that a catalyst provides a location for the reaction to take place on. The catalyst can be thought of as holding the product in an orientation so that they may react more easily. However catalysts themselves do not take part in the reaction and do not get used in the reaction. Catalysts do not provide the energy to allow a reaction to take place. Heat, concentration and particle size are never considered catalysts. Your body is at around 37 oC. Your temperature can not increase very much without causing brain damage and ultimately death. Yet without catalysts (in the body they are called enzymes) very few, if any, of the chemical reactions could take place in our cells at body temperature. What kind of heat is necessary to break down proteins? Consider the heat necessary to cook an egg. Cooking eggs is just using heat to break protein bonds (that is why they turn white). Enzymes allow you to break down the proteins in an egg at body temperature without having to reach very high temperatures. When a catalyst is used in a chemical reaction it is not consumed in the reaction. It is neither a reactant nor a product. In a car’s catalytic converter Pt is used to convert 2 poisonous gases into 2 harmless gasses. Please note how the Pt Pt is above the reaction arrow. The Pt is not consumed. The Pt 2CO + 2NO → 2CO2 + N2 never runs out. 53 The Mole An important concept used in this lab is the mole. This is not the insectivore that lives under your lawn that is spelled mole. In chemistry, when we talk about a mole it refers to a very specific number. A mole of something is 6.02x1023. A mole refers to a number. Just like a dozen means 12, a mole means 6.02x1023. Chemists use the concept of the mole to refer to the number of particles involved in a reaction. Example A: 1 mole of oxygen reacts with 2 moles of hydrogen. vs. Example B: 620,000 times a million, times a million, times a million molecules of oxygen reacts with 1,240,000 times a million, times a million, times a million molecules of hydrogen. Pre-Lab Questions: 1. Find grams a. For one mole of calcium b. For one mole of carbon c. For three moles of oxygen 2. How many moles of each element are in one mole of calcium carbonate (CaCO3)? 3. What is the mass of one mole of calcium carbonate in grams? Procedure: In this lab it is necessary to establish certain operational definitions. An operational definition is a statement regarding how an activity is to be conducted. They are very specific and are designed so another scientist can follow your procedure exactly. Write operational definitions for when the time starts and when the time stops. Also write any other definitions necessary such that if someone repeats your experiment they would get the same results. You will work in your lab groups of 4 students to share a balance at your lab station. 1. 2. 3. 4. 5. Cover your balance with a layer of plastic wrap (saran wrap©) Obtain 2 marble chips of similar size and shape. Mass the marble chips and record data. Mass a Styrofoam cup. Record this mass Place 100 mL 1.0M HCl in a Styrofoam cup. Set a thermometer up so that it is suspended in the acid solution without touching the cup. You will record the temperature every minute. Find the temperature of the acid and record this data. It should be close to 20 oC. Record the data using 3 sig figs. 6. Place cup and acid onto a balance and record mass. 7. Place one marble chip onto balance next to the Styrofoam cup (not inside the cup). 8. Prepare stopwatch to be used. When you are ready proceed. 9. Place the marble chip into the acid. Record mass every 15 seconds. 10. Record temperature every 1 minute. 11. Continue to record data until 2 concurrent data points are within 0.5 grams. At this point continue collecting data for 4 more time intervals. 9. Record the final temperature of the experiment. 10. Repeat the above steps, however use the chilled hydrochloric acid. 54 Data: (example data tables follow) Marble Chip 1 Mass of marble chip Time Mass Temp Observations Mass Temp Observations T0 T15 T30 T45 T60 Tetc Marble Chip 2 Mass of marble chip Time T0 T15 T30 T45 T60 Tetc Results: Each experiment should have its own graph. Graph: Provide a title for each graph. List the dependent variable on the Y axis (Mass) and the independent variable on the X axis (Time). Please do this graph such that it uses a full sheet of paper in a landscape orientation. Please use a ruler and equal graduations. Do not draw the graph with a (0,0) origin. The origin of the graph needs to be appropriate to the data graphed. Discussion: Compare the graphs and determine what effect each tested variable had on the reaction rate. Post-Lab Questions: 1. How does each of the following factors affect reaction rate? a. temperature b. particle size c. concentration d. catalysts 2. A tablet with a mass of 3.251 grams dissolves in 25 seconds. What is the average reaction rate? 3. Based on your graph, what would the reaction rate be if a tablet dissolves in a 40 degree water bath? 55 Experiment 8 Qualitative analysis: Flame test Purpose: To conduct a series of experiments and to make qualitative observations To identify an unknown salt using flame testing Hypothesis: (Write the hypothesis for part 2 of this experiment.) Provide a hypothesis regarding the identity of your unknown samples. Background: Qualitative observations In many sciences the primary focus of data collection is quantitative data. Quantitative data uses an instrument to act as an interface between the observer and the environment which is being observed. This tool provides a numeric value to quantify observations. The value in this kind of observations is that in theory anyone, who uses the instrument correctly, will get essentially the same data. Because science is driven by the concept of repeatability it places quantitative data collection in the forefront. Another form of data collection is qualitative observations. Qualitative observations focus on qualities: hot, cold, white, chunky, etc. These observations have limitations because an experimental situation can be described differently by different people. However qualitative observations are no less important. Because qualitative observations do not rely on instruments there is no need to consider instrumental error. People have a lifetime of experience describing things. In qualitative analysis the scientist observes the qualities of the situation or experiment and then, makes detailed notes of these observations. Then these notes are compared to other observations. In general qualitative observations can produce very satisfactory and accurate results. One particular weakness of qualitative observations is the difficulty in statistically analyzing the results. Flame Test A spectral line is the result of the emission of a photon of specific energy (thus specific frequency) when the electron moves from a higher energy state to a lower one. An atomic spectrum appears as a line rather than as a continuum because the atom’s energy has only certain levels or states. In the Bohr model the quantum numbers (1,2,3,n) are associated with the radius of the electron’s orbit, which is directly related to the atom’s energy. When the electron is in the orbit closest to the nucleus (n=1) the atom is in its lowest energy level, called the ground state. By absorbing energy equal to the energy necessary to push the electron from the first to second energy level the electron can move between levels. When the electron is in this higher than normal state it is called the excited state. In a flame test you excite the electrons from the ground state to an excited state. When the electron returns to the ground state it releases energy and this releases a photon of light. This analysis produces spectral data. The use of spectral data to identify and quantify substances is essential to modern chemistry. The terms spectroscopy, spectrometry and spectrophotometer denote a large 56 group of instrumental techniques that measure substances atomic and molecular energy levels from the spectra produced. The two types of spectra most often obtained are emission and absorption spectra. An emission spectrum, such as the H atom line is produced when atoms that have been excited to a higher energy level emit photons characteristic of the element as they return to a lower energy level. Some elements produce a very intense spectral line (or several closely spaced ones) that serve as a marker of the element’s presence. Such an intense line is the basis of flame tests, rapid qualitative procedures performed by placing a granule of an ionic compound or a drop of its solution in a flame. Some of the colors in fireworks and flares are due to the emission from the same elements as the flame tests: red from strontium, blue green from copper salts. The characteristic color of sodium vapor lamps and mercury vapor are due to one or a few prominent lines in their emission spectra. An absorption spectrum is produced when atoms absorb photons of certain wavelengths and become excited from lower to higher energy levels. Therefore the absorption spectrum of an element appears as dark lines against a bright background. Materials: Spot reaction plate (96 well tray) Cobalt glass All solutions are 0.2 Molar Copper nitrate Strontium chloride Lithium nitrate Bunsen burner Chloride test solution Sodium chloride Copper chloride Potassium chloride Lithium chloride Ni-Chrome wire loop 3.0 M HCl copper cleaner Sodium nitrate Strontium nitrate Potassium nitrate Procedure: You will use a piece of Ni-chrome wire (nickel chromium alloy) with a loop at the end to place one drop of a solution you want to test in a Bunsen burner flame. A piece of copper wire will be flame tested by placing it directly into the flame. You will also test the effects of a test solution to test for chlorides on each of the known salt solutions. Upon completion of these tests an unknown solution will be obtained and its identity will be determined. Part 1: Flame test of known chemicals Note: Do not exchange wires. For each solution only use the wire that is already in that solution. After you use the wire, be sure to put it back with the solution from which it came. To reduce the possibility of mixing up the wires, test only one solution at a time. When you have completed one test, return the solution to the test tube rack and obtain the next solution. 1. Remove the wire from the solution. 2. Place the loop of the wire with the solution on it in the flame. For testing the copper wire, first clean the wire with hydrochloric acid. Then place the wire directly into the flame using the tongs. 3. Observe and record the color of the flame. 4. Place the wire back in the solution. 5. Record results in the data table in your lab notebook. 6. Obtain next sample and repeat steps 1-6. 57 Part 2: Flame test unknown salt solution 1. Record the unknown’s Identification Number in data table. 2. Repeat above steps with unknown salt solution. 3. After determining the possible identity of the unknown solution, repeat the test with the known and unknown side by side to confirm the identity. Part 2.1: Hypothesis 1. Based on experimental evidence develop a well reasoned hypothesis based on experimental evidence. 2. Write down the hypothesis. Part 3: Test for chloride. Use a magic test solution to determine the presence of chloride in the unknown salt. 1. Place one drop of unknown solution into one of the wells of the 96 well tray. 2. Place one drop of the test solution into the same well. 3. Observe any changes 4. Record observations in data table. If chloride is present solid chunks will form when the two solutions are mixed. If no chloride is present no visible reaction will be observed. Data table (Write your observations directly into your lab notebook. Never write on scratch paper or onto a sample data table in your lab manual.) Name of salt solution Sodium chloride Potassium nitrate Copper nitrate Lithium chloride Strontium nitrate Copper chloride Sodium nitrate Lithium nitrate Potassium chloride Strontium chloride Copper wire Initial unknown flame test Unknown ID number Observations Observations Side by Side confirmation test Name or number Unknown ID number Known salt w/chloride Known salt w/nitrate Observations Chloride test Results: Unknown: Unknown chemical’s number ________ Unknown chemical’s identity____________________________ Discussion: As per “How to write a lab report.” Discuss the evidence that the metal and not the accompanying anions is responsible for the color of the flame. 58 Post-Lab Questions: 1. Group the substances based on the color of the flame produced. 2. What patterns do you notice in the groupings? 3. Predict the color of the flame for a substance called strontium sulfate. Explain your reasoning. 4. What evidence do you have that atoms of certain elements produce a flame with a specific color? 5. The yellow color of the flame for sodium indicates that the sodium atoms changed in some way when they were heated. Consider the following possibility that the electron configuration of sodium changed from [Ne]3s1 to [Ne]4s1. What is the difference between [Ne]3s1 and [Ne]4s1? 6. Do you think gold can be made by changing the arrangement of electrons in atoms? Explain. References: Finnegan, M. Place, H. Weissbart, B. (2000) Washington state university chemistry 101-102 laboratory manual. Star Publishing Company: Belmont, California Willbraham, A. Staley, D. Matta,. (1995) Chemistry 4th edition. Addison-Wesley: Menlo Park, CA Silberberg, Martin. (1996) Chemistry the molecular nature of matter and change. Mosby: New York, NY Stacy, A, Coonrod, J, Claesgens, J (2003) Alchemy: atoms elements, and compounds. Key Curriculum Press: Emeryville, CA 59 Experiment 9 Qualitative Analysis: Precipitation testing Purpose: To conduct a series of experiments and to make qualitative observations To identify an unknown salt using qualitative analysis Background: Qualitative observations As in the Flame Test laboratory activity this activity relies on qualitative observations. When conducting the flame tests it was sufficient to only note the color of the flame. However in this lab much more detailed observations will be necessary. Precipitation When a salt (a material containing a positively charged cation and a negatively charged anion) is placed in water the general tendency is for the polar nature of water to be attached to the cations and anions. These interactions are so strong that they pull the cation and anions apart from each other and into the solution making an aqueous solution. That is, a solution of water with free floating cations and anions. Here the water has not yet dissolved the salt into an aqueous solution, note that the water has both a positive (δ+) and a negative side (δ-). The salt (+ cations, -anions) are surrounded by polar water molecules. Some salts have such strong attractive forces between the positive cations and negative anions that the partial electrical charges found on the ends of the water molecule are not sufficiently strong to separate the cations and the anions from the crystalline structure. These salts are not soluble in water. That means when you place a salt that has very strong internal forces in water, the water molecule does not dissolve it and the salt remain a solid surrounded by water. Soluble means the water dissolves the salt and the solution will look clear (though it may have coloration). A material that dissolves in water is denoted with the symbol (aq), meaning aqueous. Insoluble means that water can not dissolve the salt and the solution will look cloudy and often be a flat white, yellow, pale blue, etc. These solutions may translucent (Tyndal effect). A material that does not dissolve in water is denoted with the symbol (s), meaning solid. 60 Solubility Rules 1. Salts containing Na+, K+, and NH4+ and acids are always soluble (dissolves, aqueous). 2. Salts containing NO3-, ClO3-, ClO4-, and CH3COO- are always soluble. 3. All Cl-, Br-, and I- are soluble except for those of Ag+, Pb2+, and Hg22+, which are insoluble (will not dissolve). 4. All SO42- are soluble except of those of Sr2+, Ba2+, Hg22+, Hg2+, and Pb2+. The sulfate salts of Ca2+, Ag+ are moderately soluble. 5. All OH- are insoluble expect those of alkali metals, which are soluble, and hydroxides of Ca2+, Ba2+, and Sr2+, which are moderately soluble. 6. All SO32-, CO32-, CrO42- and PO43- are insoluble except those of NH4+and alkali metals. 7. All sulfides S3- are insoluble except those of NH4+, the alkali metals and the alkaline earth metals, which are soluble. Acids and bases In the most general of terms, acids are chemicals that contain a proton that can be disassociated (falls off) when placed in water. This causes the water’s proton concentration to go up. H2O + HCl + à H2O + H+ +ClThis is most often written this way: H2O + HCl + à H3O+ +ClThis is true for all hydroxy acids (acids with free protons). These include H2SO4, HNO3 and others. Bases are chemicals that contain the polyatomic ion hydroxide (OH-). The hydroxide in a base disassociates in water. NaOH + H2O à Na+ + OH- + H2O The hydroxide, once disassociated, will wander around and scavenge any protons it can find to make water. H+ + OH- à H2O Thus any solution that OH- is placed into has a very low concentration of H+ because the OH consumes it all. Acids and base neutralize each other to form water and a salt. HCl + NaOH à H2O + NaCl In summary Acidic solutions have high proton concentrations and a low pH. Basic solutions have low proton concentrations and a high pH. 61 Pre-Lab Questions: When working on of the following questions completely write out the question and indicate the phase of each product and reactant in the answer. Use (aq) aqueous, (s) solid, (l) liquid, (g) gas. 1. Write and balance each of the following reactions. a. Ammonium hydroxide reacts with lead II nitrate to produce ammonium nitrate and lead II hydroxide. b. Iron III chloride reacts with silver nitrate to produce iron III nitrate and silver chloride. c. Sodium carbonate reacts with hydrochloric acid to produce carbon dioxide, water and sodium chloride. 2. Determine the products and then write and balance the equation. a. Sodium carbonate reacts with calcium chloride to produce… b. Sodium hydroxide reacts with copper II sulfate to produce… 3. Produce a written description for each of the unbalanced reactions. a. CuSO4 + Na3PO4 à Cu3(PO4)2 + Na2SO4 b. CaCl2 + Pb(NO3)2 à Ca(NO3)2 + PbCl2 4. Determine the products and provide a written description for each set of reactants. a. KI + AgNO3 à b. Potassium iodide reacts with lead II nitrate to produce… Materials: Spot reaction plate (well tray) Cobalt glass 0.2M Ammonium hydroxide 0.2M Hydrochloric acid 0.2M Nitric acid 1.0M Sodium Carbonate 0.2M Sulfuric acid Bunsen burner Litmus paper, blue and pink 0.2M Calcium chloride 0.2M Iron III chloride 0.2M Potassium Iodide 0.2M Sodium hydroxide 0.2M Sodium nitrate Ni-Chrome wire loop HCl cleaning solution 0.2M Copper II sulfate 0.2M Lead II nitrate 0.2M Silver nitrate Procedure: Part 1 Standardization The first part of this lab is to establish known qualitative data that will then be used to compare to the unknown sample. Here you are finding out what things look like when you know what everything is. This will help you when you are trying to figure out what unknown chemical you have. Before conducting any tests, make observations of what the chemicals look like in the pipette. These are your pre-observations. Acid-base test To conduct this test simply place a very small drop of the chemical on both the blue litmus paper and a very small drop on the red litmus paper. Red litmus- will turn blue in the presence of a base Blue litmus- will turn red in the presence of an acid Also make detailed observations regarding how the solutions absorb into the paper. 62 Precipitation test Clean your well tray very well with tap water. Then do a final rinse using distilled water. To conduct the precipitation test, place the well slide on a dark surface. Determine some method of establishing what is in each well used. I might suggest making a diagram on a separate piece of paper that incorporates a drawing of the well slide and the well labels. Following the scheme proposed in the data table, mix each of the indicated chemicals together. Use 1-2 drops of each reagent per reaction. As an example, based on the pre-made data table, you might conduct your trials in the following order: 1. NH4OH and lead II nitrate 7. NH4OH and Sodium hydroxide 2. NH4OH and hydrochloric acid 8. NH4OH and silver nitrate 3. NH4OH and sodium nitrate 9. NH4OH and calcium chloride 4. NH4OH and potassium iodide 10. NH4OH and nitric acid 5. NH4OH and copper II sulfate 11. NH4OH and sodium carbonate 6. NH4OH and sulfuric acid 12. NH4OH and iron III chloride Then start on the iron III chloride, etc. Make very detailed observations of your reactions. It will be crucial in future steps. You will use this data to help determine the identity of an unknown salt. You might consider the following abbreviations: NVR = no visible reaction (s) yellow = a yellow precipitate formed gas = bubbles formed Part 2 Unknown testing You will identify 2 unknown salt solutions. Collect your first unknown from your instructor. Record the unknown’s number in your data section immediately. Your unknown is one of the solutions for which you have already collected data. If you have made careful observations this activity will not be frustrating and will be rewarding. If you contaminated your samples or did not make careful observations during the first part of this lab, identifying your unknown salt will be difficult, if not impossible. Acid base test Conduct the litmus paper test using the same technique as before with your unknown. Make detailed observations. Not only changes in color but how the chemical absorbs into the paper. Compare this data with the data collected in part 1. Precipitation Clean your well tray very well with tap water. Then do a final rinse using distilled water. Place 23 drops of each of the different solutions into the well slide in some predetermined order. Then add 2-3 drops of your unknown to each of the known samples. Record your observations and compare them to the data collected in part 1. Part 3 Confirmation Tests Using the data collected in part 1 and 2 you should be able to establish with some level of certainty what your unknown is. It is now time to confirm your results. To do this you will perform each of the experiments side by side: the unknown with the most likely known salt solution. 63 As an example: For the precipitation test fill two rows of wells, one with the unknown and the other with the most likely known. Then one at a time add each of the different knowns to these samples first to the unknown then to the mostly likely known. Compare data. Do this for all the known chemicals. When you are done you should have two identical rows of mixtures. Then do the litmus test. Place a drop of the mostly likely known next to a drop of the unknown on blue then pink litmus. Compare. Repeat Repeat above steps for the 2nd unknown Pre-observation table Data: A large portion of your grade in this lab will be based on the correct identification of your 2 unknowns. The only possible way to identify your unknowns will be to have made excellent observations and then record those observations in a way that you can use them efficiently. There is a limited amount of time and you will have to use it carefully. Ammonium hydroxide Iron III chloride Sodium carbonate Nitric acid Calcium chloride Silver nitrate Sodium hydroxide Sulfuric acid Copper II sulfate Potassium iodide Sodium nitrate Hydrochloric acid Lead II nitrate Litmus paper test Blue Paper Pink Paper Ammonium hydroxide Iron III chloride Sodium carbonate Nitric acid Calcium chloride Silver nitrate Sodium hydroxide Sulfuric acid Copper II sulfate Potassium iodide Sodium nitrate Hydrochloric acid Lead II nitrate 64 Other observations Sample Precipitation Data Table Sodium Nitrate Sulfuric acid Sodium phosphate Sodium hydroxide Sodium carbonate Silver nitrate Potassium iodide Nitric acid Lead II nitrate Iron III chloride Hydrochloric acid Copper II sulfate Calcium Chloride Ammonium hydroxide Calcium chloride Copper II sulfate Hydrochloric acid Iron III chloride Lead II nitrate Nitric acid Potassium iodide Procedural Summary Silver nitrate Step 1 Complete known observations and tests. Sodium carbonate Step 2 Repeat tests with unknown. Sodium hydroxide Step 3 Develop a hypothesis regarding unknown. Sodium phosphate Step 4 Complete side by side confirmation tests of the unknown and the chemical identified in the hypothesis. Sulfuric acid 65 Results: Unknown number 1: Unknown chemical’s number ________ Unknown chemical’s identity____________________________ Unknown number 2: Unknown chemical’s number ________ Unknown chemical’s identity____________________________ Discussion: As per “How to write a lab report.” Post-Lab Questions: 1. Write and balance and indicate phase of the reaction of silver nitrate with each of the other chemicals used in this lab. Make sure to indicate phase of all compounds. References: Finnegan, M. Place, H. Weissbart, B. (2000) Washington state university chemistry 101-102 laboratory manual. Star Publishing Company: Belmont, California Willbraham, A. Staley, D. Matta, M. (1995) Chemistry 4th edition. Addison-Wesley: Menlo Park, CA Silberberg, Martin. (1996) Chemistry the molecular nature of matter and change. Mosby: New York, NY 66 Experiment 10 Chemical Reactions Purpose: To perform simple chemical reactions To identify the products of these chemical reactions To account for the reactions with balanced equations To name the reactants and products Background Information: Evidence of a Chemical Reaction A chemical reaction has taken place when a substance has changed its chemical composition. The evidence that a chemical reaction has taken place can be one or more of the following 1. The appearance of gas (bubbles) 5. Production of light 2. The appearance of a solid (precipitate) 6. Flame 3. A change in color 7. A change in smell 4. A change in temperature Reaction types It is useful to be able to determine which kind of reaction is taking place for it will make the prediction of the product that will form more easily. The primary reaction types discussed in this class are: 1. Neutralization reactions 2. Precipitation reactions 3. Oxidation-reduction reactions (redox reactions) a. Single replacement b. Combination reactions c. Decomposition reactions Some reactions can be classified as more than one of the reaction types listed above. Neutralization A reaction between an acid and a base to produce salt and water is called a neutralization reaction. In this experiment the only bases used will be those containing the hydroxide anion OH-. The generalized formula for this kind of neutralization reaction is: HA(aq)+ COH(aq) à HOH(l) and CA(aq) This kind of reaction of an acid and a base always yields water (HOH or H2O) and a salt. Example: Hydrochloric acid and sodium hydroxide react to produce sodium chloride and water. Since sodium chloride is soluble, its formation cannot be seen, but there is evidence that a reaction has taken place because the reaction mixture gets warmer. NaOH(aq) + HCl à NaCl(aq) + HOH(l) (Please note HOH is another way of writing H2O.) 67 The following reaction is a special type of neutralization reaction. This example illustrates how any carbonate will react with a strong acid to produce salt, water and carbon dioxide. Example: Bubbles of gas are formed when sodium carbonate reacts with hydrochloric acid. Na2CO3(aq) +HCl(aq) à CO2(g) + NaCl(aq) + H2O(l) Carbon dioxide and water is produced when any carbonate or bicarbonate reacts with any acid. This is also considered a neutralization reaction Precipitation reaction A reaction in which an insoluble substance is produced when two aqueous salt solutions are mixed together is called a precipitation reaction. A salt is a compound that consists of a positive ion, called a cation and a negative ion called an anion. When a salt dissolves in water, it separates into its constituent cations and anions. This is an aqueous salt solution and will be denoted with (aq). A list of names and formulae of the commonly occurring ions follows this section. You are solely responsible to learn each of the ions. Your success in this class depends heavily on this. Also following this background section is a set of solubility rules which will help you identify any insoluble solid precipitates formed during the reactions. You may also want to reference the nomenclature laboratory section; this may help you during the laboratory activity. Example: A white precipitate is formed when solutions of silver nitrate and sodium chloride are mixed. Write and balance the equation for the reaction showing the appropriate phases such as (s), (g), (aq) or (l). A. Write the formula for materials that will react in this reaction (the reactants). Silver nitrate is AgNO3 Sodium sulfite is Na2SO3 B. List all of the ions present in a solution of each of the two salts. Ag+ + NO3- and 2Na+ + SO32C. Write the formulae for the possible products when the salts exchange partner ions. (make sure that each cation is paired with a new anion) 2Ag+ pairs with SO32- and Na+ pairs with NO3Ag2SO3 and NaNO3 silver sulfite sodium nitrate D. Use the solubility rules to determine which of the possible products is insoluble and precipitates. The solubility rules (page 59) indicate that nitrates are soluble, so the sodium nitrate is soluble and does not form a precipitate. Silver sulfate is affected by rule #5: all sulfites are in soluble except those of ammonia and alkali metals. Silver sulfite forms a precipitate. The equation for the reaction can now be written and balanced: AgNO3(aq) + Na2SO3 (aq) à Ag2SO3 (s) + NaNO3(aq) The (s) indicates a solid. This is a precipitate. The (aq) indicates a aqueous solution, that is the two ions are dissolved in the water E. Write and balance the equation for the reaction showing the phases such as (s), (g), (aq) or (l). Example: A white precipitate is formed when aqueous solutions of sodium phosphate and calcium chloride are mixed. The final form of the equation would be: 2Na3PO4(aq) + 3CaCl2(aq) à Ca3(PO4)2(s) + 6NaCl(aq) 68 Oxidation-Reduction Reactions Any reaction in which electrons are transferred between reactants in the formation of the products is called an oxidation-reduction reaction (referred to as redox). Many reactions which also fit into other classifications involve a transfer of electrons and should be classified as oxidation-reduction reactions as well. Please note that any time an element is present as a reactant but then present in an ionic form as a product this is a redox reaction. When white hot sodium metal is placed into a flask filled with chlorine gas a very bright Example: reaction takes place. Both reactants are present as uncharged elements. After the reaction a salt is present which is made up of cations and anions. Each sodium has lost an electron to become a cation, each chlorine has accepted an electron to become an anion. Electrons were transferred from the sodium (leaving it with a positive charge) to the chlorine (making it negative). 2Na(s) + Cl2(g) à 2NaCl(aq) Redox: Single replacement reactions A very common form of a single replacement reaction is a reactive elemental metal reacting with a less reactive metal salt. In this single reaction it is possible to predict which metals can be substituted based on the metal’s reactivity. How metals react in a single replacement reaction depend on the metals’ position on the Activity Series. lithium potassium strontium calcium sodium The Activity Series of the metals is an invaluable aid to predicting the products of replacement reactions. It also can be used as an aid in predicting products of some other reactions. Pay attention to the notes below as they are provided to help you make better use of the activity series than just the list of metals by themselves. ------------------------------------------------------------ magnesium aluminum • CARBON zinc chromium • ----------------------------------------------------------- iron cadmium cobalt nickel tin lead • • ---------------------------------------------------------- HYDROGEN antimony arsenic bismuth copper -------------------------------------------------------- mercury silver paladium platinum gold • • Each element on the list replaces from a compound any of the elements below it. The larger the interval between elements, the more vigorous the reaction. The first five elements (lithium - sodium) are known as very active metals and they react with cold water to produce the hydroxide ion and hydrogen gas. The next four metals (magnesium - chromium) are considered active metals and they will react with very hot water or steam to form the oxide and hydrogen gas. The next six metals (iron - lead) replace hydrogen from HCl and dilute sulfuric and nitric acids. Their oxides undergo reduction by heating with H2, carbon, and carbon monoxide. The metals lithium - copper, can combine directly with oxygen to form the oxide. The last five metals (mercury - gold) are often found free in nature, their oxides decompose with mild heating, and they form oxides only indirectly. These single replacement reactions can also be completed using an acid. See example 2. 69 Example: When Zn metal is reacted with aqueous iron II sulfate, the zinc metal dissolves and an orange solid forms. Zn(s) + FeSO4(aq) à ZnSO4(aq) + Fe(s) Because these two metals are not far apart this reaction is not very energetic. Example: Bubbles of gas are formed when solid magnesium reacts with hydrochloric acid. Mg(aq) +HCl(aq) à MgCl2(aq) + H2(g) In this reaction the solid magnesium appears to disappear as it dissolves into an aqueous solution of magnesium chloride. Because magnesium is much higher in the series than hydrogen this reaction is fairly vigorous. Note: These single replacement reactions are the base of the electrochemistry of batteries. For instance batteries use the following materials to produce electricity: lithium and carbon (lithium ion), nickel cadmium (NiCad), Zinc and Carbon (Heavy Duty). Redox: Decomposition Reactions A single substance breaks apart to give two or more new substances. The generalized formula for this reaction is A à B + C Example: Hydrogen peroxide decomposes to form liquid water and oxygen gas. The hydrogen peroxide is initially dissolved in water this is called an aqueous solution. Hydrogen Peroxide (aqueous) à water (liquid) + oxygen (gas) The balanced equation is Fe 2H2O2(aq) à 2H2O(l) + O2(g) Please note the presence of the Fe above the reaction arrow. When something is located above the reaction arrow it is used as a catalyst in the reaction. Catalysts speed up a reaction but are not consumed in the reaction. Redox: Combination reactions Two or more substances combine together to form one new substance. The generalized formula for this reaction is A+BàC Example: When charcoal burns, carbon (charcoal) reacts with oxygen to form carbon dioxide Carbon(solid) + oxygen (gas) à carbon dioxide (gas) C(s) + O2(s) à CO2(g) 70 Names and Formulae for Commonly Occurring Ions (I would suggest you make one 3x5 flash cards for each ion and then memorize the cards. The cards should have the name on one side and the formula on the other) Cations 1+ Proton Lithium ion Sodium ion Potassium ion Silver ion Copper I Ammonium Anions 2+ + H Li+ Na+ K+ Ag+ Cu+ NH4+ Magnesium ion Calcium ion Strontium ion Barium ion Mercury I Mercury II Copper II Zinc ion Iron II Tin II Lead II Cobalt II Nickel II Manganese II 2+ Mg Ca2+ Sr2+ Ba2+ Hg22+ Hg2+ Cu2+ Zn2+ Fe2+ Sn2+ Pb2+ Co2+ Ni2+ Mn2+ Please note that there are other metal ions. For instance the formula for Manganese III is Mn3+. 3+ Aluminum ion Chromium III Iron III 3+ Al Cr3+ Fe3+ 4+ 2- - Hydride H Oxide Hydroxide OH Peroxide Fluoride FSulfide Chloride Cl Sulfite Bromide BrSulfate Iodide IThiosulfate Hypochlorite ClOCarbonate Chlorite ClO2 Chromate Chlorate ClO3 Dichromate Perchlorate ClO4Oxalate Cyanide CNCyanate CNOThiocyanate SCNNitrate NO3Nitrite NO2Permanganate MnO4Acetate CH3COOHydrogen carbonate HCO3(bicarbonate) Hydrogen Sulfate HSO4- (bisulfate) Peroxide O- 3- 4+ O2O22S2SO32SO42S2O32CO32CrO42Cr2O72C2O42- Nitride N3Phosphide P3Phosphate PO43- Tin IV Sn Lead IV Pb4+ Common Acids Hydrochloric acid Sulfuric acid Nitric acid Phosphoric acid Carbonic acid Acetic acid 1- Common Bases HCl H2SO4 HNO3 H3PO4 H2CO3 CH3COOH Sodium hydroxide Potassium hydroxide Calcium hydroxide Ammonia 71 NaOH KOH Ca(OH)2 NH3 Solubility Rules (I would suggest you memorize rule 1-3 at least) 1. Salts containing sodium (Na+), potassium (K+), and ammonium (NH4+) and acids are always soluble. 2. Salts containing nitrates (NO3-), chlorate (ClO3-), perchlorate (ClO4-), and acetate (CH3COO-) are always soluble. 3. All chlorides (Cl-), bromides (Br-), and iodides (I-) are soluble except for those of silver (Ag+), lead II (Pb2+), and mercury I (Hg22+), which are insoluble. 4. All sulfates (SO42-) are soluble except of those of strontium (Sr2+), barium(Ba2+), mercury I (Hg22+), mercury II (Hg2+), and lead (Pb2+). The sulfate salts of calcium (Ca2+), silver (Ag+) are moderately soluble. 5. All hydroxides (OH-) are insoluble expect those of alkali metals, which are soluble, and hydroxides of calcium (Ca2+) , barium (Ba2+), and strontium (Sr2+), which are moderately soluble. 6. All sulfites (SO32-), carbonates (CO32-), chromates (CrO42-) and phosphates (PO43-) are insoluble except those of ammonium (NH4+) and alkali metals. 7. All sulfides (S2-) are insoluble except those of ammonium (NH4+), the alkali metals and the alkaline earth metals, which are soluble. Please note: § Soluble means that the solid (s) will completely dissolve in water making an aqueous solution (aq). § Moderately soluble means that some of the material will dissolve, although some portion of it will remain as a solid. § Insoluble means that the solid will not dissolve and remain as a solid. Pre-Lab Questions: The stepwise process to writing a balanced reaction from a word description is a. b. c. d. Write the formula for materials that will react in this reaction (the reactants). List all of the ions present in a solution of each of the two salts. Write the formulae for the possible products when the salts exchange partner ions. Use the solubility rules to determine which of the possible products is insoluble and precipitates. e. Write and balance the equation for the reaction showing the appropriate phases such as (s), (g), (aq) or (l). Show the stepwise process towards writing a balanced equation for each reaction below. 1. Sodium phosphate solution will react with calcium chloride solution to form a white precipitate. (To confirm your balanced equation check the background information.) 2. A yellow precipitate is formed when solutions of silver nitrate and sodium phosphate are reacted. 3. Bubbles of gas are produced when a solution of potassium hydrogen carbonate (potassium bicarbonate) reacts with sulfuric acid. 4. A white precipitate appears and the solution gets warm when solutions of barium hydroxide and phosphoric acid are mixed. 72 Materials: Bunsen burner Magnesium metal 3 M KI 1.0 M HNO3 0.2 M Na2SO4 0.1 M Na2CO3 0.2 M NaCl ring stand crucible 1.0 M NaOH 0.2 M AgNO3 1.0 M HCl 0.2 M BaCl2 Iron Nail clay triangle 3% hydrogen peroxide 1.0 M H2SO4 0.2 M KBr 0.2 M CuSO4 Saturated Ca(OH)2 NaHCO3 (solid) Procedure: *******Note well******* In all cases, get material from the hood in labeled test tubes and then return to your laboratory bench. Make observations of the unreacted reagents, then conduct the experiment at your lab bench continue to make observations of the reaction. Combination Reactions 1a. and 1b. Reaction of magnesium and nitrogen and the reaction of magnesium and oxygen. When you burn magnesium it reacts both with oxygen in the air (≈ 20%) and nitrogen in the air (≈75%). You will only burn one strip of magnesium. When it burns it will complete 2 reactions. Set up a Bunsen burner, ring stand, and a clay triangle. Support a dry crucible (dry is far more important than clean in this experiment) on the triangle. Place 1 inch of loosely coiled magnesium metal in the crucible and heat the bottom of the crucible until the metal burns. It is important to make careful observations in this step as you will need them again in later experiments. !!!DO NOT LOOK DIRECTLY AT THE BURNING MAGNESIUM!!! Once the magnesium has completely burned, allow the crucible to cool to room temperature. Disturb the ash by stirring gently with a glass rod and describe the appearance of the ash in your laboratory notebook Decomposition Reaction 2. Decomposition of hydrogen peroxide This is the same reaction covered in the discussion. Without a catalyst this reaction proceeds extremely slowly. A catalyst is a substance that speeds up a reaction without being consumed by the reaction. You will use a catalyst to speed up the decomposition of hydrogen peroxide so that the bubbles of oxygen being produced can be seen. Place about 2 mL of 3% hydrogen peroxide in a small test tube. Add 3-4 drops of KI solution. Stir with a glass rod. Note the rate of evolution of bubbles of oxygen. Write careful observations into your laboratory notebook. Neutralization reactions Check each reactant with litmus paper before mixing, and check each mixture with litmus paper after the reaction has taken place. The proper technique is to dip a clean glass stirring rod into the solution and then touch the stirring rod to the litmus paper. Write the results of the tests with litmus paper into your laboratory notebook. 3. Sodium hydroxide and sulfuric acid 73 Place 2mL of aqueous sodium hydroxide in a small test tube. Do not measure this, simply use 40 drops or approximately 2mL as shown by your laboratory instructor. Carefully add 2mL dilute sulfuric acid. Stir with a glass stirring rod. Note any temperature changes by touching the test tube. It is not necessary to record the exact temperature change. 4. Nitric acid and sodium bicarbonate Place about 2 grams of solid sodium bicarbonate into an evaporating dish and then place about 2mL of aqueous hydrochloric acid onto the solid sodium bicarbonate. Please check the pH of the resulting solution using litmus paper. Please repeat experiment with aqueous sodium bicarbonate in a test tube. Precipitation Reactions Set up a warm water bath using a Bunsen burner, beaker, distilled water, a ring stand and a wire gauze. Do not let the water boil. Set up the small test tube rack containing six clean test tubes. Number each tube either with a pencil on its label spot or with a grease pencil. As you record the results of each reaction in your notebook, be sure to include the test tube number. Add about 2mL of each of the two reagents. It is not necessary to measure the volume of reagents. Write down all observations including colors of original reagent solution, amount of precipitate formed and color of precipitates. 5. Silver nitrate and potassium bromide 6. Sodium sulfate and barium chloride 7. Copper II sulfate and sodium hydroxide 8. Sodium carbonate and barium chloride 9. Copper II sulfate and sodium chloride 10. Calcium hydroxide and sulfuric acid (test before and after reaction with litmus paper) After you have run each reaction heat all test tubes in a water bath and make additional observations. Carefully record any and all changes and events that you observe, including a rough estimate of the time involved. !!! NEVER HEAT A TEST TUBE DIRECTLY IN A FLAME!!! Single replacement reactions 11. Copper II sulfate and iron Place about 2ml of copper II sulfate into a test tube. Then polish an iron nail with steel wool. Now place the polished nail into the copper II sulfate. Make careful observation regarding the rate of reaction and the differences between the submerged and not submerged portion of the nail. 12. Magnesium with hydrochloric acid Place a 1 cm strip of magnesium metal into an evaporating dish. Add about 2 mL of dilute hydrochloric acid. Write all observations into your laboratory notebook. 74 Data: This data table is to be completed before you come to class. I expect about four experiments per page; please align your table as follows (landscape). Use a ruler, be neat and above all make sure your table is arranged in numerical order (this is not necessarily the order in which you conduct the experiments). 1a 1b 2 3 4a 4b Nitrogen Oxygen Magnesium Mg N2 O2 Mg Reactant formula Magnesium Hydrogen peroxide Sodium hydroxide Sulfuric acid Sodium hydroxide (s) & Nitric acid Sodium hydroxide (aq) & Nitric acid Preobservations Observations Product names Product Formula Balanced equation (Include phase) 75 Calculations: No calculations for this lab Results: See data table for the balanced equations and names of products. Discussion: (please only address the following) In this experiment you did not perform any reactions labeled as oxidation-reduction. Explain how the reaction of magnesium burning in air and the reaction of magnesium with hydrochloric acid could both be called oxidation-reductions. Also explain why the precipitation reactions could not have been described as oxidation-reduction reactions. Post-Lab Questions: For each pair listed below, write a balanced equation and name all reactants and products. All reactions except 2c. involve aqueous salts listed above. 1. a. AgNO3 and KI b. CaCl2 and Na2SO4 c. H3PO4 and NaOH 2. a. barium hydroxide and nitric acid b. calcium hydroxide and hydrochloric acid c. burning aluminum metal and oxygen References: Finnegan, M. Place, H. Weissbart, B. (2000) Washington state university chemistry 101-102 laboratory manual. Star Publishing Company: Belmont, California 76 Experiment 11 Stoichiometry Purpose: Determine the limiting reactant and predict mass of products. Hypothesis: Develop a hypothesis regarding the conservation of mass in this reaction. Background Information: In this experiment you will compare the experimentally measured and the theoretical amount (in grams) of barium sulfate produced by the reaction of a barium chloride dihydrate solution and a potassium sulfate solution. The reaction involved is shown below. K2SO4(aq) + BaCl2·2H2O(aq) à BaSO4(s) + 2KCl(aq) + 2H2O(l) (1) The barium sulfate forms a precipitate while potassium chloride remains in solution. The barium sulfate precipitate is collected by gravity filtration because the water and the potassium chloride will pass through the filter paper (see figure 1). Figure 1 Gravity Filtration and filter paper preparation The filter paper and precipitate are placed together in a beaker or on a watch glass. The filter paper and precipitate can then be air or oven dried. The experimental mass of barium sulfate can then be determined and compared to the theoretical mass of barium sulfate. This reaction is a rarity because conversion to product is essentially complete. Most reactions do not yield complete conversion to product. In this case, however, the experimental and the theoretical yield should be fairly close. In fact the experimental amount of barium sulfate is often slightly larger than the theoretical amount of barium sulfate because the precipitate is contaminated with a small amount of the filtrate which contains potassium chloride. This contaminated filtrate could be eliminated by washing with water. But this process is very time consuming. At the end of the experiment you will need to determine the accuracy of the experiment. In this experiment accuracy is defined as the agreement between the experimental and theoretical values, in grams, of barium sulfate. We will measure accuracy by the absolute error and the relative error. Absolute error is defined by equation (2). Absolute error = E = |m-t| Where E = absolute error m = experimentally measured value t = theoretical value 77 (2) % relative error is defined by (equation 3). % relative error = |m-t|/t *100 (3) A related equation determines the % yield (equation 4). m/t*100= % yield (4) Limiting reactant: You will be asked to determine which material limited the reaction. Think of this as the material that ran out first. For instance in building bicycle if you have 200 bicycle frames and 300 wheels. Which runs out first, assuming each frame needs two wheels. The answer may seem obvious. However when you start working with grams of chemicals it can seem harder. Back to the example, another important question might be how many seats are needed? Here is a step by step explanation of how to solve a limiting reactant problem Step 1: Write and balance equation Step 2: Determine the number of moles of each reactant Step 3: Check the first reactants and determine how many moles of the second reactant is needed If you have enough moles of the second reactant go to step 5 if not go to step 4 If you have enough moles of the second the first reactant is the Limiting reactant • To reactant with _______ moles of (reactant 1) it necessary to have ________ moles of (reactant 2). • There are only ____________moles of (reactant 2). • There are not enough moles of (reactant 2) to react with all of the ______________ reactant 1). • Reactant 2 is the limiting reactant. Step 4: Check the second reactant and determine how many moles of the first reactant are needed. Step 5: Write down the identity of the limiting reactant. Step 6: Identify from above how many moles of the limiting reactant will be used in the reaction. Step 7: Convert from mole of limiting reactant to grams of the other reactant. Step 8: Convert from mole of the limiting reactant to mole of the first product. Step 9: Convert from moles of limiting reactant to moles of second product Step 10: Determine grams of excess reactant Step11: Check grams of material before reaction compare to grams material after reaction Another way of thinking about this problem is this It may be helpful to organize your thoughts and focus your efforts. In the calculations, ask and answer these questions. A.) If I have ________ grams of K2SO4 how many grams of BaCl2*2H2O do I need?________ I have this many grams of BaCl2*2H2O __________. If there are fewer grams than needed then it is the limiting reactant. B.) If I have ________ grams of BaCl2*2H2O how many grams of K2SO4 do I need?________ I have this many grams of K2SO4 _________. If there are fewer grams than needed then it is the limiting reactant The reactant that ran out first is _______________ . By determining the limiting reactant you can then predict how much of the product can be produced. 78 Pre-Lab Questions: 1. Use the provided experimental data to answer the following questions. 0.987g of potassium sulfate are reacted with 0.897 grams of barium chloride dihydrate. a. Write the balanced equation. b. How many grams of each reactant and product do you start with? (hint you don’t start with any product) c. What is the molecular mass of each reactant and product? d. How many moles of reactant and product do you start with? e. What is the limiting reactant? f. How many moles of each reactant and product do you end up with? g. How many grams of each reactant and product do you end up with? 2. Use the provided experimental data to fill in the following table. 0.925 grams of potassium sulfate are reacted with 1.125 grams of barium chloride dihydrate. K2SO4(aq) + Grams @ start Molar mass Moles @ start Change Moles @ end Grams @ end BaCl2*2H2O(aq) à BaSO4(s) + KCl(aq) +H2O(l) Grams @ start à Grams @ end 3. In the lab it indicates in step 4 and 8 to add approximately 30 ml and 10 ml water respectively. Yet it does not ask you to record your volume. Why does it not seem to matter how much water you add? Why does the volume of water not seem to affect the experiment? 4. If, at the end of the experiment, a % yield greater than 100 is calculated what might be the source of the error? What could be done to correct this error? Materials: Potassium sulfate (approximately 1.0g) Barium chloride dihydrate (approximately 1.0g) Wooden filtration rack Filter paper Wash bottle Bunsen burner Wire gauze Ring stand Ring clamp Grease pencil Barium sulfate waste receptacle Procedure: 1. 2. 3. 4. 5. 6. 7. 8. 9. Mass a clean dry 150 ml beaker. Then record mass directly into lab notebook. Obtain approximately 1.0 grams of K2SO4 and place in beaker. Record the mass of the beaker and sample directly into the notebook. Label beaker with your initials and the contents with pencil or tape. Add approximately 30 ml of distilled water and stir the solution. Bring solution to a gentle boil. Mass another clean dry 150 ml beaker. Then record mass directly into lab notebook. Obtain approx. 1.0 g of barium chloride dihydrate and place in beaker. Record the mass of the beaker and the sample directly into the lab notebook. Add 10ml of deionized water to beaker and dissolve the barium chloride dihydrate. The precipitate is formed by adding the barium chloride solution very slowly to the hot sulfate solution. To mix this reactants together very slowly use a disposable pipette or burette. Consider trying one drop every few seconds. By mixing the solutions slowly it may avoid the formation of very small barium sulfate particles. 79 10. Rinse the barium chloride dihydrate solution beaker with another 10 ml of D.I. water to remove any residual barium chloride dihydrate and transfer this rinse solution to the potassium sulfate solution beaker also. The white precipitate that forms is barium sulfate. 11. After the reagents have been mixed. Heat the solution over a Bunsen burner (do not boil), for about 15 to 20 minutes. After this time the barium sulfate precipitate should be at the bottom of the beaker under a clear solution. (This may be a good time to stop for the day.) 12. Obtain a sheet of filter paper; put your initials on the paper in pencil. Mass the filter paper to the nearest 0.001 g on the same analytical balance previously used for your other masses and record the results directly in your lab notebook. 13. Set up the gravity filtration apparatus described in Figure 1. Have an assistant or instructor show you how to correctly prepare the filter paper. Make sure that you do not punch a hole in the filter paper. If there is a hole in the filter paper, the barium sulfate will not be separated from the water and the other product, potassium chloride. NOTE: There will be a number of times during this experiment when you are not busy. Use this “dead” time to begin the analysis of the data. Calculate the number of moles of each reagent. Begin to complete the table in the Results Section. Note that the barium chloride is the dihydrate, barium chloride* water. Be sure to use the correct molecular weight. Choose the limiting reagent. On the basis of this result, complete the table in the Results Section. If you can complete this calculation during the lab period, the subsequent lab report will be simplified greatly. 14. Pour most of the clear, hot solution into the filter. Do not let the liquid level rise above the edge of the paper. Keep the remaining solution and solid hot while waiting for the filter to drain. When only a few mL remain in the beaker, agitate to suspend the Barium sulfate and pour the suspension into the filter. Do not let liquid level in filter rise above the edge of the paper. 15. Rinse the sides of the beaker with a few milliliters of deionized water, using a wash bottle. Transfer the rinsing into the filter. Rinse the beaker as required until the last of the barium sulfate is in the filter. This is a stopping point; if you must leave, place your funnel in a beaker so the filtration can continue to completion until the following lab period. 16. When the filtration is complete, rinse the filter cake with a few milliliters of deionized water. Do not let liquid level rise above the edge of the paper. 17. Air dry the filter paper & barium sulfate in your drawer until the following lab period then oven dry for about 15 minutes or as designated by the instructor. 18. Mass the filter paper & barium sulfate once they are dry on the same analytical balance previously used. Record the mass directly into your notebook. Observation Data: 1. mass of beaker and potassium sulfate mass of empty beaker mass of potassium sulfate g g g 2. mass of beaker and barium chloride mass of empty beaker mass of barium chloride g g g 3. mass of filter paper and barium sulfate mass of filter paper mass of barium sulfate g g g 80 Calculations: Show your calculations for the theoretical mass of barium sulfate produced, the mass of potassium chloride and the mass of the excess reagent. Use unit analysis and use the proper number of significant figures. Clearly indicate the limiting reagent. Table 1 Grams @ start Molar mass Moles @ start Change Moles @ end Grams @ end K2SO4(aq) + BaCl2*2H2O(aq) à BaSO4(s) + KCl(aq) +H2O(l) Theoretical yield “Change” is determined by first calculating the moles of the limiting reagent and applying it to the other reactants and products. a. Compare the theoretical and experimental masses of barium sulfate produced. b. Calculate the absolute error. c. Calculate the relative error. d. Fill in data from the above table. e. Label all arithmetic calculation pertaining to mass determination. Results: State the theoretical yield. (from Table 1) State the experimental yield. Report the absolute error. Report the relative % error. Report the % yield. Discussion: Briefly summarize your results and discuss your relative errors. If your percent yield is greater than 100 explain what factors may have caused this error. If your percent yield is less than 90 percent explain what may have happened to cause this error. Be very specific in this section. For example discuss the how the results would be effected for instance losing material before or after the initial massing. What would happen if it was not completely dried out? It is unsafe to place barium salts into the sink . How do you if it is safe to pur th filtrate down the sink. Answer the other discussion questions as presented in your “How to Write a Lab Report.” See next page for Post Lab Questions. 81 Post-Lab Questions: Soda ash (sodium carbonate) is widely used in the manufacture of glass. Prior to the environmental movement much of it was produced by the following reaction. CaCO3 + 2 NaCl à Na2CO3 + CaCl2 Unfortunately, the byproduct calcium chloride is of little use and was dumped into rivers, creating a pollution problem. As a result of the environmental movement, all of these plants closed. Assume that 125g of calcium carbonate (100.09 g/mol) and 125 g of sodium chloride (58.44 g/mol) are allowed to react. Construct a table similar to that shown in the calculation section. a. Determine how many grams of useful sodium carbonate (105.99 g/mol) will be produced. b. How many grams of useless calcium chloride (110.98 g/mol) will be produced? c. You should also determine how many grams of excess reagent are left. Show all calculations; use dimensional analysis; and use the proper significant figures. Use a table if you wish to organize your data. Clearly show how you chose the limiting reactant. References: Collins, V. Kahl, D. Perry, F. (1996) Good stuff from the chemistry laboratory. Warren Wilson College press: Swannanoa NC Silberberg, Martin. (1996) Chemistry the molecular nature of matter and change. Mosby: New York, NY 82 Experiment 12 Determining the Concentration of a Solution: Beer’s Law The primary objective of this experiment is to determine the concentration of an unknown nickel (II) sulfate solution. You will be using the Colorimeter shown in Figure 1. In this device, red light from the LED light source will pass through the solution and strike a photocell. The NiSO4 solution used in this experiment has a deep green color. A higher concentration of the colored solution absorbs more light (and transmits less) than a solution of lower concentration. The Colorimeter monitors the light received by the photocell as either an absorbance or a percent transmittance value. Figure 1 Figure 2 You are to prepare five nickel sulfate solutions of known concentration (standard solutions). Each is transferred to a small, rectangular cuvette that is placed into the Colorimeter. The amount of light that penetrates the solution and strikes the photocell is used to compute the absorbance of each solution. When a graph of absorbance vs. concentration is plotted for the standard solutions, a direct relationship should result, as shown in Figure 2. The direct relationship between absorbance and concentration for a solution is known as Beer’s law. The concentration of an unknown NiSO4 solution is then determined by measuring its absorbance with the Colorimeter. By locating the absorbance of the unknown on the vertical axis of the graph, the corresponding concentration can be found on the horizontal axis (follow the arrows in Figure 2). The concentration of the unknown can also be found using the slope of the Beer’s law curve. Objectives: In this experiment, you will • • • • Prepare NiSO4 standard solutions. Use a Colorimeter to measure the absorbance value of each standard solution. Find the relationship between absorbance and concentration of a solution. Use the results of this experiment to determine the unknown concentration of another NiSO4 solution. 83 Materials: computer Vernier computer interface Logger Pro Vernier Colorimeter one cuvette five 20 Í 150 mm test tubes tissues (preferably lint-free) stirring rod 30 mL of 0.40 M NiSO4 5 mL of NiSO4 unknown solution two 10 mL pipets (or graduated cylinders) pipet pump or pipet bulb distilled water test tube rack two 100 mL beakers Procedure: 1. Obtain and wear goggles! CAUTION: Be careful not to ingest any NiSO4 solution or spill any on your skin. Inform your teacher immediately in the event of an accident. 2. Add about 30 mL of 0.40 M NiSO4 stock solution to a 100 mL beaker. Add about 30 mL of distilled water to another 100 mL beaker. 3. For each trial place the volume of NiSO4 required into a graduated cylinder. Record the actual value of NiSO4. Then add distilled H2O to bring the volume up to 10.00mL Target volumes. Your actual volumes will vary. Trial number 1 2 3 4 5 0.40 M NiSO4 (mL) 2 4 6 8 ~10 Distilled H2O (mL) 8 6 4 2 0 Final Vol Concentration (M) 10 10 10 10 10 0.08 0.16 0.24 0.32 0.40 4. Connect the Colorimeter to the computer interface. 5. You are now ready to calibrate the Colorimeter. Prepare a blank by filling a cuvette 3/4 full with distilled water. To correctly use a Colorimeter cuvette, remember: • • • • All cuvettes should be wiped clean and dry on the outside with a tissue. Handle cuvettes only by the top edge of the ribbed sides. All solutions should be free of bubbles. Always position the cuvette with its reference mark facing toward the white reference mark at the top of the cuvette slot on the Colorimeter. 6. Calibrate the Colorimeter. a. Open the Colorimeter lid. b. Holding the cuvette by the upper edges, place it in the cuvette slot of the Colorimeter. Close the lid. 84 c. If your Colorimeter has a CAL button, Press the < or > button on the Colorimeter to select a wavelength of 635 nm (Red) for this experiment. Press the CAL button until the red LED begins to flash. Then release the CAL button. When the LED stops flashing, the calibration is complete. 7. You are now ready to collect absorbance data for the five standard solutions. Click . Empty the water from the cuvette. Using the solution in Test Tube 1, rinse the cuvette twice with ~1 mL amounts and then fill it 3/4 full. Wipe the outside with a tissue and place it in the Colorimeter. After closing the lid, wait for the absorbance value displayed on the monitor to stabilize. Then click . 8. Discard the cuvette contents as directed by your teacher. Rinse the cuvette twice with the Test Tube 2 solution, 0.16 M NiSO4, and fill the cuvette 3/4 full. Wipe the outside, place it in the Colorimeter, and close the lid. When the absorbance value stabilizes, click . 9. Repeat the Step 8 procedure to save and plot the absorbance and concentration values of the solutions in Test Tube 3 (0.24 M) and Test Tube 4 (0.32 M), as well as the stock 0.40 M NiSO4. Wait until Step 12 to do the unknown. When you have finished with the 0.40 M NiSO4 solution, click . 10. In your Data and Calculations table, record the absorbance and concentration data pairs that are displayed in the table. Using the data pairs create a graph. 11. From the data points on your graph chart a best fit line. From this best fit line use Y=mX+B to write a equation for the line. 12. Obtain about 5 mL of the unknown NiSO4 in another clean, dry, test tube. Record the number of the unknown in the Data and Calculations table. Rinse the cuvette twice with the unknown solution and fill it about 3/4 full. Wipe the outside of the cuvette, place it into the Colorimeter, and close the lid. Read the absorbance value displayed in the meter. (Important: The reading in the meter is live, so it is not necessary to click to read the absorbance value.) When the displayed absorbance value stabilizes, record its value in Trial 6 of the Data and Calculations table. 13. Using the equation for the line and the unknow data find the concentration for the unknown. Data and Calculations: Trial 0.40 M NiSO4 Distilled H2O (mL) (mL) Final Vol Conc (mol/L) Absorbance 1 2 3 4 5 6 Unknown number ____ Concentration of unknown mol/L 85 Calculations: Use M1V1=M2V2 to find the concentration for each trial. M1 = Molarity intial V1 = Volume final M2 = Molarity final V2 = Volume final Based on your data us Y=mX+b to find the slope of the line. Rearrange equation to solve for X. Insert the value of the unknown as Y (absorbation) solve for X. Results: Graph the known values Place the unknown material on the graph Report the unknown concentration and ID number. Discussion: Discuss how this technique could be used in a practical application in either an industry, research or medical application. Propose an experiment where you could use this technique to measure the concentration of something you are interested in. 86 Experiment 13 Determination of Sodium Hydroxide Concentration Purpose: Develop titration skills. Use the titration technique to determine the concentration of a unknown solution. Background: Solid sodium hydroxide has a tendency to absorb water from the atmosphere. The longer sodium hydroxide is exposed to air the more water it absorbs. This characteristic of sodium gydroxide makes it impossible to make a standard solution by weight alone. For instance if you wanted to make a 1 molar solution of sodium hydroxide all that is needed is to demtermine the molecular weight of Sodium Hydroxide (40.00 grams/mole). Thus you should be able to just place 40.00 grams of NaOH into a container and then bring the water level up to 1.00 liters. However this does not work. If you did what was previously outlined a less then 1.00 molar solution is likely to be created. Because when 40.00 grams of sodium hydroxide is massed there an unknown quantity of water that has been absorbed into the NaOH crystals. In what is supposed to be 40.00 grams of NaOH there might only be 32.00 grams of NaOH and 8 grams of water. To determine the actual concentration of NaOH in a solution it is necessary to complete a stoichiometic reaction called a titration. A titration is when a known concentration of a substance is used to determine an unknown concentration. There are many types of titrations. To determine the concentration of NaOH an Acid-Base Titration will be completed. The Nature of Acids and Bases There are multiple definitions of what constitutes an acid and a base. However the most basic definition will be sufficient for this laboratory activity. The Arrhenius Acid-Base definition asserts that: • An acid contains hydrogen and dissociates in water to yield H2O and H+ (sometimes written as H3O+). • A base contains an OH group and dissociates in water to yield OH-. Acids and bases in general focus on the H+ (proton) concentration in water. This concentration is expressed in moles/liter. But because this occurs in water it is important to note that a glass of water just sitting there minding its own business, will undergo the following reversible reaction all on its own without any help. H 2O à ß H+ + OH- Please note that the arrow towards products is very large. This is to indicate that the reaction from water to its ions is pretty infrequent. But it does happen. It is possible to measure just how much H+ has dissociated. In any given sample of water the H+ concentration is 1.00 x10-7 mole H+/liter of water That is not a lot of H+ floating around in the water. It was decided that a scale should be made to help people understand how acidic or basic things are. It was further decided that on this scale water should be considered neither acidic, nor basic, it would be neutral. Thus the scale should be easy to understand. 87 We use the pH scale for this purpose. It goes from around 0 to 14. Small numbers on this scale have a very high proton concentration and a low hydroxide ion concentration and are called acidic. Large numbers on this scale have a high hydroxide ion concentration and a low proton concentration and are called basic. pH 0 1 2 3 4 5 + Protons A lot of H 1.0 M H+ hydroxide Practically no OH1x10-14M OH- 7 6 8 9 10 11 12 13 14 - Practically no H+ 1x10-14M H+ Equal to H+ 1x10-7M OH- A lot of OH1.0 M OH- Equal to OH 1x10-7M H+ So how do you go from some simple scale of 0-14 to a molar concentration? To do this you need to understand what pH means. The p= the –log The H= the H+ concentration in mol/liter (which is often written as [H+]) pH= -log[H+] If you want to know what the pH of a solution that has a H+ molarity of 1.0 x10-7 mole/liter (or [H+]=1.0x10-7 M) all you do is pH=-log1.0x10-7 The calculator syntax for converting concentration to pH the TI 80 series calculator is (-) LOG 1 EE (-) 7* ENTER *The number following the log is the concentration of the solution in moles/liter H+ To calculate the molarity from the pH you just do it in reverse (though it is slightly harder). Let us find the concentration of H+ in a solution with a pH of 7 pH= -log[H+] 7= -log[H+] (-1)*7= (-1)*-log[H+] -7= log[H+] to get rid of the log you 10x both sides (10log = 1) 10-7 = 10log[H+] 10-7= 1*[H+] and 1x10-7= [H+] The calculator syntax for pH to concentration on the TI 80 series calculator is 2nd * LOG(10x) 7* (-) ENTER is the pH of any solution you are interested in find the H+ concentration of Because the pH scale is a logarithmic function the difference in pH between pH 7 and pH 6 is a factor of 10. That is, a solution with a pH of 6 has 10 times more H+ than a solution of pH 7. 0 1 2 3 4 5 6 7 10,000,000 1,000,000 100,000 10,000 1000 100 10 1 unit H+ 88 A solution of pH 0 has 100,000,000 times more H+ as a solution with a pH of 7. The same is true for numbers greater than seven each pH level increase H+ concentration is reduced by a factor of ten and the OH- concentration is increased by a factor of 10. Indicators An acid base indicator is a weak organic acid whose acid form, which will be denoted as HA(indicator) is a different color than its base form A-(indicator), with the color change occurring over a specific pH range. Typically, one or both forms are intensely colored, so only a tiny amount of indicator is needed and the presence does not affect the pH of the solution. Indicators are used for approximate pH monitoring in acid base titrations or in reactions. HA(indicator) à ß H+ + A-(indicator) The double arrows of this reaction indicate that it is a reversible reaction. A reversible reaction can go in either direction. The direction that is favored is often based on which side has the highest concentration. In the above example if there is a lot of H+ then the direction is likely to be moving toward the reactants. If the H+ concentration is very small the reaction will move in the direction of products. Consider phenol red as an indicator. When it is in the HAindicator form it is yellow. When it is in the A-indicator it is red. When you place phenol red in an acid solution there are very large quantities of H+. Because of the high concentrations of H+ the H+ on the indicator will not dissociate and the reaction moves in the direction of reactants. Thus there are high concentrations of HAindicator. Yellow Red HAindicator ß H+ + A-indicator à When you place phenol red in a basic solution there is a very low concentration of H+. This forces the reaction to favor products. Red Yellow HAindicator à H+ + A-indicator ß When phenol red is in high concentrations of H+ (an acidic solution) the primary species is HAindicator and the color is yellow. When phenol red is in low concentrations of H+ (a basic solution) the primary species is A-indicator and the color is red. 89 Some commonly used indicators: 1 Thymol Blue 2 Red Orange Methyl red Bromothymol Blue 3 4 pH 5 6 Yellow 7 8 9 Green 10 11 Blue Green 12 Red Orange dfds Yellow Yellow sdf Phenol red Green Blue Yellow sdfs Red Orange Clearsdf Pink Phenolphthalein Acid Base Reactions A frequently used definition of the term acid and base is attributed to BrØnsted and Lowry An acid is any substance which donates protons (H+ ions) A base is any substance which accepts protons When an acid dissolves in water, some or all of the protons are given up. These bare H+ ions quickly react with water to form an ion usually written as H3O+ (hydronium ion). Thus when hydrogen chloride dissolves in water to form Hydrochloric acid the process maybe presented as: HCl + H2O à Cl- + H3O+ This dissolution process is an acid-base reaction with hydrogen acting as an acid and water acting as a base. The acid is donating protons (HClàCl- + H+) and the base is accepting protons (H2O + H+ à H3O+). When this solution is mixed with another base which will accept protons it is the hydronium ion H3O+, which acts as an acid. When a weak acid dissolves in water, the proton transfer is not complete. For example, the weak acid HNO2, dissociates partially. Only a small fraction of the HNO2 molecules transfer their protons to water. HNO2 + H2O à NO2- + H3O+ Many bases form the hydroxide ion, OH- in water. NaOH à Na+ +OHThe reaction of an acid with a base is a proton transfer which converts both reacting species into water molecules. H3O+ + OH- à 2H2O The overall reaction of hydrochloric acid with sodium hydroxide can be written: H3O+ + Cl-+ Na+ + OH- à 2H2O + Na+ + ClThe formula of the salt produced is NaCl and the name of the salt is sodium chloride. Similarly, when a weak acid reacts with hydroxide ions, the H+ ions are pulled of the acid. The reaction is complete when all of the acid has lost its H+ ions or when the base is used up. An example is given below for the reaction of NaOH and HNO2: HNO2 + NaOH à H2O + Na+ + NO2- 90 Titration A titration is a process used to determine the volume of a solution needed to react with a given amount of another substance. In this experiment, you will titrate hydrochloric acid solution, HCl, with a basic sodium hydroxide solution, NaOH. The concentration of the NaOH solution is given and you will determine the unknown concentration of the HCl. Hydrogen ions from the HCl react with hydroxide ions from the NaOH in a one-to-one ratio to produce water in the overall reaction: H+(aq) + Cl–(aq) + Na+(aq) +OH–(aq) H2O(l) + Na+(aq) + Cl–(aq) Typical titration setup Titration calculation 1. Determine moles of known (acid in this case) 2. Convert moles of known to moles of unknown (acid to moles of base) 3. Convert moles of unknown to Molarity (to find molarity of the base) For this example assume that 10.00 mL of a 0.200 Molar solution of HCl was used to neutralize 15.25 mL of sodium hydroxide. Moles of acid (the known) The number of moles of acid can be found using the morality and volume data. Thus: 10.00 mL HCl 1 Liter HCl 0.200 moles HCl =2.00 x10-3 moles HCl 1000 mL HCl 1 Liter HCl Equation 1a Moles of Base (the unknown) From the balanced equation NaOH + CH3COOH à NaCH3COO + H2O please observe that for every one mole of NaOH 1 mole of CH3COOH is neutralized Thus: 2.00 x10-3 mol HCl 1 mol NaOH 1 mol HCl = 2.00x10-3 mol of NaOH Equation 1b Convert moles of base to molarity Now that the moles of base has been determined it is not possible to use the volume of base measured using the burette to determine the molarity. Keep in mind molarity is a measure of the number of moles per liter of solution. Thus: 1000 mL NaOH 2.00x10-3 mol of NaOH = 0.13114 = 0.131 M NaOH 15.25 mL NaOH mole liter 1 liter NaOH Equation 1c Or you can just do it all at once 10.00 mL HCl 15.25 mL NaOH 0.200 mol HCl 1000 mL HCl 1 mol NaOH 1 mol HCl 1000 mL NaOH 1 liter NaOH = 0.131 M NaOH Equation 1d From the data in the lab Known molarity of the acid From the balanced equation, it is not always a 1:1 relationship Converting mL to Liters is necessary M= Mole/liter of solution 91 Statistical Analysis The first is a rough titration this is only to determine the approximate quantities needed. This data is discarded. You will run at least three experimental trials. Then statistically analyze these three trials to determine the mean and the average deviation from the mean. Statistics allow us to determine the quality of the data. One reason random errors occur is due to the estimated nature of all measurements. Errors occur whenever measurements are made. It is useful to know how large the errors are. In this experiment precision of the results will be determined. First perform the experiment three or more times. Then calculate the average or mean as shown in Equation 3. _ X1 + X2 + X3 Mean = X n = Where X1 is the value determined from experiment trial 1 etc n= the number of trials Equation 3 If an experiment has good precision, all of the results will be clustered closely around the mean or average. If the experiment has poor precision, the results will be widely scattered around the mean. Precision can be determined by calculating the deviation from the mean. Deviation is defined in equation 4 as: _ i takes on the values of 1,2,3… for each trial δi = |Xi –X| Equation 4 _ X = the mean for all trials δi = the deviation for measurement i Xi = the value of measurement i The average deviation provides a measurement of how closely each of your datas points are to each other. This is defined below in equation 5 _ |δ 1| + |δ 2| + |δ 3| Average deviation = δ = n Equation 5 Where |δi| = absolute value of deviation (no algebraic sign is used) n = number of measurements Precision is often calculated as the relative average deviation. To determine the overall precision for the experimental data it is necessary to take the average deviation and divide it by the mean (found in equation 3) then multiple this number by 100. See equation 6. _ Equation 6 δ _____ Precision = relative average deviation = X * 100 92 If the relative average deviation is 5% or less, the data is acceptable. If the relative average deviation is 2% or less, the precision is good. If the precision is 1% or less, the quality of the data is outstanding. Pre-Lab Questions: Write the balanced equation for each of the following acid base reactions. In each case, you are expected to predict the formula of the salt produced. 1. Nitric acid reacted with sodium hydroxide 2. Sulfuric acid reacted with potassium hydroxide 3. Acetic acid reacted with calcium hydroxide Materials: NaOH unknown concetration ~ 0.200 M Burette clamp Pipette pump Phenolphthalein 150 mL beaker Hydrochloric acid Burette 10.00 volumetric pipette 100 mL standardized sodium hydroxide 125 mL Erlenmeyer flask Ring Stand Procedure: 1. Obtain a burette, burette clamp and large ring stand from the lab bench. Set up material as per instructor’s demonstration. Describe this laboratory apparatus in your procedure section. 2. Record the molarity of the hydrochloric acid solution directly into your lab notebook. 3. Obtain in a clean and dry beaker enough sodium hydroxide to fill the burette, Only take what you need. Do not waste the sodium hydroxide. If the class runs out there will be no more to complete your experiments. 4. Ask the instructor if the burette needs to be cleaned. If the burette needs cleaning rinse the burette three times with 10 mL of the solution. After each rinsing, allow the solution to drain through the stopcock. Then fill the burette so that the liquid level is above the zero line. 5. With the burette in the clamp place the zero line at eye level and the end of the burette over an empty beaker. Slowly drain the solution until the liquid level is between 1.00 and 2.00 mL. When finished, wipe the tip of the burette. 6. Transfer 10.00 mL Hydrochloric acid to a clean but not necessarily dry 125 mL Erlenmeyer flask. Use the technique demonstrated in class. Describe this technique in your lab note book in the procedure section. 7. Add two to three drops of phenolphthalein indicator to the Erlenmeyer flask. Now add approximately 50 mL of distilled water to the flask. Make sure to use the same amount of phenolphthalein and water for each trial. 8. Record the NaOH burette reading as the initial burette volume. Proceed to titrate the sample. The indicator will change from colorless to pink when the equivalence point has been reached (the color will change when the moles of H+ are the same as the moles of OH-). When the indicator changes to pink and the color persists for about 30 seconds, you have completed the titration. 93 Note: Titrate rapidly at first. When the pink color begins to persist, slow the addition of sodium hydroxide to about one drop per second. As the equivalence point is approached, add sodium hydroxide so that only one drop of NaOH is needed to turn the solution pink. If you happen to overshoot the end point, do not despair, just start over. 9. Record the burette reading. If the volume is less than 10mL, you should read to 3 sig figures; example, 6.35 mL. If the volume is greater than 10mL you should read to 4 sig figures; example, 14.31mL. Record the volume directly into your laboratory notebook. If you have trouble reading the burette to the proper number of significant figures, ask the instructor for assistance. 10. After the titration is complete, you may pour the contents of the Erlenmeyer flask down the drain with running water. 11. Repeat the titration as needed until you have a total of at least three titrations. The excess sodium hydroxide solution and acid should be drained into the sink with running water. Data: Stated Molarity of base___________ (this is the theoretical molarity of the base) Molarity of hydrochloric acid _____________ Trial one Volume of HCl _____ Final burette vol __________ Initial burette vol __________ Vol of NaOH ___________ Observations: Trial 2 Volume of HCl _____ Final burette vol __________ Initial burette vol __________ Vol of NaOH ___________ Observations: Trial 3 Volume of HCl _____ Final burette vol __________ Initial burette vol __________ Vol of NaOH ___________ Observations: DO NOT RECORD DATA IN YOUR Include color of when titration is completed LABORATORY MANUAL Calculations: Trial 1 Determine moles of base (see equation 1a) ________ Trial 2 Determine moles of base (see equation 1a) ________ Trial 3 Determine moles of base (see equation 1a) ________ Determine moles of acid (see equation 1b) ________ Determine moles of acid (see equation 1b) ________ Determine moles of acid (see equation 1b) ________ Determine molarity of Determine molarity of Determine molarity of unknown unknown unknown (See equation 1c) _________ (See equation 1c) _________ (See equation 1c) _________ Statistical analysis A. Molaritytrial1 ________ above) Molaritytrial2 _______ molaritytrial3 ________ ß (From calculations ß (see equation 3) B. Molarity mean ___________ 94 Trial 3 δ3=_______ ß(see C. Deviation trial 1 δ1= _______ equation 4) Trial 2 δ2=_________ D. Standard deviation ____________ ß (see equation 5) E. Relative average deviation ____________ ß (see equation 6) F. Percent error of the base ______________ ( % error calculations can be found in previous lab) Results: 1. Molarity of sodium hydroxide Trial 1 ___________ Trial 2 ___________ Trial 3 ___________ 2. Deviation Trial 1 ___________ Trial 2 ___________ Trial 3 ___________ 3. Average Deviation _________________ 4. Relative average deviation _________________ Discussion: • • Discuss your relative average deviation. If it is greater than 1% explain possible sources of error. Discuss topics presented in the discussion section of “How to write a lab report” Post-Lab Questions: 1. 10.00mL of a solution of hydrochloric acid requires 12.81 mL of a solution of 0.1365 M sodium hydroxide to neutralize it. Give the concentration of the hydrochloric acid in molarity and mass %. Assume that the hydrochloric acid solution has a density of 1g/mL. 2. You have used 25.38 mL of 0.1073 M potassium permanganate to titrate 9.86 mL of hydrogen peroxide. Given the balanced equation below and the fact that the molecular weight of hydrogen peroxide is 34.01 g/mole and the density of hydrogen peroxide is 1.000g/mL, calculate the molarity and the % mass of the hydrogen peroxide. Show all work; use unit analysis, the proper significant figures, etc. 3H2SO4 + 2 KMnO4 + 5H2O2 à 2MnSO4 + K2SO4 + 5O2 + 8H2O (Hint: you can disregard all information save the mole relationship between the potassium permanganate and the hydrogen peroxide.) References: Collins, V. Kahl, D. Perry, F. (1996) Good stuff from the chemistry laboratory. Warren Wilson College press. Swannanoa NC Finnegan, M. Place, H. Weissbart, B. (2000) Washington state university chemistry 101-102 laboratory manual.Star Publishing Company: Belmont, California Silberberg, Martin. 1996. Chemistry the molecular nature of matter and change. Mosby: New York, NY Willbraham, A. Staley, D. Matta, M. 1995 Chemistry. Addison-Wesley: Menlo Park, CA 95 Experiment 14 Comparison of the geometries and shapes of molecules and ions Purpose: To make models of various molecules and ions To study the geometry and shape of these models Background: Valence shell electrons Valence electrons are the electrons in the outermost energy level. “The number of valence electrons largely determines the chemical properties of an element” (pg. 374 Willbraham). The periodic table is arranged such that elements with the same number of valence electrons are arranged in vertical columns. In the group A elements the number of valence electrons equals the group number thus all elements in group 1A all have 1 valence electron. The elements in group 2A all have 2 valence electrons, etc. As a note of interest all transition metals in the B group have 2 valence electrons. That is to say all transition metals have 2 valence electrons. Electron dot diagram The Lewis structure of a single atom, often called a “dot diagram,” consists of a chemical symbol for an element and a dot for each valence electron. When assigning electrons in an electron dot configuration, use the following conventions: 1. There are four orbitals represented as four sides surrounding the atom. 2. Place one electron into each side (orbital). 3. After all sides have one electron you may, if necessary, place a second electron in an orbital. 1 valence electron 2 valence electrons 3 valence electrons 4 valence electrons 5 valence electrons 6 valence electrons 7 valence electrons 8 valence electrons When there is only one electron on a side (orbital) the electron is called a free radical. If there are two electrons on a side it is called a lone pair. Thus: when there are two valence electrons there are two free radicals. When there are seven valence electrons there are three lone pairs and one free radical. Lewis structures When attempting to draw Lewis structures please use the following easy steps: 1. Find the sum of all valence electrons in all atoms of the molecule 2. Determine the central atom and use a pair of electrons to form a bond between each atom bound to the central atom. 3. Add remaining electrons in pairs starting with the highest electronegativity (right and up on the periodic table) to satisfy the octet rule (except, 2 e- for H and 6 e- B). 4. Rearrange electrons to from double and triple bonds to satisfy octet rule (except for H and boron). There are exceptions to the rule that all elements needed 8 electrons, for instance hydrogen only needs 2, boron only needs 6 and phosphorus, sulfur and chlorine can have more than 8 electrons, because of there large atomic radius and the presence of the “d” subshell. 96 Bond polarity in covalent bonds In each bond between two nonmetals the electrons are being shared. This shared pair of electrons is called a covalent bond. There are three types of covalent bonds. • • • Covalent: electrons are shared equally Polar covalent: electrons are shared but one atom has more contact than the other atom. This makes one atom slightly positive and the other atom slightly negative. Ionic in nature: one atom in the bond is in so much more contact with the electron. That bond takes on an ionic nature, where one atom maintains a nearly full negative charge and the other atom maintains a nearly full positive charge. Each bond in a molecule has a shared pair of electrons. However some atoms do not share equally. Electronegativity is a measure of how equally the atoms in a bond share the electrons. You will find the electronegativity value for each atom on the periodic table on the back of this lab manual. It is in italics in the bottom right corner. When evaluating the polarity of a bond find the difference between the atoms’ electronegativities. For instance in an O-H bond, hydrogen has an electronegativity of 2.1and oxygen has an electronegativity of 3.5. The difference is found by subtracting the two values and taking the absolute value. Thus in an O-H bond the difference is 1.4. After calculating the difference in electronegativity consult the table below to determine the type of bond that exists between two atoms. Covalent Polar covalent Ionic in nature 0.0 0.4 2.0 ------------------------------------|---------------------------------------------|----------------------------------From 0.0 up to and Above 0.4 and less than 2.0 From 2.0 and greater including 0.4 The above information and table only applies to nonmetal-nonmetal bonds. When considering combinations of metals and nonmetals the bonds are referred to as ionic. Geometry The Lewis structure represents the molecule in two dimensions. The actual arrangement of the atoms in space occurs in three dimensions. The actual arrangement of atoms within the molecule can be approximated by considering the arrangements which maximizes the distance between regions of the negative charge in the central atom. These regions of negative charge are attributes to the outer, or valence, shell electrons. After determining the number of electron pairs around an atom, Valence Shell Electron Pair Repulsion Theory (VSEPR) is used to predict the geometry. The VSEPR theory states that the geometric arrangement of atoms or groups of atoms, around a central atom, is determined solely by the repulsion between electron pairs present in the valence shell of the central atom. This means that the things attached to the central atom will try to stay as far away from each other as possible and yet stay connected to the central atom. They attempt to stay away from each other because each electron pair is negatively charged and is repelled by the negative charge on the other electron pairs. Atoms and lone pairs connected to the central atom are considered electron pairs. There are only five possible geometries: 97 # of electron groups (domains) Geometry 2 Linear 3 Trigonal planar 4 Tetrahedral 5 Trigonal bipyramidal 6 Octahedral Using water as an example, there are four lone pairs of electrons arranged around the central oxygen and these four pairs of electrons repel each other. The resulting arrangement allows the four pairs to be as far away from each other as possible. This particular arrangement is called tetrahedral. In general, the geometry can be predicted from the number of electron pairs around the central atom. However the possibility of double and triple bonds must be accounted for. Rather than count the pairs of electrons around the central atom, the number of groups of electrons (domains) around the central atom are counted. Thus a double bond with two pairs of electrons contributing to the bond counts as just one group of electrons. The three pairs of electrons in a triple bond count as only one group. Note: Each lone pair attached to the central atom counts as one group (or domain). Each atom attached to the central atom counts as one group (or domain). 98 VSEPR Class Once the Lewis structure has been determined, it is easy to define the VSEPR class. Use an A to represent the central atom. Represent each atom surrounding the central atom as an X. Then represent each lone pair of electrons surrounding the central atom as an E. Thus a molecule which has a central atom with four atoms around it and no lone pairs would be AX4. A molecule with a central atom that is connected to two atoms and has two lone pairs would be AX2E2. Geometry Linear AXE type Trigonal planar AXE type Tetrahedral AXE type 0 lone pairs 1 lone pair 2 lone pairs 3 lone pairs AX2 AXE AX3 AX2E AXE2 AX4 AX3E AX2E2 AXE3 AX5 AX4E AX3E2 AX2E3 AX6 AX5E AX4E2 Trigonal Bipyramidal AXE type Octahedral AXE type Example: H2O has a tetrahedral geometry but the shape is described as bent Lewis structure Geometry Tetrahedral or Shape After the geometric arrangement of the electron groups has been determined, it is possible to describe the shape. The shape describes the three dimensional arrangement of the atoms in a molecule. If the molecule has no lone pairs, the shape is the same as the geometry. For molecules with lone pairs, the unshared pairs are used to determine the geometry but are ignored when describing the shape. The following table illustrates the different geometries and subsequently possible shapes. 99 Shape Geometry No lone pairs 1 lone pair 2 lone pair 3 lone pair 2 electron groups (domains) Linear Linear Linear Trigonal Planar Bent Linear Tetrahedral Trigonal Pyramidal Bent Linear Trigonal Bipyramidal Seesaw T-shaped Linear Octahedral Square pyramidal Square planar 3 electron groups (domains) Trigonal Planar 4 electron groups (domains) Tetrahedral 5 electron groups (domains) Trigonal Bipyramidal 6 electron groups (domains) Octahedral Example: H2O has a tetrahedral geometry but because it has two lone pairs the shape is described as bent. It’s AXE type is AX2E2 Lewis structure Geometry Shape or Tetrahedral Consider all electron domains 100 Bent Considers the orientation of only the bound atoms Molecular Polarity Once the shape of a molecule or ion has been determined, it is possible to predict whether the molecule or ion is polar. A covalent bond in a molecule is polar when two atoms of different electronegativity share the pair of electrons in the bond. The atom with greater electronegativity will be partially negative, while the other atom will be partially positive. A molecule or ion with a polar bond will be polar unless there is a bond (or sum of bonds) which is equally polar in the opposite direction. Any asymmetric molecule will be polar if it has polar bonds and any symmetric molecule will be nonpolar whether or not it has polar bond. Thus: • • In nearly all cases lone pairs on a molecule will cause that molecule to be polar. In nearly all cases unequal distribution of polar bonds will cause the molecule to be polar If a molecule is polar, the positive and negative areas of the molecule should be shown. There are two methods of showing this described below. Example: Water is a polar molecule The oxygen side of the molecule has a slight negative charge because when it shares the electrons with hydrogen in the bond it keeps them in the oxygen orbitals for longer periods of time. This makes the oxygen side have a negative charge. Because the electrons are pulled more frequently to the oxygen side it leaves the hydrogen side with a low electron density which makes the hydrogen side slightly positive. Bond angle The bond angles are an important part of the description of the geometry and shape of a molecule in three dimensions. It is possible to predict the approximate bond angle from the geometric arrangement of the atoms in a molecule or ion. Geometry Linear Trigonal Planar Tetrahedral Trigonal bipyramidal Octahedral Approximate bond angle 180o 120o 109.5o 90o, 120o, 180o 90 o 101 Hybridization of the Central Atom In atoms, the areas where the electrons are likely to be found are called orbitals. In molecules, these atomic orbitals become hybridized so that the bonding electrons can be located between the atoms. The hybridization about the central atom can be predicted if the number of electron domains is counted. The number of electron domains is always the same as the number of orbitals that make up the hybridization. Geometry # of electron groups (domains) Hybridization of central atom Linear 2 sp Trigonal Planar 3 sp2 Tetrahedral 4 sp3 Trigonal bipyramidal 5 sp3d Octahedral 6 sp3d2 In a “sp” hybridized orbital one orbit from the “s” subshell and 1 orbit from the p subshell are combined to form the sp hybridized orbital. There are a total of 2 orbits in a sp hybrid. Please note that there are 3 p orbitals. In the sp hybrid 2 p orbitals have not been used In a “sp2” hybridized orbital one orbit from the “s” subshell and 2 orbits from the p subshell are combined to form the sp2 hybridized orbital. There are a total of 3 orbitals in a sp2 hybrid. Please note that there are 3 p orbitals. In the sp2 hybrid 1 p orbitals have not been used In a “sp3” hybridized orbital one orbit from the “s” subshell and 3 orbits from the p subshell are combined to form the sp3 hybridized orbital. There are a total of 4 orbitals in a sp3 hybrid. In a “sp3d” hybridized orbital one orbit from the “s” subshell, 3 orbits from the p subshell and one orbit from the d subshell are combined to form the sp3d hybridized orbital. There are a total of 5 orbitals in a sp3d hybrid. In a “sp3d2” hybridized orbital one orbit from the “s” subshell, 3 orbits from the p subshell and two orbit from the d subshell are combined to form the sp3d2 hybridized orbital. There are a total of 6 orbitals in a sp3d2 hybrid. Pre-Lab Questions: 1. When drawing Lewis structures it is important to know the octet rule. a. State the octet rule. b. Which atoms are allowed to take less than eight electrons when drawing Lewis structures? c. Which atoms are allowed to obtain more than eight electrons when drawing Lewis structures? 102 Procedure: Atoms Look at your atom kit and determine what model atom will be used for each of the following Boron Hydrogen Phosphorus Fluorine Nitrogen Sulfur Carbon Oxygen Chlorine Model kit electron pairs Parts Short straight Paddle Long Curved Electron pairs 2 electrons in a single bond 2 electrons as a pair, but not involved in a bond (lone pair) 2 electrons as part of a double or triple bond For each of the molecules below, draw the Lewis structure in your notebook, build the model, then answer questions A-J. Finally, do all the comparisons. Model 1 BF3 (boron trifluoride) Model 7 SF6 (sulfur hexafluoride) Model 2 CH4 (methane) Model 8 CO2 (carbon dioxide) 2Model 3 NH3 (ammonia) Model 9 CO3 (carbonate ion) Model 4 H2 O (water) Model 10 CH3Cl (chloromethane) Model 5 NH4+ (ammonium) Model 11 BeCl2 (Beryllium chloride) Model 6 PCl5 (phosphorus pentachloride) Model 12 ICl5 (Iodine pentachloride) Questions Identify the central atom and answer each of the following questions for each model. You must completely copy the question into your laboratory notebook before working on the first model. You may find it easiest to make a large table with headings to do this exercise. A. How many electron pairs (domains) are present F. What is the shape of the molecule? around the central atom? G. Are there any polar bonds? B. How many of the e- pairs are atoms bonded to the H. Is the molecule polar or nonpolar? central atom (count double bonds as one bond)? C. How many electron pairs connected to the central I. List all bond angles 180o or less. atom are lone pairs? D. What is the VSEPR class (AXmEn)? J. Hybridization around the central atom? E. What is the geometry of the molecule? K. Make a 3-D sketch of the model. When making the 3-D sketch please use the following conventions as an example. Let us consider the chloroform molecule CH2Cl2: A standard 3-D sketch These two lines and attached H atoms are on the same plane as the central atom This dashed wedge indicates a bond and Cl atom that is going away from you. 103 This wedge indicates a bond and Cl atom that is coming toward you In general keep double bonds on the same plane as the central atom. Data Table: Please construct your data table to look like this, (landscape oriented). It will make organizing and comparing your data easier and will simplify grading. If it is not like this example it may not be accepted. Model Example: H3 O+ Model #1 BF3 Model Model #2 CH4 Model #3 NH3 Lewis structure Lewis structure A. B. 1 C. C. 3 B. 4 A. D. E. Tetrahedral E. F. Trigonal Pyramidal F. H+àO or δ+H – Oδ- G. G. D. AX3E H. Polar H. I. I. J. Sp3 J. K. K. 104 H-O 109.5 Results: Do all of the following comparisons. Be sure to consider your answers to the questions on the previous page. Always use the formula or name of the compound, not the model number when making comparisons. Write in complete sentences to answer each question. 1. Compare models for CH4, NH3 and H2O. Are there similarities in the number of electron pairs, the geometry, the shape, hybridization, etc? Differences? Are any of the molecules polar? 2. Compare CH4 and NH4+. Are there similarities in the number of electron pairs, the geometry, the shape, hybridization, etc? Differences? Do the sketches look similar? Is either molecule polar? 3. Compare the models for NH3 and NH4+. Are there similarities in the number of electron pairs, the geometry, the shape, hybridization, etc? Differences? DO the sketches look similar? Is either molecule polar? 4. Compare models for PCl5 and SF6. What do they have in common? Is either molecule polar? 5. Compare CH4 and CO2. Are there similarities in the number of electron pairs, the geometry, the shape, hybridization, etc? Differences? What do they have in common? Is either molecule polar? 6. Compare H2O and CO2. Are there similarities in the number of electron pairs, the geometry, the shape, hybridization, etc? Differences? Is either molecule polar? 7. Compare models for BF3 and CO32-. Are there similarities in the number of electron pairs, the geometry, the shape, hybridization, etc? Differences? Is either molecule polar? 8. Compare the models for NH3 and CO32-. Are there similarities in the number of electron pairs, the geometry, the shape, hybridization, etc? Differences? Is either molecule polar? 9. Compare CH4 and CH3Cl. Are there similarities in the number of electron pairs, the geometry, the shape, hybridization, etc? Differences? Is either molecule polar? Discussion: What information is shown by a model but is not shown in a Lewis structure? Does a Lewis structure show whether a molecule is polar? Explain your answer using example from the models. Post-Lab Questions: create a data table like the one used previously in this laboratory activity to record the following answers. 1. 2. 3. 4. Predict the answers to questions A-K for a model of CCl4. Predict the answers to questions A-K for a model of CHCl3 Predict the answers to questions A-K for a model of CH2Cl2 Compare CH4, CH3Cl, CH2Cl2, CHCl3, CCl4 References: Finnegan, M. Place, H. Weissbart, B. (2000) Washington state university chemistry 101-102 laboratory manual. Star Publishing Company: Belmont, California Willbraham, A. Staley, D. Matta, M. (1995) Chemistry 4th edition. Addison-Wesley: Menlo Park, CA Silberberg, Martin. (1996) Chemistry the molecular nature of matter and change. Mosby: New York, NY 105 Experiment 15 Vapor Pressure of Liquids In this experiment, you will investigate the relationship between the vapor pressure of a liquid and its temperature. When a liquid is added to the Erlenmeyer flask shown in Figure 1, it will evaporate into the air above it in the flask. Eventually, equilibrium is reached between the rate of evaporation and the rate of condensation. At this point, the vapor pressure of the liquid is equal to the partial pressure of its vapor in the flask. Pressure and temperature data will be collected using a Gas Pressure Sensor and a Temperature Probe. The flask will be placed in water baths of different temperatures to determine the effect of temperature on vapor pressure. You will also compare the vapor pressure of two different liquids, ethanol and methanol, at the same temperature. Objectives: In this experiment, you will • • Investigate the relationship between the vapor pressure of a liquid and its temperature. Compare the vapor pressure of two different liquids at the same temperature. Figure 1 Materials: computer Vernier computer interface Logger Pro Vernier Gas Pressure Sensor Vernier Temperature Probe rubber-stopper assembly plastic tubing with two connectors 20 mL syringe two 125 mL Erlenmeyer flasks methanol ethanol ice four 1 liter beakers 106 Procedure: 1. Obtain and wear goggles! CAUTION: The alcohols used in this experiment are flammable and poisonous. Avoid inhaling their vapors. Avoid contacting them with your skin or clothing. Be sure there are no open flames in the lab during this experiment. Notify your teacher immediately if an accident occurs. 2. Use 1 liter beakers to prepare four water baths, one in each of the following temperature ranges: 0 to 5°C, 10 to 15°C, 20 to 25°C (use room temperature water), and 30 to 35°C. For each water bath, mix varying amounts of warm water, cool water, and ice to obtain a volume of 800 mL in a 1 L beaker. To save time and beakers, several lab groups can use the same set of water baths. 3. Prepare the Temperature Probe and Gas Pressure Sensor for data collection. a. Plug the Gas Pressure Sensor into CH1 and the Temperature Probe into CH2 of the computer interface. b. Obtain a rubber-stopper assembly with a piece of heavy-wall plastic tubing connected to one of its two valves. Attach the connector at the free end of the plastic tubing to the open stem of the Gas Pressure Sensor with a clockwise turn. Leave its two-way valve on the rubber stopper open (lined up with the valve stem as shown in Figure 2) until Step 9. c. Insert the rubber-stopper assembly into a 125 mL Erlenmeyer flask. Important: Twist the stopper into the neck of the flask to ensure a tight fit. Figure 2 Figure 3 4. Follow your instructions on how to calibrate the sensors. 5. The temperature and pressure readings should now be displayed in the meter. While the twoway valve above the rubber stopper is still open, record the value for atmospheric pressure in your data table (round to the nearest 0.1 kPa). 6. Finish setting up the apparatus shown in Figure 3: a. Obtain a room-temperature water bath (20–25°C). b. Place the Temperature Probe in the water bath. c. Hold the flask in the water bath, with the entire flask covered as shown in Figure 3. d. After 30 seconds, close the 2-way valve above the rubber stopper as shown in Figure 4—do this by turning the white valve handle so it is perpendicular with the valve stem itself. open closed Figure 4 7. Obtain the methanol container and the syringe. Draw 3 mL of the methanol up into the syringe. With the two-way valve still closed, screw the syringe onto the two-way valve, as shown in Figure 3. 107 8. Introduce the methanol into the Erlenmeyer flask. a. Open the 2-way valve above the rubber stopper—do this by turning the white valve handle so it is aligned with the valve stem (see Figure 4). b. Squirt the methanol into the flask by pushing in the plunger of the syringe. c. Quickly return the plunger of the syringe back to the 3 mL mark of the syringe, then close the 2-way valve by turning the white valve handle so it is perpendicular with the valve stem. d. Remove the syringe from the 2-way valve with a counter-clockwise turn. 9. To monitor and collect pressure and temperature data: a. Click . b. When the pressure and temperature readings displayed in the meter stabilize, equilibrium between methanol liquid and vapor has been established. Click . The first pressuretemperature data pair is now stored. 10. To collect another data pair using the 30–35°C water bath: a. Place the Erlenmeyer flask assembly and the temperature probe into the 30–35°C water bath. Make sure the entire flask is covered. b. When the pressure and temperature readings displayed on the computer monitor stabilize, click . The second data pair has now been stored. 11. For Trial 3, repeat the Step-10 procedure, using the 10–15°C water bath. Then repeat the Step-10 procedure for Trial 4, using the 0–5°C water bath. 12. Click to end data collection. Record the pressure and temperature values in your data table, or, if directed by your instructor, print a copy of the table. 13. Gently loosen and remove the Gas Pressure Sensor so the Erlenmeyer flask is open to the atmosphere. Remove the stopper assembly from the flask and dispose of the methanol as directed by your teacher. 14. Obtain another clean, dry 125 mL Erlenmeyer flask. Draw air in and out of the syringe enough times that you are certain that all of the methanol has evaporated from it. 16. Repeat Steps 6–8 to do one trial only using ethanol in the room temperature water bath. When the pressure stabilizes, record the measured pressure of ethanol displayed in the meter in your data table. 17. Gently loosen and remove the stopper assembly from the flask and dispose of the ethanol as directed by your teacher. 108 Processing the Data: 1. Convert each of the Celsius temperatures to Kelvin (K). Write the answer in the space provided. 2. To obtain the vapor pressure of methanol and ethanol, the air pressure must be subtracted from each of the measured pressure values. However, for Trials 2–4, even if no methanol was present, the pressure in the flask would have increased due to a higher temperature, or decreased due to a lower temperature (remember those gas laws?). Therefore, you must convert the atmospheric pressure at the temperature of the first water bath to a corrected air pressure at the temperature of the water bath in Trial 2, 3, or 4. To do this, use the gas-law equation (use the Kelvin temperatures): P2 P1 = T2 T1 where P1 and T1 are the atmospheric pressure and the temperature of the Trial 1 (room temperature) water bath. T2 is the temperature of the water bath in Trial 2, 3, or 4. Solve for P2, and record this value as the corrected air pressure for Trials 2, 3, and 4. For Trial 1 of methanol and Trial 1 of ethanol, it is not necessary to make a correction; for these two trials, simply record the atmospheric pressure value in the blank designated for air pressure. 3. Obtain the vapor pressure by subtracting the corrected air pressure from the measured pressure in Trials 2-4. Subtract the uncorrected air pressure in Trial 1 of methanol (and Trial 1 of ethanol) from the measured pressure. 4. Plot a graph of vapor pressure vs. temperature (°C) for the four data pairs you collected for methanol. Temperature is the independent variable and vapor pressure is the dependent variable. As directed by your teacher, plot the graph manually, or use Logger Pro. Note: Be sure to plot the vapor pressure, not the measured pressure. 5. How would you describe the relationship between vapor pressure and temperature, as represented in the graph you made in the previous step? Explain this relationship using the concept of kinetic energy of molecules. 6. Which liquid, methanol or ethanol, had the larger vapor pressure value at room temperature? Explain your answer. Take into account various intermolecular forces in these two liquids. 109 Data and Calculations: Atmospheric pressure _______ kPa Substance Trial Methanol 1 2 Ethanol 3 4 1 Temperature (°C) °C °C °C °C °C Temperature (K) K K K K K kPa kPa kPa kPa kPa Measured pressure Air pressure no correction corrected corrected corrected no correction kPa kPa kPa kPa kPa kPa kPa kPa kPa kPa Vapor pressure Results: Draw a graph that reports your data. Discussion: What does the graph and the differences in vapor pressure tell you about the strength of intermolecular forces of the different chemicals? What might account for these differences? How does temperature effect vapor pressure and why? Propose a possible follow-up experiment. Post Lab Questions: 1. What kind are the molecular weights for each chemical tested in this lab. 2. Draw Lewis structures for each chemical tested in this lab. a. Indicate all polar bonds in each Lewis structure. b. Indicate molecular polarity 3. What strongest type of intermolecular forces are present in each of the chemicals tested in this lab 4. Provide an explanation based on molecular polarity and molecule weight to account for the differences in vapor pressure. 110 Experiment 15 Calorimetry: Heat of Fusion of Ice Purpose: To determine the energy (in Joules) required to melt one gram of ice and to determine the molar heat of fusion for ice (in kJ/mol). Background: When a chemical or physical change takes place heat is either given off or absorbed. That is, the change is either exothermic or endothermic. It is important for chemists to be able to measure this heat. Measurements of this kind are made in a device called a calorimeter (kal rim’əә təәr). The technique used in making these measurements is called calorimetry. In simplest terms, a calorimeter is an insulated container made up of two chambers. The outer chamber contains a known mass of water. In the inner chamber, the experimenter places the materials that are to lose or gain heat while undergoing a physical or chemical change. The basic principle on which the calorimeter works is that when two bodies at different temperature are in contact with one another, heat will flow from the warmer body to the colder body. Thus, heat lost by one body will be gained by the other. This exchange of heat will continue until the two bodies are at the same temperature. In a calorimeter heat is exchanged between the water and the materials undergoing change. Unlike most calorimeters, the simple Styrofoam-cup calorimeter used in this experiment will have only one chamber. The ice will be placed directly into a measured amount of water. The heat required to melt the ice will be supplied by the water. By measuring the temperature change (ΔT) of the water, you can calculate the quantity of heat exchanged between the water and the ice. Using these experimental data, you will calculate the heat of fusion of ice. Melting and freezing behavior are among the characteristic properties that give a pure substance its unique identity. As energy is added, pure solid water (ice) at 0°C changes to liquid water at 0°C. In this experiment, you will determine the energy (in joules) required to melt one gram of ice. You will then determine the molar heat of fusion for ice (in kJ/mol). Excess ice will be added to warm water, at a known temperature, in a Styrofoam cup. The warm water will be cooled down to a temperature near 0°C by the ice. The energy required to melt the ice is removed from the warm water as it cools. To calculate the heat that flows from the water, you can use the relationship q = m• Cp •Δt where q stands for heat flow, Cp is specific heat, m is mass in grams, and Δt is the change in temperature. For water, Cp is 4.18 J/g°C. Materials: 100 mL graduated cylinder Vernier computer interface warm water Temperature Probe Styrofoam cup ring stand Logger Pro tongs 111 stirring rod ice cubes utility clamp 250 mL beaker Procedure: 1. 2. 3. 4. Connect the probe to the computer interface. Place a Styrofoam cup into a 250 mL beaker as shown in Figure 1. Use a utility clamp to suspend the Temperature Probe on a ring stand as shown in Figure 1. Use a 100 mL graduated cylinder to obtain 100.0 mL of water at about 60°C from your instructor. Record this as V1. 5. Obtain 7 or 8 large ice cubes. 6. Lower the Temperature Probe into the warm water (to about 1 cm from the bottom). 7. Click to begin data collection. Wait until the temperature reaches a maximum (it will take a few seconds for the cold probe to reach the temperature of the warm water). This maximum will determine the initial temperature, t1, of the water. As soon as this maximum temperature is reached, fill the Styrofoam cup with ice cubes. Shake excess water from the ice cubes before adding them (or dry with a paper towel). Record the maximum temperature, t1, in your data table. Figure 1 8. Use a stirring rod to stir the mixture as the temperature approaches 0°C. Important: As the ice melts, add more large ice cubes to keep the mixture full of ice! 9. When the temperature reaches about 4°C, quickly remove the unmelted ice (using tongs). Continue stirring until the temperature reaches a minimum (and begins to rise). This minimum temperature is the final temperature, t2, of the water. Record t2 in your data table. Click when you have finished collecting data. 10. Use the 100 mL graduated cylinder to measure the volume of water remaining in the Styrofoam cup to the nearest 0.1 mL. Record this as V2. 11. You can confirm your data by clicking the Statistics button, . The minimum temperature (t2) and maximum temperature (t1) are listed in the floating box on the graph. Calculations: 1. 2. 3. 4. Use the equation Δt = t2 – t1 to determine Δt, the change in water temperature. Subtract to determine the volume of ice that was melted (V2 –V1). Find the mass of ice melted using the volume of melt (use 1.00 g/mL as the density of water). Use the equation given in the introduction of this experiment to calculate the energy (in joules) released by the 100 g of liquid water as it cooled through Δt. 5. Now use the results obtained above to determine the heat of fusion—the energy required to melt one gram of ice (in J/g H2O). 6. Use your answer to Step 5 and the molar mass of water to calculate the molar heat of fusion for ice (in kJ/mol H2O). 7. Find the percent error for the molar heat of fusion value in Step 6. The accepted value for molar heat of fusion is 6.01 kJ/mol. 112 Data and Calculations: Initial water temperature, t1 °C Final water temperature, t2 °C Change in water temperature, Dt °C Final water volume, V2 mL Initial water volume, V1 mL Volume of melt mL Results: Mass of ice melted Heat released by cooling water (q = Cp•m•Δt) g J/g ice melted (heat of fusion) J kJ/mol ice melted (molar heat of fusion) J/g Percent error kJ/mol Discussion: Please discuss issues suggested in “how to write a lab report.” % Post-Lab Questions: 1. List possible sources of error. Describe the design of a calorimetry experiment that would reduce some of the error potential. 2. One source of error is the flow of heat between the water in the cup and the surroundings. Explain how this error is reduced by starting with water at 50 oC. 3. In what way does the calorimetry make use of the law of conservation of energy? 4. Define the following terms: a. exothermic, b. endothermic, c. heat of fusion, d. specific heat capacity. 5. Is the process of melting exothermic or endothermic? Give evidence to support this answer. 6. What is the difference between heat and temperature? 113 Experiment 17 Boyle’s Law: Pressure-Volume Relationship in Gases The primary objective of this experiment is to determine the relationship between the pressure and volume of a confined gas. The gas we use will be air, and it will be confined in a syringe connected to a Gas Pressure Sensor (see Figure 1). When the volume of the syringe is changed by moving the piston, a change occurs in the pressure exerted by the confined gas. This pressure change will be monitored using a Gas Pressure Sensor. It is assumed that temperature will be constant throughout the experiment. Pressure and volume data pairs will be collected during this experiment and then analyzed. From the data and graph, you should be able to determine what kind of mathematical relationship exists between the pressure and volume of the confined gas. Historically, this relationship was first established by Robert Boyle in 1662 and has since been known as Boyle’s Law. Please note that cm3 is often written “cc” in the medical field. Also, keep in mind that cm3 is the same as mL. You may find it also useful to know that 101.325 kPa equals 1 atm (this relationship has infinite sig figs). Objectives: In this experiment, you will Use a Gas Pressure Sensor and a gas syringe to measure the pressure of an air sample at several different volumes. • Determine the relationship between pressure and volume of the gas. • Describe the relationship between gas pressure and volume in a mathematical equation. • Use the results to predict the pressure at other volumes. • Figure 1 Materials: Lab Quest 1. 2. 3. 4. 5. Vernier Gas Pressure Sensor 20 mL gas syringe Obtain Lab Quest, Gas Pressure Sensor and a 20mL syringe. Connect Gas Pressure Sensor to the Lab Quest into port 1. Turn Lab Quest on. Make sure Lab Quest is reading pressure in kPa. If it is not in kPa please change units to kPa. The pressure in kPa should be between around 90 and 120 kPa. If it is not, please ask your instructor for help. Please do not forget to record both your observations and the quantitative data. 114 6. Obtain a syringe. DO NOT CONNECT YOUR SYRINGE to the gas pressure sensor. Please notice the front of the black rubber portion of the plunger is the side that measurements are taken from (see figure 2) (figure 2) This is a measurement of 5.0 mL 7. With your syringe NOT CONNECTED to the gas pressure sensor move the plunger to the 20.0 mL mark. 8. With the syringe at 20.0 mL now CONNECT the syringe to the gas pressure sensor. 9. Record the gas pressure in kPa. Now record the volume. Please add 0.8ml to the volume to account for the volume inside the gas pressure sensor. 10. Do not disconnect the syringe from the gas pressure sensor. Move the plunger to 15.0 mL and hold it long enough to record the pressure in kPa. One person should move the plunger and one person should record the gas pressure. Record the volume. Remember to add 0.8mL to all volume data readings. Move the plunger back to 20.0 mL. Do not remove syringe from Gas Pressure Sensor. 11. Repeat step 10. However in this step move plunger to 12.0 mL. 12. Repeat step 10. However in this step move plunger to 10.0 mL. 13. Repeat step 10. However in this step move plunger to 8.0 mL. 14. Graph your data on graph paper. Place the volume on the X axis and the pressure in atm on the Y axis. Please use all graphing conventions. Data: Volume (mL) add 0.8mL Pressure (kPa) Calculations: Convert all kPa values into atm, using dimensional analysis. Calculate P/V for all samples. This may produce a constant value for K Calculate P*V for all samples. This may produce a constant value for K Report values in a table such as Pressure in Atm P/V P*V 115 The correct mathematical relationship between gas volume and pressure will produce a constant value. This value will be labeled as k. Results: Provide a graph of the data. Draw a best fit smoothed line. Do not connect the dots with straight lines. Report the Value found for k. Discussion: Cover topics described in how to write a lab report. Propose an experiment that would demonstrate Charles’s Law. Post lab questions: 1. Using the same the best fit line on your graph. What does your data show would happen to the pressure if the volume is changed from 20.0 mL to 13.5 mL? Show the pressure values in your answer. 2. From your data and the shape of the curve in the plot of pressure vs. volume, do you think the relationship between the pressure and volume of a confined gas is direct or inverse? Explain your answer. 3. Based on your data, what would you expect the pressure to be if the volume of the syringe was increased to 40.0 mL? Explain or show work to support your answer. 4. Based on your data, what would you expect the pressure to be if the volume of the syringe was decreased to 2.5 mL? Explain or show work to support your answer. 5. What experimental factors are assumed to be constant in this experiment? 6. How constant were the values for k you obtained in Question 8? Good data may show some minor variation, but the values for k should be relatively constant. 7. Using P, V, and k, write an equation representing Boyle’s law. Write a verbal statement that correctly expresses Boyle’s law. 116 Experiment 18 Quantitative Reaction of HCl and Mg Purpose: To calculate the molar volume of a gas collected over water Hypothesis: Develop a hypothesis stating how many liters of gas will be produced for each mole of gas produced. Background: Pertinent Gas Laws To understand the gas laws at an intuitive level it is valuable to understand the kinetic molecular model of gas behavior. The kinetic molecular model it makes three basic assumptions: 1. Gas is composed of particles and there is no attractive force between gas particles. 2. The particles of gas move rapidly in straight lines and are in constant random motion. 3. All collisions are perfectly elastic, meaning that kinetic energy is transferred from one particle to another in a collision but no energy is lost. Temperature in this model is a measure of how much kinetic energy the speeding particles have. The equation for kinetic energy is Ekinetic = ½ mass * Velocity2. From the equation you can see that as temperature goes up particles move faster. While working with gas laws, temperature is always measured in Kelvin. To convert from Celsius to Kelvin add 273.14 to convert from Kelvin to Celsius subtract 273.14. Pressure in this model is how much force is pressed against the container when the speeding gas particles collide with the surfaces of the container. Pressure is measured in many different units. The three units commonly used in this class are millimeters mercury (mmHg), and kilopascal (kPa), atmospheres (atm) 760 mmHg = 101.3kPa = 1atm Boyle’s Law Consider the effect of pressure on the volume of a contained gas while temperature is held constant. As volume decreases the pressure in the container goes up. As volume increases pressure goes down. Boyle’s Law states that for a given mass of gases at constant temperature, the volume of gas varies inversely with the pressure. Mathematically it is expressed as: P1V1=P2V2 In this equation P1and V1 represent the temperature and pressure at the initial set of conditions and P2 and V2 are the final set of conditions. Thus if under one set of conditions the pressure is 100 atm and the volume is one liter the same mass of gas at the same pressure would have a pressure of 50 if it was contained in a 2 liter container. P1V1=P2V2 100 atm* 1L = 50 atm *2 L 117 By applying the kinetic model to the behavior of gas the pressure increases in a smaller container (if temperature is held constant) because as the volume decreases the surface area decreases a lot. With a smaller surface area each small area of surface is getting hit more often. Imagine 5 pool balls constantly moving on a pool table. If the pool table is huge, say the size of a soccer field, the balls infrequently hit the sides. If you keep the balls moving at the same speed but shrink the table down to the size of your desk it is easy to see that the balls now hit the sides of the table more frequently. Remember pressure is a measure of how often and how hard the particles hit the surfaces of the container. Ideal Gas Law The Ideal Gas Law allows you to consider the effect of changing the number of molecules present in a sample. There are different ways to increase pressure, for instance you could speed the molecules up so they hit the surface of the container harder (increase temperature), or you can decrease the surface area of the container so that the particles hit more often (decrease volume) or you can do one other thing to increase pressure. You can put more particles in the container so that the surfaces of the container are hit more often because there is more stuff inside to hit the walls. Using the pool table analogy to change pressure you can change the size of the table. Rolls the balls faster or just put more balls on the table. The Ideal Gas Law states that in any gas system if you multiply the pressure times the volume and divide that by the product of the temperature and the moles it will equal a constant value. This is true for any ideal gas. Thus P*V Where P = pressure (atm) V= Volume (L) =R T*n T= temperature (K) n= moles of gas Thus P1 * V1 = R T1 * n1 R= 0.08206 or expressed in common form: PV=nRT atm*L mol*K This is an amazingly powerful equation. It is the superman of gas law equations. It can pretty much do anything. Dalton’s Partial Pressure Law Consider the pool table. Pressure is measured by how often and how hard the balls hit the side of the container. It is easy to picture the balls rolling around in straight lines and colliding with the bumpers and bouncing back to keep on moving. Dalton’s Partial Pressure Law states that each gas in a container contributes to the total pressure. So if on the pool table there are blue balls and some red balls and some green balls the total pressure exerted on the sides of the pool table is equal to the sum of the pressure exerted from each of the different balls. Thus Ptotal =P1 +P2 +P3 etc The contribution each gas in a mixture makes to the total pressure is the partial pressure exerted by that gas. Dalton’s Law of Partial Pressures: at constant volume and temperature, the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the component gases. 118 So if on the pool table you want to know how much pressure is exerted you need to know how much pressure is being exerted by all of the different gases in the container. So in your class room, let us assume that the pressure is 1.00 atm. Let us also assume that the air in the classroom is 78% nitrogen, 21% oxygen and 1% carbon dioxide. In this situation the total pressure in the room is 0.78 atm nitrogen + 0.21 atm oxygen + 0.01 atm carbon dioxide = 1 atm air Calculations The most challenging part of this lab will be to complete the calculations. A. To find grams of Magnesium ribbon 1. Convert cm of Mg to grams of Mg Let us assume that 1 meter of Mg ribbon has a mass of 0.750 grams. If you have 5.03 cm of magnesium ribbon you can calculate how many grams of Mg as follows: 5.03 cm Mg 1m 100 cm 0.750 grams Mg 1 meter Mg Sample equation 1 = .0377 g Mg 2. Convert grams of Mg to Moles Mg Use dimensional analysis and the atomic weight of Mg found on the periodic table. 0.0377 g Mg 1 mol Mg Sample equation 2 24.30 g Mg = 0.00155 mol Mg (Result A) B. To find theoretical moles of hydrogen gas produced 0.00155 mol Mg 1 mol H2 1 mol Mg = 0.00155 mol H2 Sample equation 3 (Result B) C. Find experimental ideal gas constant. The ideal gas law is PV=nRT. To solve for R rearrange the equation into this form R=PV/nT 1. To solve for n, convert mole of Mg to moles of hydrogen gas 2HCl(aq) + Mg(s) à H2(g) + MgCl2(aq) 1. Find T T = the temperature in Kelvin to convert Celsius to Kelvin add 273.14 to the Celsius measurement 273.14 + 19.0 oC = 292.2 K 3. Find Volume Read V off of your gas collection tube to 4 sig figs. In this example the value is 37.85mL The units of this value must be liters so convert mL to liters. 4. Determine pressure of hydrogen gas To do this you will need to find the total pressure for today and then subtract the pressure due to the water vapor in the gas collection tube. To do this you will use Dalton’s partial pressure 119 law. Dalton’s partial pressure law states that P1 + P2 + Pn = Ptotal. This just means that the partial pressure of each gas in a system contributes to the total pressure in the container. Vapor Pressure of water at various temperatures Temperature (K) Pressure (atm) Temp. (K) Pressure (atm) 288.2 289.2 290.2 291.2 292.2 293.2 294.2 295.2 0.0168 0.0179 0.0191 0.0204 0.0217 0.0230 0.0245 0.0261 296.2 297.2 298.2 299.2 300.2 301.2 302.2 303.2 0.0276 0.0295 0.0312 0.0332 0.0351 0.0372 0.0395 0.0418 In this case you have Phydrogen + Pwater vapor = Ptotal The total pressure in this experiment will equal the room pressure. The water vapor pressure will be determined based on room temperature and a table. These two pieces of information will allow you to find the partial pressure of the hydrogen gas. In this example assume the pressure today is 0.996 atm and the temperature is 19.0 oC which then is converted to 292.2 K (from Celsius to Kelvin add 273.14). Thus: Ptotal - Pwater vapor = 0.996 atm - .0217 atm = 0.974 atm Phydrogen Sample equation 4 5. To solve for R Using the information from above R=PV/nT R= 0.08314atm*L/mol*K P=0.974 atm T= 292.2 K V= 0.03785 liters n= 0.00155 moles D. Percent Error To find the percent error you just use the percent error formula. Percent Error = | Theoretical – Experimental | *100 Theoretical | 0.08206mol*K/atm*L- 0.08314mol*K/atm*L | 0.08206 L/mol * 100 120 Sample equation 5 =1.32% error Sample equation 6 Pre-Lab Questions: 1. a. Convert 0.879 atm to kPa b. Convert 787 mmHg to atm 2. The pressure on 2.50 L of gas changes from 0.987 atm to 0.400 atm. What is the new volume if the temperature is held constant? 3. When a rigid hollow sphere containing 6.80 liters of gas is heated to 601 K the pressure is 180.0 atm. How many moles of helium are in the sphere? 4. Determine the total pressure of a gas mixture that contains oxygen nitrogen and helium if the partial pressures of the gases are PO2 = 0.20 atm, PN2= 0.467 atm, PHe= 0.267 atm? Materials: Magnesium ribbon Volumetric gas collection tube Battery jar Ring stand Copper wire 400 mL beaker 10.0 mL 3.0 M HCl Burette clamp Procedure: 1. Obtain a piece of magnesium ribbon that is 1-2 cm long. Measure the length of the ribbon carefully and record the value to the nearest 0.01cm. 2. Record the mass of 1 meter of ribbon this information may be posted on the white board. 3. Fold the magnesium ribbon so that it can be encased in a small spiral cage made of fine copper wire. Let enough wire serve as a handle so that the cage can be placed approximately 2cm form the stopper. 4. Set up a ring and burette clamp in position to hold a 50 mL gas measuring tube which has been fitted with a one or two hole rubber stopper as per figure 2a. Place a 400 mL beaker about two thirds full of tap water near the ring stand. Figure 1 5. Tilt the gas measuring tube slightly and pour in 3M HCl to about the 10mL mark. 121 Figure 2 Figure 3 6. With the tube still tilted, slowly fill it with tap water from the beaker. While pouring rinse any acid that may be on the sides of the tube so that the final liquid near the opening of the tube contains very little acid. Avoid stirring up the acid layer in the bottom. You are attempting to have a layer of less dense water on top of a layer of more dense acid (See figure 3) . 7. Holding the copper coil by the handle, insert it into the tube until the wire case is positioned between the 49mL -50mL marks. Hook the copper wire over the edge of the tube and clamp it there by inserting the rubber stopper. When properly set up, the tube will have no air bubbles and the water will completely fill the holes in the stoppers as well as the tube. If the tube is not full, attempt to add water through the stopper holes by gently using a wash bottle. (see figure 3) Again, be certain that the acid layer stay, at the bottom undisturbed. 8. Cover the holes in the stopper with your finger and invert the tube in the container of water. Clamp the tube in place. The acid being denser than water will begin to stream down through the water and will eventually react with the metal. 9. Record observations, feel the gas tube with your hand and observe and record any changes in temperature. 10. After the reaction has stopped, wait about 5 minutes to allow the tube and its contest to come to room temperature. Bubbles clinging to the sides of the tube can be dislodged by gently tapping the tube. 11. Cover the hole(s) in the stopper with your finger and transfer the tube to a large cylinder or battery jar which is almost filled with water at room temperature. Raise or lower the tube until the level of the liquid inside the tube is the same as the level outside of the tube. This permits you Figure 4 to measure the volume of gases in the tube (hydrogen and water vapor) at the air pressure in the room. Read the volume with your eyes at the same level as the bottom of the meniscus. Record the volume of gas to the nearest 0.01mL. 12. Place your finger over stopper holes and remove the gas collection tube from the water. Pour the remaining acid solution down the sink. Rinse the tube with tap water. 13. Record the room temperature, record the air pressure in the room. 122 Data: Air pressure in room Air temperature in room Length of Mg ribbon used in experiment Length of magnesium standard Mass of magnesium standard Volume of gas produced Observations: ___________ ___________ ___________ ___________ (AS WRITTEN ON BOARD) ___________ (AS WRITTEN ON BOARD) ___________ Calculations: A. To find grams of Magnesium ribbon 1. Convert cm of Mg to grams of Mg 2. Convert grams of Mg to Moles Mg 3. Calculate temperature in Kelvin B. 1. 2. 3. To find theoretical moles of hydrogen gas produced Find the partial pressure of hydrogen gas Find experimental ideal gas constant. Percent Error C. Determine experimental molar volume 1. Determine the volume in liters that gas would occupy at STP. 2. You may wish to use PV=nRT. STP is defined as 1atm, 273 K Results: A. B. C. D. E. F. Grams of Magnesium ribbon Theoretical Moles of H2 gas produced Experimental volume of gas produced Experimental Gas Constant Percent error Molar Volume (liters/Mol, at STP) ___________ ___________ ___________ ___________ ___________ ___________ (STP= temp@273K and 101.2 kPa) Discussion: Discuss topics presented in “How to write a lab report” Post-Lab Questions: 1. In a second experiment, 0.113 grams of zinc reacted with 3.0 Molar HCl. The pressure is 0.995 atm and the temperature is 21 oC. 38.35 mL of gas were collected over water. From this data calculate the experimental value for R. What is the percent error for the experimental R (the theoretical value of R is 0.08206 mol*K/atm*L)? References: Unknown (unknown). Experiment 10 Reaction of a metal with hydrochloric acid. unknown Willbraham, A. Staley, D. Matta, M. (1995) Chemistry 4th edition. Addison-Wesley: Menlo Park, CA 123 Experiment 19 Heat of Neutralization Purpose: The purpose of this laboratory activity is to compare the heat of neutralization of two different acids and to calculate the experimental heat of formation for water. Hypothesis: To develop the hypothesis, compare the net ionic equations for the reaction of nitric acid and sodium hydroxide to the reaction of hydrochloric acid and sodium hydroxide. From the net ionic equation develop a hypothesis which compares the heat of neutralization for nitric acid to the heat of neutralization for hydrochloric acid. Background: Driving Forces There are two primary reasons that reactions take place: enthalpy and entropy. Some chemists refer to these as driving forces, as in the question posed in many chemistry classes, "What is the primary driving force in this reaction…?" Entropy is the measure of disorder in a system. This is beyond the scope of this laboratory activity. However it is sufficient to think that in general the universe tends towards disorder. In terms of energy, generally energy is released as things move toward disorder and it takes energy to either organize things or to fight off entropy. Enthalpy is the measure of heat (molecular movement) that is present in a substance. The universe has a fixed amount of enthalpy. There is only so much energy, and you cannot create or destroy energy. When a bond is formed the atoms become more stable for each atom now has each of its orbitals Energy diagram filled. As a result when a bond is formed the atoms have less energy. Thus when a bond is formed energy is released to the environment. The atoms constantly attempt to reach the lowest possible energy state, just as a ball rolls down hill to the lowest possible energy state, a state of rest with no potential energy. Atoms react to make bonds releasing energy in effect rolling down hill and giving off energy as they do so. For instance two materials that have very little energy are water and carbon dioxide. Carbon dioxide and water are the products of so many reactions Figure 1 because they have a low potential 124 energy. Water and carbon dioxide are like the rock at the bottom of the hill. Reactions which release energy are called exothermic (see Figure 2). Exothermic reactions release energy to the environment. They feel hot to us; this is because energy is released when the new bonds are formed to create the product bonds. Exothermic reaction gives off energy (Figure 2) Some reactions are Diagram of a endothermic reaction endothermic (see Figure 3). A reaction that releases energy to the environment That is, they require energy from the environment to proceed. This indicates that the product bonds formed release less energy than is required to break the reactant bonds. Thus energy from the environment is needed to keep breaking bonds. As energy is repeatedly removed from the environment the environment starts to feel cold. This is because heat energy is removed from it. As heat energy is removed the molecules begin to move more slowly and this is what is measured as cold. Endothermic reactions feel cold to the touch, because they require energy Figure 3 from the environment to proceed. They get this energy from the surrounding air or any chemist who happens to touch the reaction vessel. Why would a reaction that constantly requires heat be able to proceed spontaneously? A spontaneous reaction is a reaction that will take place without further addition of heat energy. That is, once the activation energy is met (the energy required to start the reaction) the reaction will proceed without any additional input of energy from the surroundings. If a reaction is endothermic (meaning it draws energy from its surroundings to proceed) what could possibly be its driving force? The answer is entropy. 125 Heat: Heat is the transfer of energy due to differences in temperature. Heat is not the same as temperature. When heat is released in a chemical reaction the reaction is referred to as an exothermic reaction. When heat is absorbed the reaction is termed endothermic. The heat released in a reaction can be measured using a known quantity (mass) of water. As heat is released from the chemical reaction it is absorbed by the water and the water’s temperature goes up. The amount that the temperature goes up indicates how much energy was released in the reaction. There are several different types of reactions that release energy in the form of heat. Heat of formation: The amount of heat involved in the production of one mole of a substance. Heat of combustion: The amount of heat produced when 1 mole of a material is reacted with oxygen. Heat of solution: The amount of energy produced when one mole of a material is dissolved in water. Heat of vaporization: The amount of energy involved when one mole of a substance changes phase from a liquid to a gas (vaporization) or changes phase from gas to a liquid (condensation). Heat of fusion: The amount of energy involved when one mole of a substance changes phase from a liquid to a solid (freezing) or changes phase from solid to a liquid (melting). Heat of neutralization: The amount of heat involved when on mole of water is created when an acid reacts with a base, the literature value is -55.90 kJ/mol. The heat that is released or absorbed in a chemical reaction (which maintains a constant pressure) is called ∆H. The ∆ is a capital delta (Greek letter), which symbolizes “a change in”. The ∆ indicates a difference in the amount of heat in a system. This is expressed mathematically as ∆H= Hf-Hi Hf = heat final Hi = heat intial If heat is released the ∆H is negative and the reaction is exothermic. If heat is absorbed the ∆H is positive and the reaction is endothermic. In the lab you will be using a simple calorimeter. When acid is neutralized water is produced. This reaction releases heat. Because both the acid and the base in this reaction are dissolved in water the heat given off by this reaction will be absorbed by the surrounding water. If this temperature increase is carefully measured the amount of energy released can be calculated. Some heat energy may be lost to the surroundings. This heat loss will be small because of the insulated containers. The amount of heat lost to the surrounding air will be accounted for by plotting a cooling rate graph. 126 The heat from the reaction is absorbed by the water. The temperature of water increases. However some of the heat escapes to the surrounding air. Using the graph a line can be drawn to extrapolate the temperature at the time the solutions are mixed. Temp final The greatest difference between the temperature of the reactants before they were mixed and the products after they were mixed is the ∆T, or change in temperature. Temp initial This will be used in the equation q=mc∆T. Cooling-Rate Graph Pre-Lab Questions: 1. Explain the following in regards to enthalpy and entropy. a. When sodium chloride is dissolved into water the surrounding water gets colder. Is this physical change endothermic or exothermic? Why is this reaction spontaneous? Provide evidence for your answer. b. Many gases are held under enough pressure that the molecules come so close together that they become liquid. When the pressure is released the compressed gas in liquid phase returns to its gaseous state. When this happens the surroundings become very cold. This is in essence how a refrigerator works. Is this physical change endothermic or exothermic? Why is this reaction spontaneous? Provide evidence for your answer. 2. The temperature of 100.00mL of water was raised 25.5 oC. How much heat in joules was added to the water? How much heat in calories was this? (the specific heat of water is 4.184 j/g oC or 1.00 cal/g oC) q=mc∆T Where q = energy m = mass c = the specific heat ∆T = change in temperature 3. If 50.0mL of 1.80 M HCl was added to 50.0 mL of 1.80 M NaOH, calculate the molarity of the resulting NaCL solution. (Hint: Calculate the moles of NaCl made. The resulting solution has a volume of 100mL.) Materials: Calorimeter Thermometer Watch with second hand 1.0 M NaOH 1.0 M HCl 1.0 M HNO3 127 100.0 mL graduated cylinder Procedure: Heat of neutralization: HCl and NaOH 5. Set up calorimeter as shown in illustration. This is the same way it was set up in the previous experiment. Make sure the thermometer does not touch the bottom or sides of the cups. 6. Obtain about 50.0mL of 1.0 M HCl with the graduated cylinder. Pour the acid into the calorimeter. Record actual volume. 7. Measure about 50.0 mL of 1.0 M NaOH with a graduated cylinder. Pour contents into a beaker. Allow this solution to sit in the beaker while reading the temperature of the HCl in the calorimeter. Record actual volume. 8. Measure and record the temperature of the acid in the calorimeter to the nearest 0.1 oC. This it the first temperature, T0. After 30 seconds record the temperature again, this is T1. Continue recording data points for 2 minutes and 30 seconds. Assume the NaOH solution is at the same temperature. Is this a valid assumption? 9. At exactly 3 minutes, T6, quickly pour the NaOH into the calorimeter begin the chemical reaction. You cannot measure the temperature accurately at this time because it will be changing very rapidly. After you have poured the NaOH into the calorimetric immediately replace the lid, thermometer and stirrer and gently stir the mixture. 10. 30 seconds after mixing, T7, read the thermometer and record the temperature. Read and record the temperature at 30 second intervals. Stir the mixture occasionally. Continue stirring and recording data for 4 more minutes. 11. Draw a graph of temperature (Y-axis) vs. Time (X-axis). Connect the points from T0 to T5. The line should be nearly horizontal. Use this line to extrapolate what the initial temperature at T6 was. 12. Draw a best fit line between points T7 -T16. Use this line to extrapolate a second value for T6. This line is the final temperature. 13. The difference between the two extrapolated lines is the ∆T. Heat of neutralization: HNO3 (nitric acid) and NaOH. 1. Repeat the above using 50 mL of 1.0 M HNO3 and 50mL of 1.0 M NaOH. Calculate ∆T as before with extrapolated lines on a graph. Data: Trial 1à Volume HCl and NaOH HCl NaOH Trial 2 à Write Data In HNO3 Note NaOH Book Temperature at various times… T0 T1 T2 T3 T4 T5 T6 T7 base Write Data In Lab Note Book added T8 T9 Etc Record all actual volumes used, use these values for all calculation. Record all observation. 128 Calculations: 1. Write and balance the reaction for aqueous HCl reacting with NaOH. 2. Determine how many moles of water are produced when 1 mole of HCl reacts with 1 mole of NaOH. 3. For trial one use the volume of HCl or NaOH which ever one is less (it acts as a limiting reactant) and find the moles. 4. Using dimensional analysis find the moles of water produced based on item 3 above. 5. Using dimensional analysis calculate the mass of solution used in the reaction. Assume that the solution formed is volume of acid plus vol of base. The density of the salt solution is 0.9982 g/mL. 6. From the mass of the solution (m), the specific heat (c)of the solution and its change in temperature (∆T). Calculate the amount of heat evolved by the reaction. The specific heat of this solution is 4.017 J/goC, Heat absorbed by water=mass of solution * specific heat * ∆T q= mc∆T 7. Calculate the amount of heat evolved by the reaction. This is the amount of heat absorbed by the water, multiplied by -1. This ∆H is for the number of moles of water your reaction produced. Heat evolved by the reaction = -(heat absorbed by the water) = ∆H ∆H = -q 8. Calculate the value of ∆H for 1 mole of water being produced by this reaction. Refer to the answers from calculation steps 3 and 7. This value of ∆H is the heat of neutralization. Be certain that the units are J/mole. 9. Is the reaction exothermic or endothermic? What sign should the ∆H have, - or +. 10. Repeat the above steps for the neutralization of acetic acid. Assume that the sodium acetate solution produced has the density and specific heat as the sodium chloride solution. Results: A. B. C. D. E. ∆H of water from the HCl reaction. ∆H of water from the acetic acid reaction. Report the theoretical ∆H of water. Report % error from the HCl reaction. Please background information for theoretical value. Report % error from the HNO3 reaction. Please background information for theoretical value. Discussion: Discuss the results of your experiment in regards to supporting or refuting the stated hypothesis (do not use the word “prove”). Is the heat of neutralization the same regardless of which acid was used? Explain your answer using the net ionic equation. Answer the discussion questions as presented in your “How to Write a Lab Report.” Post-Lab Questions: 1. When 0.25 mole of NH4Cl is dissolved in 1.00 L of water the temperature drops 12.5 oC. (Assume 1013g of solution and a specific heat of 4.02 J/g oC). What is the heat of solution for ammonium chloride? 129 Experiment 20 Molecular Model Mobile Project I want to get this assignment to you as soon as possible because some of you may be going places for spring break. I would strongly urge you to read this before spring break arrives so that any misunderstandings get cleared up before you leave. Essentially you will research a chemical, determine its molecular structure, and build a model. I would strongly urge you to not spend a lot of money. Also I am rather partial to space filling models, though a ball and stick model is fine. Molecule Four different elements 1 Three different shapes 1 Two different bonds 1 Choose a molecule that will not embarrass me to hang in the classroom, this is subjective so be careful, I like cool but be reasonable. Really cool molecules might earn extra credit. Your model may not be scored if another student is doing the same chemical. Please check the sign up sheet so that you will receive credit. Paper If your paper even appears to be copied from some other source you will be asked to rewrite the whole document for reduced credit. Typed, Times New Roman, 1" margins, 12 point font, clean clear paper. 1 Paragraph of the chemical’s history, who discovered it and why. If it is a naturally 1 occurring material who first determined the chemical formula? Trivial info The societal context of this chemical? (Why is it important? How is it used?) 1 A complete reference section in APA format 1 internet citation 1 2 print book references (textbook, encyclopedia, other books, etc.) 1 Journal citation 1 One reference citation (CRC, Merck) 1 The Model Your model will hang from the ceiling from no more than 2 points and the lowest point of model will be less than 75cm (including sign) from the ceiling. If it does not meet these requirements you will receive no credit. The model is sturdy; it will not fall apart The model is not made of food Bond angles are correct Relative size of atoms is close to accurate. Things in one energy level are all about the same size, things in lower energy levels are smaller; things in higher energy levels are larger. A sign with the name of the molecule, the MW, and formula in IUPAC notation Different colors to represent different atoms Presentation I am dressed professionally, speak professionally, act professionally Presentation lasts between 1 1/2 and 3 minutes The presentation is informative and accurate 130 1 1 1 1 1 1 1 1 1 Molecular Model Mobile Project SCORE SHEET TO BE TORN OUT AND TURNED IN WITH PAPER. PAPER IS DUE BEFORE YOU GIVE YOUR TALK. Scoring Molecule Four different elements Three different shapes Two different bonds 1 1 1 Choose a molecule that will not embarrass me to hang in the classroom, this is subjective so be careful, I like cool but be reasonable. Really cool molecules might earn extra credit. Your model may not be scored if another student is doing the same chemical. Please check the sign up sheet so that you will receive credit. Paper If your paper even appears to be copied from some other source you will be asked to rewrite the whole document for reduced credit. Typed, Times New Roman, 1" margins, 12 point font, clean clear paper. 1 Paragraph of the chemical’s history, who discovered it and why. If it is a naturally 1 occurring material who first determined the chemical formula? Trivial info The societal context of this chemical? (Why is it important? How is it used?) 1 A complete reference section in APA format 1 internet citation 1 2 print book references (textbook, encyclopedia, other books, etc.) 1 Journal citation 1 One reference citation (CRC, Merck) 1 The Model Your model will hang from the ceiling from no more than 2 points and the lowest point of model will be less than 75cm (including sign) from the ceiling. If it does not meet these requirements you will receive no credit. The model is sturdy; it will not fall apart The model is not made of food Bond angles are correct Relative size of atoms is close to accurate. Things in one energy level are all about the same size, things in lower energy levels are smaller; things in higher energy levels are larger. A sign with the name of the molecule, the MW, and formula in IUPAC notation Different colors to represent different atoms Presentation I am dressed professionally, speak professionally, act professionally Presentation lasts between 1 1/2 and 3 minutes The presentation is informative and accurate 131 1 1 1 1 1 1 1 1 1 132 Appendix Solubility Rules: 1. 2. 3. 4. 5. 6. 7. Salt containing sodium, and potassium ions, ammonium and acids are always soluble. Salts containing nitrates, chlorate, perchlorate and acetate are always soluble. All chlorides, bromides and iodides are soluble except for those of silver, lead II, and mercury I, which are insoluble. All sulfates are soluble except of those of calcium, silver, strontium, barium, mercury I, mercury II and lead. All hydroxides are insoluble expect those of alkali metals, which are soluble, and hydroxides of calcium, barium, and strontium, which are moderately soluble. All sulfites , carbonates, chromates and phosphates are insoluble expect those of ammonium and alkali metals. All sulfides and insoluble expect those of ammonium, the alkali metals and the alkaline earth metals, which are soluble. Bond Energy kJ/mol H-H H-F H-Cl H-Br H-I C-H C-C C-Si C-N C-O C-P C-S C-F C-Cl 432 565 427 363 295 413 347 301 305 358 264 259 453 339 Bond C-Br C-I C=C C=N C=O C=O (CO2) F-F F-Cl F-Br F-I N-H N-N N-P N-O Selected Specific heat capacities [J/g* oC] H2O(liquid) 4.18 Iron 0.46 Acetone Benzene Hofus kJ/mol 5.72 9.87 Energy kJ/mol Bond Energy kJ/mol Bond Energy kJ/mol Bond Energy kJ/mol 276 216 614 615 745 799 159 193 212 263 391 160 209 201 N-F N-Cl N-Br N-I O-H O-P O-O O-S O-F O-Cl O-Br O-I O2 N=N 272 200 243 159 467 351 204 265 190 203 234 234 498 418 N=O N=N S-H S-S S-F S-Cl S-Br Si-Si Si-O Si-S Si-Cl Si-Br Si-I P-H 607 945 347 266 327 271 218 226 368 226 381 310 234 320 P-P P-F P-Cl P-Br P-I C=C C=N C=O Cl-Cl Cl-Br Cl-I Br-Br Br-I I-I 200 490 331 272 184 839 891 1070 243 215 208 193 175 151 H2O(solid) Aluminum Hovap kJ/mol 29.1 30.8 1Atm = 101.3 kilo pascal 1Atm = 760 mmHg 1Atm = 760 Torr 1Atm = 1013 mBar PV = nRT P1 V 1 = P2 V 2 (P1V1)/T1 = (P2V2)/T2 Pt = P1+P2+P3+Pn R = 0.08206 Density of water Temp oC g/ml 0 (solid) 0.9150 0 (liquid) 0.9999 4 1.000 20 0.9982 40 0.9922 60 0.9832 80 0.9718 100(steam) 0.0006 % error = 100* (|Exp-Theo|/Theo) % yield = 100* (Exp/Theo) Absolute error = |Theo – Exp| 133 [J/g* oC] 2.1 0.90 Water Methanol H2O(gas) ethanol [J/g* oC] 1.7 2.4 Hofus kJ/mol 6.01 .94 Hovap kJ/mol 40.7 8.2 1cm = 2.54 inches (∞ sig figs) cm3 = 1 mL 1 mole = 6.022 x1023 Q = mCΔT λ* ν = c 2.998x108 meter per second Planck’s constant = 6.626x10-34 J*sec ΔE = h * ν RH = 2.178x10-18 J M1V1=M2V2 1 1 2 ΔE = -RH Nf Ni 2 ( )