ChemLab Manual

ChemLab Manual - Chemistry 1800U Laboratory Manual Winter,...

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Unformatted text preview: Chemistry 1800U Laboratory Manual Winter, 2012 Senior Laboratory Instructor: Richard Bartholomew UA 4070 Table of Contents LABORATORY CALENDAR . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1 SAFETY RULES . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3 SAFETY PROCEDURES . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6 GOOD LABORATORY PRACTICE . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10 LABORATORY REPORTS . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13 PROCESS FOR RE-SCHEDULING MISSED LABORATORIES . . . . . . . . . . . . . . . . . . . . . . 17 WEIGHING TECHNIQUES . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 19 VOLUMETRIC TECHNIQUES . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 23 OPERATION OF THE BUNSEN BURNER . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 29 SUCTION FILTRATION . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 31 SIGNIFICANT FIGURES . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 33 LABORATORY REPORT RECEIPT SHEET . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 39 SAFETY ACKNOWLEDGEMENT . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 40 LOCKER EQUIPMENT LIST . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 41 1. LABORATORY SAFETY AND ORIENTATION . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 43 i 2. ATOMIC EMISSION SPECTRA . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 52 3. MEASUREMENT OF THE IDEAL GAS CONSTANT, ‘R’ . . . . . . . . . . . . . . . . . . . . . . . . . 64 4. EFFECT OF CONCENTRATION AND TEMPERATURE ON REACTION RATE . . . . . . 80 5. THERMOCHEMISTRY AND THE MEASUREMENT OF THE ENTHALPY OF DISSOCIATION FOR ACETIC ACID . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 93 6. ELECTROCHEMISTRY AND VOLTAIC CELLS . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 116 ii LABORATORY CALENDAR Winter, 2012 Laboratories begin January 13, 2012. The exact dates of your laboratories depend on the CRN in which you are registered. These dates can be found on your schedule or on the preview of available courses on MyCampus. CRN’s 72282 and 72283: Dates Experiment January 13 1 - Laboratory Safety and Orientation January 20 2 - Atomic Spectroscopy February 3 3 - Effect of Concentration and Temperature on Reaction Rate February 17 4 - Measurement of the Ideal Gas Constant March 9 5 - Thermochemistry August 23 6 - Electrochemistry 1 All other CRN’s: Dates Experiment January 16 - 27 1 - Laboratory Safety and Orientation January 30 - February 10 2 - Atomic Spectroscopy February 13 - 16 3 - Effect of Concentration and Temperature on Reaction Rate February 20 - 24 Reading Week - no labs February 27 - March 2 3 - Effect of Concentration and Temperature on Reaction Rate March 5 - 16 4 - Measurement of the Ideal Gas Constant March 19 - 30 5 - Thermochemistry April 2 - 13 6 - Electrochemistry 2 UOIT Chem. 1800U, S11; Introduction - 8.13 SAFETY RULES Safety in the laboratory is of the utmost importance to your instructors and also must be to you, the student. Everybody’s safety depends on each student adhering to the safety rules and procedures outlined in this manual. The following rules must be obeyed and will be rigorously enforced. 1. Students arriving more than 10 minutes late to a laboratory period will not be allowed access to the laboratory. Students must arrive promptly to receive instructions about the safe conduct of the experiments. 2. ALL students must wear eye protection. For people with eyeglasses it must be worn over regular glasses. This rule will be vigorously enforced. 3. Contact lenses should not be worn in the chemistry laboratory. 4. All students must wear a lab coat. A 100% cotton lab coat is best. Lab coats protect skin and clothing from chemicals. The front of he lab coat should be buttoned while working in the laboratory. 5. Open shoes or sandals are forbidden. They expose the feet to spilled chemicals. 6. Loose or bulky clothing presents a hazard and should not be worn in the laboratory. 7. Clothing that exposes large areas of the body (e.g. shorts, tank tops) must not be worn in the laboratory. 3 UOIT Chem. 1800U, S11; Introduction - 8.13 8. Long hair must be tied back. The simple rule: if it can be tied back, it must be tied back! 9. Aisles must be kept clear of boots, coats, knapsacks, etc. 10. Students must know the locations of: a) fire extinguishers b) eye wash stations c) emergency showers d) emergency exits e) fire blanket f) fire alarms 11. Eyewash stations, shower, fire blanket and fire extinguishers must not be obstructed. 12. Absolutely no smoking, eating or drinking in the laboratory. Neither food nor drink may be brought to the laboratory. 13. No laboratory is to be started without an instructor present. Unauthorized experiments are strictly forbidden. Experiments must not be left unattended. 14. When diluting acid be sure to add the acid to the water. Add the acid slowly and with plenty of stirring. Diluting acids generates huge amounts of heat (it is a very exothermic process). Keeping the water in excess allows this heat to be more effectively dissipated. If the heat is not adequately dissipated the rapid heating may cause the solution to be ejected from the beaker, injuring the experimenter. If the acid is added to the water, the ejected solution will be comparatively dilute and therefore less dangerous. 4 UOIT Chem. 1800U, S11; Introduction - 8.13 15. NEVER, EVER pipette by mouth. Safety bulbs are provided for pipetting. When pipetting by mouth it is very easy to lift the tip of the pipette above the level of the liquid being pipetted As a result, solution will enter the mouth. 16. Never point the mouth of a test tube at yourself or others. This is especially true if the test tube is being heated. 17. NEVER taste chemicals or solutions. Should you get any chemicals in your mouth, do not swallow. Rinse your mouth thoroughly and then consult an instructor. 18. If you must smell a chemical, waft the vapours toward your face with your hand. Do not place the container directly under your nose - you may be overcome by the gas. 19. Never put broken glass in the regular garbage. To do so presents risks to janitors and cleaning staff. Broken glass should be swept up with a dust pan and brush and disposed in the receptacle for broken glass. ONLY broken glass should be placed in this container. NOTE: Students who arrive for laboratories inappropriately addressed may be forbidden from performing the experiment. Students who violate the safety rules may be dismissed from the laboratory period. In such cases the student will be given a grade of zero for the laboratory report. 5 UOIT Chem. 1800U, S11; Introduction - 8.13 SAFETY PROCEDURES Fire Fire is one of the most serious problems that may be faced in the chemistry laboratory. For no other safety issue is the adage “an ounce of prevention is worth a pound of cure” more applicable. In the event of a small fire, use the fire extinguisher. Remove the pin by twisting and pulling it out. Direct the nozzle of the extinguisher at the base of the fire and squeeze the trigger. Ensure that at all times you are between the fire and your escape route. Small fires can rapidly and easily become large fires. If the fire cannot be safely extinguished in 30 s, leave the room. For large fires pull the fire alarm. If in doubt, pull the fire alarm and leave the building. The more time people have to escape the more likely they will be successful. Turn off all services (burners, hot plates, water, etc.); leave the room The last person out of the room must close the door – this is the single most effective way to slow the spread of fire. Leave the building by the nearest exit and move well away from the doors. The lab demonstrator should call emergency services and give all necessary details to the fire crews. Clothing on Fire DO NOT RUN! Wrap the victim in the fire blanket or use a lab coat to smother the flames. “Drop and Roll” is effective. The emergency shower can be used. Have someone call the emergency number and get medical assistance. 6 UOIT Chem. 1800U, S11; Introduction - 8.13 Fire Alarm UOIT uses a “two stage” fire alarm system. The first stage is a warning that there may be an emergency in the building. In the first stage the fire alarm rings at a rate of 30 rings per minute. At the first stage of the alarm, stop all your experiments when it is safe and convenient to do so and prepare to leave the building. Await further instructions. If the fire alarm progresses to the second stage (louder, more frequent alarm), immediately stop all experiments. Turn off all services (gas, electricity, water, etc.). Before leaving the laboratory make sure the hallway is free of smoke and fire. If it is, leave the laboratory and exit the building by the quickest route. Close the door to the laboratory. Once outside, move away from the building and assemble at the western end of the first floor atrium of UB. If UB is closed, meet in the car park immediately to the east of UA. Your TA will do a headcount to ensure everyone is out. Do NOT wander away without telling your TA! Under no circumstances should you return to the building until told by the fire department that it is safe. Escape Routes for First-year Laboratories Room Primary Escape Secondary Escape UA 3420 Turn left, descend staircase at Turn right, then left along southern northwest corner of atrium. Exit corridor. Descend staircase at southeast building by north doors. corner of building. Exit building by doors at southeast corner. UA 3480 Turn left, descend staircase at Go straight along southern corridor. northwest corner of atrium. Exit Descend staircase at southeast corner of building by north doors. building. Exit building by doors at southeast corner. 7 UOIT Chem. 1800U, S11; Introduction - 8.13 Medical Situations UOIT maintains a service called the Campus Emergency Response Team (CERT). The CERT is staffed by volunteers who have had extensive first aid training. If CERT assistance is required, the CERT team can be contacted at x-2400. Pregnancy If you are, think you may be or are planning to become pregnant, you should be aware that the risks to a mother and a developing foetus from the chemicals used in these laboratories are generally unknown. You may wish to consider deferring taking this course until after your baby has been born. If you decide to take this course, consultation with your family doctor is strongly encouraged. If you would like the Material Safety Data Sheets for the chemicals in this course, they will be provided. Fainting If at any time you feel light headed, dizzy or that you may faint, immediately sit down on the floor. Fainting itself is rarely harmful; falling because of it may cause injury. Call for medical assistance. If someone should faint, try to assess whether they have been injured by the fall. Make them as comfortable as possible and call for medical assistance. Cuts Minor cuts should be treated with plenty of cool water. Ensure no foreign objects are present (such as glass) and apply an appropriate dressing. Seek medical attention at student services. For more serious cuts the victim should sit or lie down and keep the cut elevated. Bleeding 8 UOIT Chem. 1800U, S11; Introduction - 8.13 should be kept to a minimum by applying direct pressure (assuming no foreign objects are present in the wound). Call for emergency medical assistance. Burns Burns can be either thermal (caused by heat) or chemical. In both cases the first step is to apply plenty of cool water. If the chemical reacts with water, remove it by brushing it from the skin. If necessary, seek medical attention. Allergies and Other Pre-existing Conditions If you suffer from severe allergies or other pre-existing medical conditions such as epilepsy, diabetes, etc., it may be helpful (although not required) to alert your instructor(s) and to advise them of any necessary precautions or first-aid treatment. 9 UOIT Chem. 1800U, S11; Introduction - 8.13 GOOD LABORATORY PRACTICE 1. Use clean equipment. Dirty apparatus can lead to poor results (and therefore, poor grades!) or to unexpected and potentially dangerous reactions. Beakers and flasks can be cleaned with soap and water followed by thorough rinsing with tap water and deionized water. Soap should be avoided when cleaning volumetric glassware such as pipettes, burettes and volumetric flasks. 2. Once a chemical has been removed from its original container it must NEVER be returned to the container. To do so risks contaminating the entire stock and thereby ruining the experiments of others. If you have excess, it is cheaper to throw it away (appropriately) than to risk contaminating the stock. 3. Chemicals are expensive and solutions can be time consuming to prepare. DO NOT WASTE CHEMICALS. Read the manual very carefully and take only what you need. 4. Carefully read the labels of reagent bottles to ensure you are getting the right chemicals. Be sure to properly label apparatus containing chemicals or solutions. Accidents may result from the mistaken mixing of chemicals. 5. Remember to re-cap reagent bottles immediately after use. This prevents wastage of chemicals by contamination and spillage. Do not allow caps to become contaminated through contact with the bench top or other chemicals. 6. Do not remove reagent bottles to your benches. This causes frustration and unnecessary delay for other students. 10 UOIT Chem. 1800U, S11; Introduction - 8.13 7. Certain chemicals used in these laboratories are hazardous or present an environmental risk. Dispose of these chemicals in the appropriate waste dumps provided. 8. Set-up apparatus so that it is well back from the edge of the bench. All services (gas taps, water taps, electrical outlets, etc.) should be readily accessible. 9. Keep your work area clean, tidy and well organized. Cluttered work areas can lead to accidents. Clean up all spills quickly. At the end of the lab period clean your work area. 10. Do not dispose of solid or insoluble materials in the sinks. To do so will inevitably lead to clogs of the plumbing and to floods. 11. Large spills of acids and bases can be treated with the appropriate spill kit. Afterwards the bench should be washed with plenty of cold water. 12. Mercury spills should be cleaned up immediately. The vapour from mercury is quite toxic. Mercury spills are generally treated with powdered sulphur which reacts with the mercury to form mercury sulphide. The resulting solid is collected and treated as chemical waste. 13. Do not wander aimlessly in the laboratory. 14. Never interfere with the work of other students unless that work presents an immediate hazard to yourself or to others. 15. At the end of the laboratory period clean all equipment and return it to its appropriate place. 11 UOIT Chem. 1800U, S11; Introduction - 8.13 Handling Glassware 1. Apparatus that can roll (such as thermometers, etc). should be placed on the bench at right angles to the bench to prevent it rolling onto the floor. 2. Suction flasks (or Büchner flasks) may collapse violently under vacuum if cracked or otherwise weakened. Inspect suction flasks before using. Do not strike or tap a suction flask while it is under vacuum. 3. Chipped, broken or cracked glassware should be discarded. Heating cracked glassware is very dangerous - the glassware may shatter. Inspect glassware before using. 4. When inserting glass tubing into a stopper match the tubing to the size of the hole. Sometimes the tube can be lubricated with water or glycerol. To protect hands from being cut, wrap tube in a towel before inserting into the stopper. Apply force to the tube lengthwise while slowly twisting the tube. 5. To break glass tubing, use a triangular file to scratch the tubing at the point of the break. Moisten the scratch and wrap the tube with a towel. Place thumbs against the glass tubing on the opposite side of the scratch. Press against the tube while pulling hands apart. Fire polish the ends of the tubing before using. 12 UOIT Chem. 1800U, S11; Introduction - 8.13 LABORATORY REPORTS 1. Attendance at laboratories and submission of a laboratory report are compulsory. Failure to attend a laboratory or to submit the report may result in a grade of zero. A student may be excused from a laboratory only by providing appropriate documentation. 2. Documentation for missed laboratories should be submitted to the senior laboratory instructor . 3. In exceptional circumstances it may be possible to re-schedule an experiment. When possible (in the case, for example, of religious holidays), requests for re-scheduling should be made in advance. Laboratories may only be re-scheduled with the permission of the senior laboratory instructor. See below for the process for re-scheduling laboratories. 4. A student who fails to submit more than TWO (2) laboratory reports will not receive credit for the laboratory portion of the course. This may result in failure of the course. 5. With the exception of the last experiment laboratory reports are due at the beginning of the next period. The report for the last experiment is due at the end of the laboratory period. 6. Marks will be deducted from late lab reports at a rate of 2 marks per day. No report will be accepted if it is more than 4 days late. 7. The laboratory report must be typewritten or written in ink - the choice is yours. Two (2) marks will be deducted if the report is not typewritten or written in ink. 8. All original experimental data must be recorded in ink. Before leaving the laboratory have 13 UOIT Chem. 1800U, S11; Introduction - 8.13 your data signed by an instructor. The signed data must be submitted with the laboratory report. If no signed data are submitted with the report, two (2) marks will be deducted. 9. Any errors in recorded data should be corrected by drawing a single line through the erroneous data and writing the corrected data close by. The original data should remain legible. 10. White-out (or similar product) must never be used on laboratory reports. Two (2) marks will be deducted for using white-out (or similar product). 11. Laboratory reports should be legible and well organized. The grammar, style, and spelling will be assessed. Persistently poor spelling or grammar will be penalized. 12. Sample calculations must always be shown. In some cases the same calculation will be repeated (in titrations, for example) - it is not necessary to show every individual calculation. All steps should be shown so the marker can understand how the calculation was done and locate any errors in the method or arithmetic. Failure to show calculations will result in substantial loss of marks even if the final answers are correct. 13. Academic misconduct is a serious offence and will be punished. Academic misconduct includes (but is not limited to): plagiarism (copying) of lab reports, submitting false data, mispresenting data or using data from other students without the permission of the instructor. Details about academic misconduct, punishments and appeals procedures are given in the university calendar. It is the student’s responsibility to read and understand these regulations. 14. Teaching assistants are provided with uniform marking schemes for all the experiments. However, within the context of these marking schemes teaching assistants will execute their own judgement. At the end of the semester laboratory grades will be adjusted to account for 14 UOIT Chem. 1800U, S11; Introduction - 8.13 variation between individual teaching assistants. This may result in an increase or a decrease in your laboratory grade. For semesters in which there is more than one teaching assistant, an overall class average for the laboratory marks will be calculated and the marks for each individual teaching assistant will be scaled to this value. No correction will be applied if the correction would be less than one mark to the final laboratory grade for the course. 15. A “performance evaluation” (worth 5 marks in 30) is part of your assessment in the laboratory. This is an assessment of your general conduct, preparation, safety, attitude etc. Graphs for First Year Laboratories Graphs may be drawn by hand or by computer - the choice is yours. If doing by hand, you must use graph paper that uses at least ten divisions per centimetre and each data point should be circled. When using a computer, the spreadsheet from which the graph was generated should also be submitted. This will allow the marker to assess whether the data have been correctly plotted. Whether drawn by hand or by computer follow these guidelines: 1. Unless instructed otherwise, the independent variable should be on the x-axis and the dependent variable should be on the y-axis. 2. Use appropriate scales for the axes. Scales should be chosen so that the data fill the available space. When plotting by hand, the scale should also be chosen to allow easy plotting and reading of values; a scale in which 10 divisions represents 2/3 of a unit is not convenient! The intersection of the axes does not have to take place at the origin (0,0). 15 UOIT Chem. 1800U, S11; Introduction - 8.13 3. The axes must be labelled. The name of the quantity and the units in which it is measured should be written beside the axis, e.g.: “temperature / 0C” or “concentration / mol L-1". 4. Major divisions on the axes should be labelled. 5. A title must be added to each graph. Enough information must be in the title to make it clear what the data represent. For example: Experiment 5: Concentration vs. Reaction Rate for the Reaction between Hydrochloric Acid and Thiosulphate Ion (2HCl(aq) + Na2S2O3(aq) 6. v S(s) + SO2(aq) + H2O(l)) at 200C Do not play “connect-a-dot” with the data. When appropriate, draw a smooth curve to “best fit” the points. In many cases the curve may be a line. When using the computer, the line or curve may be calculated by regression analysis. If this is the case, the function used for the regression should be shown with the graph. The values of the adjustable parameters should also be included. 7. DO NOT use data points for the calculation of the slope. If the slope is calculated “by hand” the points used for the calculation must be clearly indicated on the graph. Clearly label the points with ‘x’ and ‘y’ co-ordinates. Format for First-Year Chemistry Laboratory Reports Laboratory reports in first-year chemistry are not “formal” reports. That is, you do not need to submit reports with a “Introduction”, “Method”, “Results and Discussion” sections. You need to submit your original data (signed by the TA), your sample calculations (including graphs) and the answers to any questions posed in the laboratory manual. 16 UOIT Chem. 1800U, S11; Introduction - 8.13 PROCESS FOR RE-SCHEDULING MISSED LABORATORIES If miss your laboratory period (for whatever reason), you should complete the “Request for Re-Scheduled Laboratory” form available in the “Laboratories” section of the WebCT site for the course. If missing the laboratory period is foreseeable (for religious observance, varsity athletics, etc.), the request should be submitted in advance. To complete the form, you must suggest an alternative time to attend a laboratory period. It is solely the responsibility of the student to find a laboratory period when they can attend. A list of scheduled laboratory periods is available on the “Preview of Available Courses” at MyCampus: There must be sufficient space in the laboratory period. In other words, the number of students enrolled in the section must be less than the maximum number of students allowed in the section. Note, re-scheduling is done on a first-come, first-served basis; a laboratory period may be full because of re-scheduling even if MyCampus indicates a vacancy. Once completed, the form should be submitted via e-mail to the senior laboratory instructor ( for the course for approval. Do NOT submit the form through WebCT. You should use your MyCampus email account to submit the form. Your request should be submitted at least three days in advance of the laboratory period you wish to attend. You should use the following subject line in your e-mail: <Chem XXXX> - request for lab re-schedule, <your name> 17 UOIT Chem. 1800U, S11; Introduction - 8.13 If your request is approved, the senior laboratory instructor will sign and return the form to you by e-mail. Approval is solely at the discretion of the senior laboratory instructor and not all requests will necessarily be approved. If your request is approved, you must present the signed form to the teaching assistant when you arrive at your re-scheduled laboratory; it is your permission to attend the period. At the end of the period have the teaching assistant sign the form (and any relevant data sheets, laboratory notebooks, etc.). When you submit your laboratory report to your regular teaching assistant, attach the completed form (with all relevant signatures). Failure to do so will result in a grade of zero for the laboratory report. The report will be due at your next regularly scheduled laboratory period. Any laboratory reports that are submitted late as a result of missing a laboratory period will have marks deducted at the usual rate for late submission of reports. If it is not possible to re-schedule the laboratory, you may be excused from the laboratory if you have an acceptable reason for missing the laboratory. Acceptable reasons include: religious observance, medical reasons or bereavement. To be excused on these grounds, documentation must be submitted to the senior laboratory instructor. You may be excused for other reasons at the discretion of the senior laboratory instructor who may then require other documentation or proof for your absence. Other remedies such as quizzes or assignments in lieu of laboratory reports may also be used (but not necessarily). 18 UOIT Chem. 1800U, S11; Introduction - 8.13 WEIGHING TECHNIQUES Good weighing technique is vital to success in all aspects of chemistry and the measurement of mass is probably the most fundamental measurement in chemistry. Because of advances in technology mass can be one of the most accurately known quantities in any experiment. For greatest efficiency the proper balance must be chosen and the proper procedure must be followed. In chemistry two types of balances (along with specialized variations) are commonly used: the open pan balance and the analytical balance. Open Pan Balance Open pan balances are used when i) a mass with less accuracy is required; ii) the mass is heavier than can be accommodated on the analytical balance; iii) an approximate mass of a substance must be transferred from one container to another. Operating a open pan balance is very simple. Press the “tare” button to re-zero the balance, place the object on the balance pan and record the mass. To measure an approximate mass of substance into a receiving vessel place the receiving vessel on the balance pan and press the “tare” button. This will re-zero the balance. Essentially, the mass of the receiving vessel has been subtracted. As a substance is added to the receiving vessel, the balance displays the mass of added material. This avoids transferring chemicals on the analytical balance. 19 UOIT Chem. 1800U, S11; Introduction - 8.13 The Analytical Balance The Mettler-Toledo analytical balance is designed to weigh relatively small masses (less than 120 g) with very great accuracy (to ±0.0001 g). These balances are remarkably easy to use but are very delicate and must be handled with great care. 1. To ensure the greatest accuracy the balance should be level. This can be confirmed by observing the “level indicator” on the balance (which functions exactly like a carpenter’s spirit level). The bubble should be exactly in the middle of the circle. If not, consult the teaching assistant. Do not move the balances because this will alter the level. 2. Ensure all the doors are closed. If it is not on already, turn on the balance. The balance will run through a series of checks and eventually should display 0.0000 g. 3. Press the “tare” button. This will set the balance to zero. 4. Gently slide open one of the doors and place the item to be weighed on the pan. It must not touch the sides of the balance or the outer ring of the pan. 5. Close all the doors of the balance. Air currents affect the measured mass. 6. Wait until the balance indicates the reading has stabilized and then record the mass. 7. Remove the object and close all the doors. 8. Chemicals must NEVER be placed on the balance pan. They must be held in or on some suitable container. Preferably, it is a closed container. 20 UOIT Chem. 1800U, S11; Introduction - 8.13 9. If chemicals are spilled in or on the balance, they must be cleaned up immediately. Consult a teaching assistant. 10. Liquids and solutions should generally not be weighed on an analytical balance. If they are weighed on the balance, they MUST be in closed containers. 11. Chemicals must not be transferred to a weighing container on the analytical balance. The transfer should be done using a open pan balance. 12. Objects which are hot should not be weighed on the analytical balance. The air currents caused by the hot object will cause erroneous readings of the mass. 13. Make sure no one is leaning on the balance table when the measurement is being made as this will adversely affect the measurement. 14. When using the balance for a sequence of measurements on the same object (as, for example, when performing a “weight by difference”), it is good practice to use the same balance. This will minimize error because any (constant) error in the balance will be cancelled or partially cancelled in the difference. 15. The analytical balance is a very accurate and sensitive instrument. Handling glassware leaves fingerprints and the analytical balance can detect this difference in mass. For greatest accuracy, the weighing vial should be handled with tongs, while wearing gloves or with a piece of paper towel wrapped around the vial. 21 UOIT Chem. 1800U, S11; Introduction - 8.13 Weight by Difference This technique allows very accurate measurement of a relatively small mass transferred to a receiving vessel (e.g., a flask or a beaker). Usually an analytical balance is used, but a open pan balance may also be used if the masses involved are comparatively large. 1. Using a open pan balance weigh the (approximate) amount of the substance into a weighing vial. A small beaker may also be used, but a weighing vial is preferable. 2. Weigh the vial on the analytical balance and record the mass. 3. Pour the contents of vial into the receiving vessel. 4. Re-weigh the now (mostly) empty vial on the same analytical balance (see above for the reason why) and record the mass. The mass transferred to the receiving vessel is simply the difference between the two masses. It does not matter if the weighing vial is completely emptied into the receiving vessel. However, none of the substance can be spilled; all of it must be in either the weighing vial or in the receiving vessel. Weight by difference can also be done in “reverse”. That is, an empty weighing vial can be weighed and its mass recorded. The weighing vial is then filled and the mass of the filled vial is recorded. The mass of the substance in the vial is the difference of the two masses. An accurate weight by difference is best achieved by ensuring the mass of the vial (or beaker) is not too large compared to the mass of the substance in the vial (or beaker). In other words the difference between the two masses should not be small compared to the two masses. 22 UOIT Chem. 1800U, S11; Introduction - 8.13 VOLUMETRIC TECHNIQUES To properly use volumetric glassware it is important to understand how to read a “meniscus”. When most liquids (including water) are confined to a vessel, they will form a concave surface; the liquid will be higher at the edges than at the centre. This curvature is called the “meniscus”. It is caused by the interaction of the molecules of the liquid with the molecules of the vessel and is more pronounced for narrower vessels. The proper technique is to read the meniscus at its lowest point. To avoid parallax error the eye should be at the same level as the bottom of the meniscus. In some unusual cases reading the top of the meniscus is required. The bottom of the meniscus can sometimes be more easily read by placing a white card with a dark stripe behind the glass tube with the top edge of the dark stripe level with the bottom edge of the meniscus. This makes the bottom of the meniscus sharper and easier to see. Transfer Pipette The transfer pipette is one of the fundamental tools of analytical chemistry. They are accurately calibrated “to deliver” the prescribed volume to the receiving vessel. When used properly the error in volume transferred is probably less than 0.2%. The trick is to use it properly! For beginners this may take a little bit of practice. To prepare a pipette for use it should be rinsed with tap water, then with deionized water, and finally with the solution that is to be transferred. The rinsings should be performed three times. Unless absolutely necessary, soap should not be used when cleaning a pipette. Soap can be difficult to rinse completely from a pipette and may leave a contaminating film. 23 UOIT Chem. 1800U, S11; Introduction - 8.13 Squeeze the pipette bulb and place it over the thick end of the pipette. Place the thin end (the tip) in the solution. Slowly release your grip on the bulb to draw liquid into the pipette. Fill the pipette about half full. Remove the bulb and quickly place your finger over the end. Hold the pipette nearly horizontal (but do not hold it by the tip) and rotate it so the liquid inside coats the entire inner surface. Drain the rinsings into a waste beaker. Repeat the rinsings twice more. A clean pipette should drain smoothly and leave no droplets on the inner surface of the pipette. If droplets do form, the pipette needs more “vigorous” cleaning. Consult an instructor. Once the pipette has been properly rinsed it is ready to be used to transfer a very precise volume. Use the bulb (NEVER use your mouth!!) to draw the liquid into the pipette until the liquid level is well above the graduation mark. For large volume pipettes you may have to squeeze and use the bulb a second (or third!) time. If so, when you remove the bulb quickly place your first finger over the top of the pipette. During the swap you may find it helpful to gently rest the bottom of the pipette on the bottom of the beaker. Once the level of the liquid is above the graduation mark, remove the bulb and quickly cover the end of the pipette with your first finger. Do NOT use your thumb!! While filling the pipette, do not lift the bottom of the pipette above the level of the liquid. If you do, liquid will squirt into the bulb. This is a very bad thing! It can lead to cross contamination of solutions. If solution has entered the bulb, remove and clean the taper and remove and liquid in the bulb. Rinse the inside of the bulb a few times with deionized water and shake out the excess water. Allow the bulb to dry before re-using. Lift the tip of the pipette out of the solution. Using a Kimwipe or piece of paper towel, wipe off the tip. While touching the tip of the pipette to the side of the beaker slowly release the pressure of your finger on the top of the pipette to lower the level of the liquid. Keep the pipette vertical while doing this. Keep your eye level with the graduation mark. Lower the level of the liquid in the pipette until the bottom of the meniscus is exactly at the mark. 24 UOIT Chem. 1800U, S11; Introduction - 8.13 To transfer to the receiving flask vessel, place the tip of pipette against the side of the receiving vessel (again, keep the pipette vertical). Remove your finger allow the pipette to drain. Once the pipette is empty wait ~10 s (to ensure complete and consistent draining) and touch the tip to the side of the receiving flask. Do NOT blow out the small amount of liquid in the tip - the pipette is calibrated for the small amount that remains in the tip. Mohr Pipette Unlike a transfer pipette (which delivers a fixed volume) a Mohr pipette can be used to deliver variable volumes. The level of accuracy of a Mohr pipette is slightly lower than for a transfer pipette. The technique for using a Mohr pipette is very similar to using a transfer pipette. Two things must be kept in mind. First, for a Mohr pipette the volume delivered is determined by the difference between graduations. So, if the meniscus is lowered from the graduation at 1.00 mL to the graduation at 4.20 mL, the volume delivered to the receiving vessel is 3.20 mL. Second, a Mohr pipette is NOT calibrated for the liquid that remains in the tip. If you fill a Mohr pipette and then drain it, you will transfer more solution than you think! Serological Pipette These types of pipettes are not used extensively in chemistry, but do find wide application in biology and biochemistry. Like a Mohr pipette, they are capable of delivering a variable volume. Unlike a Mohr pipette, they do not operate “by difference”. If filled to the 7.00 mL mark, for example, and drained, the pipette delivers 7.00 mL to the receiving vessel. Some serological pipettes require “blow out” of the liquid in the tip; some do not. If you use a serological pipette, you must check to see which type you are using. 25 UOIT Chem. 1800U, S11; Introduction - 8.13 Burette A burette is a volumetric device used primarily for titration. It allows the variable, but accurately measured, volumes to be added to a receiving vessel. The volume delivered is determined from the difference between a final and an initial volume. A 50 mL burette is typically graduated in 0.1 mL intervals; the volume should be recorded to 0.01 mL by estimating between graduation marks. A burette is cleaned in a similar way to a pipette. It should be rinsed several times with small portions of tap water, deionized water and the solution that will be used to fill the burette. As with the pipette hold the burette horizontally and rotate it so that the liquid covers the entire inner surface of the burette. Drain the rinsings through the stopcock into a waste beaker. Repeat the rinsings at least twice more. Using a funnel and a beaker fill the burette with solution to just above the ‘0' graduation. Remove the funnel and open the stopcock and drain some of solution through the burette tip into a waste beaker. Ensure there are no bubbles in the tip. Touch the tip of the burette to the side of the waste beaker to remove the last drop. The level of the liquid must be below ‘0'! Record the “initial volume”. When reading the volume, the burette must be vertical and the meniscus should be at eye level. Read the volume at the bottom of the meniscus. You should be able to estimate the volume to 1/10 of a gradation. On a burette the volume increases as you read down the burette. The burette should be set up with the stopcock facing to the right (the directions here are for a right handed person; “lefties” should reverse all directions). The stopcock is controlled with the left hand - the palm of the hand is behind the burette and the stopcock is controlled by the thumb and first finger. Initially, this will feel quite awkward but with practice will become more comfortable. The right hand is used to swirl the receiving vessel (normally, a flask). The tip of the burette should be below the rim of the receiving vessel. 26 UOIT Chem. 1800U, S11; Introduction - 8.13 Initially, the titrant (the solution in the burette) may be added quickly. As the end-point (the point at which the reaction is judged to be “complete” and the addition of titrant is stopped) is approached, the additions should be slowed down. Very close to the end-point titrant should be added drop-wise (or even ½ drop-wise!). This can be done by allowing a drop to form slowly on the tip of the burette and then using a wash bottle to wash the drop into the solution. The wash bottle should be used to rinse down the sides of the flask to ensure all of the titrant run out of the burette has reacted with the solution in the flask. Once the end-point is reached wait a few seconds for the burette to drain properly and then read the final volume. Read the meniscus in exactly the same way as when the initial volume was read. To clean the burette after the experiment is finished, drain out any remaining solution into a waste beaker. Rinse several times with tap water and then with deionized water. Store the burettes with the stopcocks open. Graduated Cylinder A graduated cylinder measures volumes with moderate accuracy; it is better than a beaker, but not as accurate as a pipette. Care should be taken to note whether the cylinder is calibrated “to deliver” (TD) or “to contain” (TC). 27 UOIT Chem. 1800U, S11; Introduction - 8.13 Volumetric Flask A volumetric flask is used to prepare solutions of very accurately known concentration. These solutions can be made by dissolving solids (see, for example, the volumetric analysis lab) or by diluting other solutions. Prior to preparing a solution the volumetric flask should be rinsed several times with deionized water. Volumetric flasks do NOT need to be dried prior to use. In fact, it is better not to attempt to dry them; many volumetric flasks are broken by attempts to unnecessarily dry them. Transfer the solid (or solution) to the volumetric flask using a funnel. Ensure transfer is complete by rinsing the original vessel thoroughly and transferring the rinsings to the volumetric flask. Rinse the funnel. Be careful not to use too much solvent when performing the rinsings. Fill the flask ~3/4 full and swirl the mixture to ensure good mixing. Rinse the neck of the flask with deionized water. Continue filling the flask until the meniscus is ~1 cm below the graduation. Complete the filling by adding water drop-wise until the bottom of the meniscus is exactly on the graduation mark. Be very careful. If the meniscus is above the graduation (by even a little bit), you must start again. Once the meniscus is on the graduation, stopper the flask and invert it several times (10 - 15) to ensure good mixing. Volumetric flasks are calibrated “to contain” (TC) their designated volume at their designated temperature. If you pour out the solution, you will not get quite the same volume. Never put hot (or cold) solutions in a volumetric flask. Never expose the volumetric flask to extremes of temperature. The resulting expansion or contraction of the glass will ruin the calibration of the volumetric flask. 28 UOIT Chem. 1800U, S11; Introduction - 8.13 OPERATION OF THE BUNSEN BURNER A Bunsen burner is an excellent method for heating materials rapidly and efficiently. However, it is an open flame and this presents risks. A Bunsen burner should not be used to heat inflammable material, especially highly volatile solvents such as ethers and alcohols. The operator must be careful to keep long hair, clothing, and other combustible material away from the flame. A lit Bunsen burner must never be left unattended! Lighting a Bunsen burner is a fairly straightforward operation. First, make sure the air flow is at a minimum. Turn the “barrel” of the burner clockwise until the air inlet ports at the bottom of the barrel are completely closed. Turn slightly anti-clockwise to open the ports slightly. If there is too much air flow, the flame of the burner will be blown out like a candle. Adjust the air flow on a burner by grasping the barrel at the very bottom. Close the needle valve for the gas (underneath the barrel) by turning it anti-clockwise (if you are looking down on the burner from above). Open the gas tap fully (the handle of the tap will be straight in line with the nozzle). Open the gas needle valve slightly (turn clockwise) to allow the flow of some gas. Use the gas lighter to create a spark over the top of the barrel. The burner should light easily and produce a bright yellow flame. If the burner does not light easily (after 2 or 3 attempts with the lighter) turn off the gas at the tap, wait a few minutes for the gas to dissipate and then try to light the burner again. Once the burner is lit, you can control the height of the flame by adjusting the gas flow with the needle valve (turning it clockwise to increase the flow) at the bottom of the barrel. Once the gas flow is set, adjust the air flow by turning the barrel anti-clockwise. Remember, only touch the barrel at the very bottom; the upper parts of the barrel get very hot! Increase the air flow until the flame is pale blue with a well defined, blue, inner core. The hottest part of the flame is just above the inner blue core. If the air flow is too low the flame will be yellow and produce soot. If the air flow is too high, the flame may blow itself out. 29 UOIT Chem. 1800U, S11; Introduction - 8.13 To turn off the Bunsen burner, reduce the air flow until the flame is yellow. Next, close the gas needle valve and finally turn off the gas at the tap. As an aside, the burners used in these laboratories are more properly called “Tirrell” burners, not “Bunsen” burners. “Bunsen” burners use a slightly different method for controlling air flow. However, throughout this manual all burners will be referred to as “Bunsen” burners. 30 UOIT Chem. 1800U, S11; Introduction - 8.13 SUCTION FILTRATION Suction filtration is an efficient method to separate a precipitate from a solution (the “supernatant”). It can be used when either the precipitate or the supernatant is required. The technique employs a partial vacuum in a flask to draw the supernatant through a filtering device. Suction filtration is used most often to quantitatively separate the precipitate from the solution. When using a suction flask, it is an excellent idea to stabilize it with a ring clamp or an adjustable clamp. Otherwise there is considerable risk it will topple over. Always inspect the suction flask for damage or cracks before use. A damaged flask can implode violently when placed under vacuum and cause serious injury. For the recovery of small amounts of product, a Hirsch funnel or a sintered glass crucible can be used and for larger quantities a Büchner funnel is used. The sintered glass crucible requires no filter paper and is often used when the precipitate must be dried by heating in an oven. For the Hirsch and Büchner filters, a one hole stopper is put on the top of the filter flask and the appropriate funnel inserted into the hole. Sintered glass crucibles use special holders. A thick walled rubber tube connects the side arm of the suction flask to an aspirator which generates the vacuum in the flask. Place a filter paper sufficiently large to cover all the holes in the funnel and moisten it with a small amount of the solvent (from which the crystals have precipitated) to ensure the filter paper adheres well to the bottom of the funnel. Turn on the aspirator. This will create a moderate vacuum in the flask that will draw the supernatant solution through the filter paper. To recover the precipitate, decant the supernatant solution onto the filter paper and allow it to be drawn through the filter paper. Next, remove as much as possible of the solid from the beaker 31 UOIT Chem. 1800U, S11; Introduction - 8.13 onto the filter paper by scraping it out with a “rubber policemen”. Small portions of cold solvent can be used to wash out the crystals. Do not use too much solvent or the crystals may re-dissolve. Get the crystals as dry as possible with the suction filtration. The crystals are then washed with small amounts of cold solvent. Break the vacuum by lifting the funnel out of the hole (or by removing the tube from the side arm of the flask). Wash the crystals with the small amount of cold solvent and then re-establish the vacuum. Re-apply the suction until the crystals are once again dry. To facilitate better drying of the crystals, they can be pressed with the blunt end of a scoopula to squeeze out any solvent. Large “clumps” of crystals can also be broken apart (take care not to tear the filter paper). Repeat these washings 2 - 3 times, breaking the suction before each washing. Several washings with small amounts of solvent are better than one washing with a large amount of solvent. Once the crystals are dry, break the suction and then turn off the aspirator. It is important to break the suction first. Otherwise, water from the aspirator line may be sucked back into the suction flask thereby contaminating the filtrate. Carefully remove the precipitate from the filter paper onto a watch glass or into a sample vial. Be careful not to tear the filter paper. When the supernatant is to be kept, it is a good idea to employ a trap between the aspirator and the suction flask. The trap is a second flask with a two hole stopper and two glass tubes in the holes. One glass tube is connected to the aspirator and the other is connected to the side arm of the suction flask. Any water sucked in from the aspirator will be caught in the trap and not contaminate the filtrate. The trap must be able to withstand a vacuum - an ordinary Erlenmeyer flask will not do. 32 UOIT Chem. 1800U, S11; Introduction - 8.13 SIGNIFICANT FIGURES No measurement in science can be considered absolutely exact; that is, without error. All measurements will include some uncertainty. Significant figures are used to reflect the degree of uncertainty in a measurement. In general, more significant figures implies greater confidence in the value. Proper understanding and use of significant figures is an essential aspect of scientific communication. As an example, consider the measurement of a piece of string with a ruler. Let us assume the smallest division on the ruler is 0.1 cm. The end of the string will quite likely lie between two divisions of the ruler (say, between 11.5 and 11.6 cm). Hence, you must estimate another digit to get a better idea of the length of the string. One person might estimate the length as 11.52 cm while another might estimate the length as 11.53 cm. It is commonly accepted that the last digit reported is estimated. However, this last digit is still recorded as a “significant figure”. The number of significant figures allows scientists to distinguish between more and less precise measurements. The mass of a coin measured on an analytical balance (which measured to 0.0001 g) might be recorded as 7.0164 g (five significant figures), but if the same coin were weighed on a open pan balance (which measured to 0.01 g) the recorded mass would be 7.02 g (three significant figures). It would be incorrect to record the mass from the open pan balance as 7.020 g because this implies, falsely, that the open pan balance weighs to 0.001 g. Rules for Counting Significant Figures 1. All non-zero figures are significant. 33 UOIT Chem. 1800U, S11; Introduction - 8.13 11.82: 3.75: 2. 4 significant figures 3 significant figures All zeros between non-zero figures are significant 410.58: 256.032: 3. 5 significant figures 6 significant figures If a decimal point is present, all zeros to the right of a non-zero figure are significant 211.00: 5 significant figures 40.0: 3 significant figures 2500.: 4 significant figures If a decimal point is absent, the “significance” of the zeros is ambiguous. For a number such as 101 000 the number of significant figures is unclear. Such numbers are usually written in scientific notation so that the number of significant figures can be assessed. Thus, the above number might be written as: 1.01 x 105: 1.010 x 105: 4. 3 significant figures 4 significant figures Zeros to the left of a non-zero figure (but not between significant figures) are not significant. 0.0806: 3 significant figures 0.0007: 1 significant figure 34 UOIT Chem. 1800U, S11; Introduction - 8.13 Significant Figures in Calculations Rules for Rounding off Numbers When rounding off numbers follow these rules: 1. If the digit following the last digit to be kept is greater than 5 increase the last digit kept by one. 2. If the digit following the last digit to be kept is less than 5, leave the last digit kept unchanged. 3. If the digit following the last digit to be kept is equal to 5 and any of the digits following the 5 is greater than zero, increase the last digit kept by one. 4. If the digit following the last digit to be kept is equal to 5 and all the digits following the five are zero, round the last digit to be retained to the nearest even number. Multiplication and Division The final result should be reported with the same number of significant figures as the number with the fewest significant figures used in the operation. Examples: 2.12 3 significant figures x 3.025 4 significant figures 6.413 35 UOIT Chem. 1800U, S11; Introduction - 8.13 Reported as: 6.41 (3 significant figures) 12.053 5 significant figures ÷ 6.2 2 significant figures 1.94403 Reported as: 1.9 (2 significant figures) Addition and Subtraction The number of figures to the right of the decimal point in the final result must equal the number of digits to the right of the decimal in the number with the fewest digits to the right of the decimal. For example: 8.102 3 digits to the right of the decimal + 10.11 2 digits to the right of the decimal + 111.1 1 digits to the right of the decimal 229.312 Reported as: 229.3 (1 digit to the right of the decimal place; four significant figures) Be careful when adding numbers written in scientific notation. When adding or subtracting, ensure all the numbers are written with a common exponent before adding them. 8.139 x 10-5 + 2.16 x 10-4 0.08139 x 10-3 Y 36 + 0.216 x 10-3 UOIT Chem. 1800U, S11; Introduction - 8.13 + 1.218 x 10-2 + 12.18 x 10-3 12.47739 x 10-3 Reported as: 12.48 x 10-3 or 1.248 x 10-2 Logarithms and Exponentials When taking the logarithm of a number, the number of decimal places in the result equals the number of significant figures in the original number. For example: log10 (1.35 x 10-2) = 1.86967 Reported as: 1.870 (the original number has 3 significant figures and so this number is reported with 3 decimal places). ln (2.5) = 0.91629 Reported as: 0.92 Only the digits after the decimal point are considered significant. The numbers to the left of the decimal point locate the decimal point. For raising numbers to powers the rules are “reversed”. The number of significant figures in the result must equal the number of significant figures after the decimal point in the original number. e3.140 = 23.10387 Reported as: 23.1 (3 significant figures) 37 UOIT Chem. 1800U, S11; Introduction - 8.13 100.31 = 2.0417 Reported as: 2.0 (2 significant figures) “Exact” Numbers: Some numbers in science are considered to be “exact”. Exact numbers can be considered to have an “infinite” number of significant figures. A number is “exact” when the objects can be counted individually. For example, 112 jellybeans in a jar, 21 students in a laboratory, 43 apples in a basket. The coefficients and subscripts in chemical equations are “exact”: Zn(s) + 2HCl(aq) 6 ZnCl (aq) + H (g) 2 2 All the numbers in this equation are considered exact. Numbers that are defined are considered exact. This occurs most commonly for conversion from one unit to another (for example from inches to centimetres or from kilograms to grams): 1 inch = 2.54 cm, exactly 1 kg = 1000 g, exactly So, to convert 12.12 inches (four significant figures) to centimetres: 12.12 inches x 2.54 cm / inch = 30.7848 cm Reported as: 30.78 cm (four significant figures) In a sequence of calculations, at each step you should report the value with the correct number of significant figures. However, to avoid “rounding errors” (which may become considerable if there are many steps), all the figures provided by the calculator should be carried through the calculation until the end. At the end, the number should be rounded off to give the correct number of significant figures. 38 UOIT Chem. 1800U, S11; Introduction - 8.13 LABORATORY REPORT RECEIPT SHEET This sheet can be used as proof that you have submitted a laboratory report. Whenever you hand in a report, have an instructor sign and date this sheet. Keep the sheet in safe place (e.g., staple it into your lab notebook). If an instructor loses a report, this sheet will act as proof that you handed in the report. Name (print): __________________________________________________ Student Number: __________________________________________________ Report Instructor Signature Date 1 2 3 4 5 6 Check-out At the end of the semester you must “check-out” your locker equipment. Ensure that all of the equipment is present, clean, and in good repair. Have the teaching assistant check it. Failure to “check-out” may result in the withholding of your laboratory marks. 39 UOIT Chem. 1800U, S11; Introduction - 8.13 SAFETY ACKNOWLEDGEMENT Carefully read the following and print and sign your name on the form. You must sign this form and present it to your laboratory instructor at the beginning of your first laboratory period.. I have read the safety rules and good laboratory practices outlined at the beginning of this laboratory manual and agree to abide by these rules and practices. I acknowledge that failure to follow these rules may result in dismissal from the laboratory period, a mark of zero for the experiment and no opportunity to repeat the experiment. I accept that persistent failure to abide by the safety rules and good laboratory practices will result in dismissal from the laboratory portion of the course. I have read the guidelines on laboratory reports. I acknowledge that failure to abide by these instructions may lead to loss of marks. I have read the instructions on the use of laboratory equipment and agree to use the equipment in accordance with those instructions. I acknowledge that using the equipment in such a way that is dangerous or that is potentially damaging to the equipment may result in dismissal from the laboratory period, a mark of zero for the experiment and no opportunity to repeat the experiment. I accept that persistent abuse of equipment will result in dismissal from the laboratory portion of the course. Name: _______________________________________________________ Signature: _______________________________________________________ Student Number: _______________________________________________________ Date: _______________________________________________________ Laboratory Period: _______________________________________________________ 40 UOIT Chem. 1800U, S11; Introduction - 8.13 LOCKER EQUIPMENT LIST Item Number 50 mL beaker 2 150 mL beaker 2 250 mL beaker 2 400 mL beaker 1 600 mL beaker 1 50 mL Erlenmeyer flask 1 125 mL Erlenmeyer flask 2 250 mL Erlenmeyer flask 3 10 mL graduated cylinder 1 25 mL graduated cylinder 1 100 mL graduated cylinder 1 Gas lighter 1 Crucible tongs 1 Thermometer, alcohol 1 Tweezers 1 Funnel, short stem, plastic 2 W atchglass, 100 mm 2 Medicine dropper 2 Scoopula 1 Stirring rod 1 Stirring rod with rubber policeman 1 Test tube brush, large 1 Test tube bush, small 1 Test tube rack 1 Test tube holder 1 Test tube, 20 mm x 150 mm 12 W ash bottle, 500 mL 1 Plastic bottle, 500 mL 1 41 UOIT Chem. 1800U, S11; Introduction - 8.13 Beakers Erlenmeyer Flasks Graduated Cylinders Gas Lighter Thermometer Tweezers Plastic Funnel Watchglass Medicine dropper Scoopula Stirring Rods Test Tube Brushes Test Tube Holder 20 x 150 mm Test Tubes and Rack Wash Bottle (l) and Plastic Bottle (r) Equipment photographs courtesy of Kaitlyn Yarrow 42 UOIT Chem. 1800, W12: Exp. 1-5.9 1. LABORATORY SAFETY AND ORIENTATION Objectives 1) To become familiar with the safety equipment and procedures in the laboratory; 2) to learn to recognize the glassware and equipment in the lockers; 3) to prepare a standard solution of sodium chloride. Introduction One of the primary roles of the teaching assistant is to ensure that all students are informed about the safety rules of the laboratory and to enforce those rules. Everyone’s safety and security in the laboratory depends on each student knowing and obeying the rules and working safely. At the beginning of this laboratory period the teaching assistant will spend 25 - 30 minutes explaining the safety rules of the laboratory, pointing out safety equipment, and describing procedures. This information is also presented in the introductory pages of this manual. You should thoroughly familiarize yourself with the safety rules and procedures of the laboratory. Be aware that failure to comply with the safety rules may result in your dismissal from the laboratory. At the beginning of the semester, the technicians ensure that all the equipment listed in the locker equipment list is in each locker and in good repair. After the safety talk, please check your locker. If you discover that anything is missing, ask the teaching assistant or the technician for a replacement. At the end of each laboratory period, make sure all the equipment that is supposed to be in the locker (no more, no less!) is there and that your bench top has been wiped clean with a damp sponge. The first-year chemistry technician will do random inspections of lockers and benches and if the workspace is dirty or equipment is missing, marks will be deducted from the last student 43 UOIT Chem. 1800, W12: Exp. 1-5.9 who used the locker (2.5 marks from the laboratory report). Preparing a solution with an accurate concentration is a fundamental skill in chemistry. Many solutions that need to have accurately known concentrations are made by dissolving a known mass of a solid (the solute) in a known volume of liquid (the solvent). These solutions (called primary standards) are then used to determine the concentrations of compounds in other solutions. Thus the primary standard solutions must be made with great care and great accuracy. The calculation of the concentration of the solid in the primary standard solution is straightforward. The number of moles of solid is given by: (1.1) where mcompound Mcompound = mass of the compound (in g) = molar mass of the compound (in g mol-1) The concentration (in mol L-1) is given by: (1.2) where Vsolution = volume of the solution (in L) In this experiment a solution of ~0.1 mol L-1 NaCl will be prepared. Sodium chloride is not commonly used as a primary standard, but the technique used here to prepare the sodium chloride solution is the same as would be used for making a primary standard. 44 UOIT Chem. 1800, W12: Exp. 1-5.9 Procedure Part I: Safety Orientation In the first part of the laboratory period the teaching assistant will explain the safety protocols and procedures that are outlined in the introduction of this manual. Once the teaching assistant has completed the introduction you should answer the following questions. Submit these answers on a separate piece of paper with your laboratory report. 1. Using the laboratory template provided, indicate the locations of: a. The fire extinguisher b. The eyewash station c. The first aid kit d. The fire blanket e. The spill kit 2. List five (5) restrictions and / or requirements for clothing in a chemistry laboratory. 3. What is the proper way to dispose of broken glass? 4. What steps should you take if the fire alarm sounds? 5. What are the appropriate treatments for i) cuts and ii) minor burns? Part II: Locker Equipment Check At the end of the introductory pages to the laboratory manual is a list of equipment that should be in each locker. Review this list and ensure that all the equipment that is on the list is in 45 UOIT Chem. 1800, W12: Exp. 1-5.9 your locker kit. If you do not know what a piece of equipment looks like, ask the teaching assistant. If a piece of equipment is missing, ask the teaching assistant or technician for a replacement. During the semester the technicians will perform random spot checks of the lockers to ensure that all the equipment that should be in the lockers has been returned. If equipment is missing or if there is extra equipment, you will have marks deducted from your report. Likewise, if your work area is not cleaned after the experiment, you will lose marks. Part III: Preparation of a Standard Solution of NaCl 1. Place a 50 mL beaker on the open pan balance and “tare” (zero) the balance. Add ~1.5 g of NaCl to the 50 mL beaker. 2. Weigh the 50 mL beaker + NaCl on the analytical balance. The analytical balance will give a mass to four decimal places. Write down all these numbers. 3. Pour the NaCl into a clean 250 mL beaker. Be careful not to spill any of the NaCl. It is more important not to spill any of the NaCl than to get all of it out of the 50 mL beaker. 4. Re-weigh the now (mostly) empty 50 mL beaker on the same analytical balance. Again, write down all the numbers. The difference between this number and the number recorded in step two is the mass of NaCl transferred. This is known as the “weight-by-difference” technique. 5. Add ~125 mL of water to the 250 mL beaker (this volume does not have to be measured accurately). 6. Using a glass stirring rod, carefully stir the solution (do not splatter the solution) until all the 46 UOIT Chem. 1800, W12: Exp. 1-5.9 NaCl dissolves. If, at any point, you remove the stirring rod from the beaker, rinse it with deionized water from a wash bottle so that the rinsings flow into the beaker. This ensures no NaCl is lost from the beaker. 7. Obtain a 250.0 mL volumetric flask and rinse it thoroughly three times with small portions of deionized water. 8. Now the quantitative transfer of the NaCl solution in the beaker must be performed. You need to transfer all the NaCl to the flask. 9. Place a funnel in the neck of the volumetric flask and, using the stirring rod as a guide, carefully pour the NaCl solution so that it runs down the rod through the funnel and into the flask. Do not place the stirring rod on the bench once you have completed this step. 10. Using a wash bottle carefully wash the inside walls of the beaker with deionized water. Using the stirring rod as you did in step 9 transfer the washings to the volumetric flask. Repeat the washings twice more. Do not use too much deionized water at this stage - the total volume added to the flask must be less that 250 mL! 11. Rinse the end of the stirring rod so that the rinsings go through the funnel into the flask. 12. Finally, rinse the inside surface of the funnel with deionized water. If steps 9 - 12 have been performed properly, all of the NaCl should be transferred to the flask. 13. Carefully fill the flask with deionized water so that the bottom of the meniscus is exactly on the mark. If this is done properly, there will be 250.0 mL of solution in the flask (within the systematic error in the flask). The bottom of the meniscus cannot be above nor below the mark. Be very careful as the solution nears the mark. The last few millilitres should be 47 UOIT Chem. 1800, W12: Exp. 1-5.9 added with a medicine dropper. If the flask is overfilled, the procedure must be started again. If you overfill the flask, empty it, rinse three times with deionized water and start again. 14. Once the solution is prepared, stopper it and have the teaching assistant check your work. The stopper needs to fit snugly but should not be put in with excessive force. To do so, risks injury or jamming the stopper in the neck of the flask. 15. Invert the flask 20 - 25 times (while holding the stopper in place with your hand) to ensure thorough mixing. 48 UOIT Chem. 1800, W12: Exp. 1-5.9 Experiment 1 - Lab Safety and Orientation Name: ____________________________ Student Number ____________________________ Day ____________________________ Time: ____________________________ TA Name ____________________________ 49 UOIT Chem. 1800, W12: Exp. 1-5.9 50 UOIT Chem. 1800, W12: Exp. 1-5.9 Results and Questions Part I - Safety Orientation On a separate piece of paper answer the questions posed in the procedure. Include the completed laboratory template showing the locations of the safety equipment. Part III - Preparation of Standard Solution mass of 50 mL beaker + NaCl / g ____________________________________ mass of “empty” 50 mL beaker / g ____________________________________ mass of NaCl used / g ____________________________________ Final volume of solution / L ____________________________________ Checked by the TA (TA initials) ____________________________________ Calculate the concentration of the NaCl solution (in mol L-1 and with the correct number of significant digits). 51 UOIT Chemistry 1800, W12; Exp. 2-9.13 2. ATOMIC EMISSION SPECTRA For this experiment, bring your laptop to the laboratory. You will need to construct graphs to analyze the data. Objectives 1) To observe the atomic emission spectra of two elements: mercury and hydrogen; 2) to calculate the Rydberg constant from the spectrum of the hydrogen atom; 3) to observe the spectra for several cations and note that cations give characteristic colours to flames; 4) to understand spectra are a valuable tool for identifying chemical species. Introduction In the late 19th century Johann Balmer observed that when electric current was passed through hydrogen gas, the gas emitted light. When the gas was viewed through a spectroscope, he observed discrete, bright lines against a dark background in the visible region of the electromagnetic spectrum. Discrete lines indicate that the hydrogen atom must exist in discrete energy levels. Otherwise, the emitted light would resemble a rainbow. He found he could relate the wavelengths (ë) of the lines to a constant, k, using the relationship: (2.1) where ‘n’ was an integer > 2. The series of lines that he observed was called the Balmer series. 52 UOIT Chemistry 1800, W12; Exp. 2-9.13 Subsequently, series of lines were found in the ultraviolet (the Lyman series) and in the infrared (the Paschen, Brackett and Pfund series). These extra lines allowed the generalization of equation (2.1): (2.2) where n2 > n1. RH is known as the Rydberg constant and in this equation has units of reciprocal wavelength (commonly expressed in cm-1). The Rydberg constant was a purely empirical value and its literature value is 1.0973 x105 cm-1. It had no theoretical basis until Bohr’s model of the hydrogen atom. In Bohr’s model of the atom the electron is restricted to specific “orbits” circling the nucleus. When it is in the orbit closest to the nucleus, the electron has its lowest energy. The lowest energy state (electron in the first orbit) is called the “ground state”. The electron in the lowest orbit can be raised (excited) to a higher state by the addition of energy. However, the energy must exactly match the energy difference between the initial and final states; the electron can only absorb energy in discrete (“quantized”) amounts. When the electron returns to the ground state (or any lower energy state) it will emit energy exactly equal to the energy difference between the upper and lower states. This is why the spectrum of the hydrogen gas appears as discrete lines. Bohr also postulated that electrons existed with discrete (or quantized) amounts of angular momentum as they circled the nucleus. Based on this assumption Bohr derived a theoretical expression for the energy levels of the electron in the hydrogen atom: (2.3) where: me = mass of an electron e = charge on an electron h = Planck’s constant 53 UOIT Chemistry 1800, W12; Exp. 2-9.13 g0 = permittivity of free space Z = atomic number n = integer The first four of these are constants and Z = 1 for hydrogen. The (1 / n2) factor means the electron is limited to discrete energy levels. If the constants are evaluated, then for hydrogen the energy difference between any two orbits (or energy states) is given by: (2.4) The negative sign is a convention used to indicate energy is lost (emitted) by the system when an electron moves to a lower energy state. When the Rydberg constant is converted to units of Joules, it is equal to 2.179 x 10-18 J. Bohr’s model successfully reproduced empirical observation, showed the validity of the quantum model and gave a theoretical basis to Rydberg’s constant. Bohr’s model, while a great step forward, is quite simple. It explains many of the observations made by physicists at the end of the 19th and beginning of the 20th centuries. While his model works well for the hydrogen atom, it is inadequate for multi-electron atoms. Bohr’s model has been replaced by the quantum mechanical model of the atom. Quantum mechanics does not regard electrons as moving in “orbits” around the nucleus. Electrons are viewed as being in “orbitals” which are mathematical constructions that describe the probability of finding an electron in a region of space surrounding the nucleus. Patterns of discrete lines are observed for many other elements and compounds when they have been “excited”. Atoms or molecules can become excited by several methods including electrically or thermally. The pattern of the discrete lines (the “spectrum”) is unique to each species and therefore can be used as a “fingerprint” for identifying species. This is the basis of branch of science known as “spectroscopy”. 54 UOIT Chemistry 1800, W12; Exp. 2-9.13 In this experiment you will use a simple handheld spectroscope to observe the emission lines of hydrogen gas. The spectroscope will be calibrated by measuring the well-known positions of lines from mercury gas. From the spectrum of the hydrogen gas the Rydberg constant will be calculated. The emission of several cations will also be observed. A spectroscope is a simple device that resolves (or splits) light into its component wavelengths using a reflection grating. A reflection grating separates the light into its component wavelengths based on the phenomena of constructive and destructive interference. At a particular angle of reflection from the grating only one wavelength will constructively interfere and be seen; other wavelengths destructively interfere and are not seen. Different wavelengths constructively interfere at different angles, so the light is dispersed into its component wavelengths, creating the spectrum. A similar effect can be observed by holding a compact disc up to the light. As the angle of the disc is changed, different colours (or wavelengths) are observed reflected from the surface. The discrete nature of the spectrum of hydrogen (and other elements), the photoelectric effect and the “ultraviolet catastrophe” were three natural phenomena that could not be explained by classical physics. The stage was set for the birth of quantum mechanics which, in turn, revolutionized physics and our understanding of the natural world. Apparatus and Materials 1) Spectroscope; 2) Nichrome wire; 3) Mercury discharge lamp and power supply (set up in the fume hood); 4) Hydrogen discharge lamp and power supply; 5) Five small, clean test tubes for collecting metal chloride solutions; 6) 0.1 mol L-1 solutions of NaCl, LiCl, KCl, CaCl2 and an unknown metal chloride; 7) Tungsten filament (incandescent) desk lamp. 55 UOIT Chemistry 1800, W12; Exp. 2-9.13 Procedure You will perform this experiment with a partner. Part I: Calibration of the Spectroscope 1. Do NOT raise the fume hood shield. The mercury lamp emits ultraviolet radiation which may be damaging to the eyes. The shield absorbs this radiation. 2. Holding the spectroscope level, point the slit of the spectroscope at the mercury lamp. Try to keep the slit 30 - 40 cm away from the lamp. If you are too close, the lines will “wander” across the scale; if you are too far away, the lines may be too faint to observe. 3. Record the scale position of all the lines you see. Four lines should be visible, although the violet line at 404.7 nm may be difficult to observe. 4. Create a calibration plot for your spectroscope. Plot the scale reading on the abscissa (x-axis) and wavelength on the ordinate (y-axis). Be sure to follow the graphing guidelines given at the beginning of the lab manual. You will use this graph to convert scale readings to wavelengths. Fit a linear function to the data using the regression wizard in Microsoft Excel. Part II: Emission Spectrum of Hydrogen 1. Point the slit of the spectroscope at the hydrogen lamp. 2. Record the scale positions and colours of all the lines you can see. You may see a faint yellow line in the spectrum - this line is a “ghost” that arises from imperfections in the spectroscope. It does not arise from the hydrogen gas. The hydrogen gas should produce 56 UOIT Chemistry 1800, W12; Exp. 2-9.13 lines that are red, green, blue and purple. The latter two may be close together and difficult to resolve. 3. Using the calibration curve, convert the scale readings to wavelengths. Record the data in the table below. Part III: Observation of the Spectra of Several Cations Obtain a small portion of each salt in small, clean, labelled test tubes. 1. Form a circle or loop of wire at one end of the nichrome wire. 2. Dip the loop into one of the salts (start with KCl). Next, hold the loop in the flame. Observe and record the colour of the flame. 3. Using the spectroscope observe the light emitted by the flame. Record the colour of the flame and the scale positions for each line that you can see. The colour imparted to the flame by the salt lasts only a few seconds, so this may be quite difficult and may require several attempts. If you leave the wire in the flame too long it will lend a yellow colour to the flame that comes from the nichrome wire. If you are unable to record the positions after four or five attempts, proceed to the next salt. 4. If you have recorded scale positions, convert them to wavelengths using your calibration curve and regression line. Convert the scale readings to wavelengths using your calibration curve. Record these wavelengths below. Use the regression line calculated by Excel to determine the wavelengths. 5. Rinse the nichrome wire with some deionized water and repeat the procedure for the next 57 UOIT Chemistry 1800, W12; Exp. 2-9.13 three salts. Do the salts in the following order: KCl, LiCl, CaCl2, NaCl. 6. Obtain an unknown solution and record the unknown number. Perform a flame test with the unknown and identify the cation based on your observations. Part IV: Comparison of Fluorescent and Incandescent Lamps 1. Use the spectroscope to observe light emitted by both a fluorescent and an incandescent lamp. Record what you observe. If you observe any discrete lines, record the scale reading and calculate the wavelength using your calibration graph. Record the wavelengths. 58 UOIT Chemistry 1800, W12; Exp. 2-9.13 Experiment 2 - Atomic Spectra Name: ____________________________ Student Number ____________________________ Partner ____________________________ Day ____________________________ Time: ____________________________ TA Name ____________________________ 59 UOIT Chemistry 1800, W12; Exp. 2-9.13 DATA SHEET Part I: Calibration of Spectroscope Colour of Line Wavelength / nm yellow 435.8 violet Wavelength / nm 546.1 blue Scale Reading 579.0 green Scale position 404.7 Part II: Hydrogen Spectrum Colour (from graph) 60 UOIT Chemistry 1800, W12; Exp. 2-9.13 Part III: Spectra of Cations Solution Flame Colour Scale Positions Wavelengths / nm (from graph) KCl LiCl CaCl2 NaCl unknown Part IV: Fluorescent and Incandescent lamps Describe what you see for: I) Fluorescent lamp ii) Incandescent lamp 61 UOIT Chemistry 1800, W12; Exp. 2-9.13 Results, Discussion and Questions 1. For each line in the hydrogen spectrum deduce n2 (the upper state) and n1 (the lower state). Hint: the lines are in the visible regions; to which series must they belong (consult your textbook)? 2. Spectroscopists define a quantity called the “wavenumber” (= 1 / ë) and give it the symbol G The units are cm-1. Photon energies are linearly dependent on G í. í: (2.5) The Rydberg equation can be re-written as: (2.6) For each line of the hydrogen spectrum calculate the Rydberg constant in wavenumbers (in cm-1) using equation (2.6). You must first convert the wavelength of each transition from nanometres to centimetres. Show a sample calculation. 3. Re-arrange equation (2.6) into the form of a line (Hint: (1 / n12) is constant. Put x = (1 / n22) and re-arrange the equation to the form y = mx + b). Plot your data and find the line of best fit (i.e., the regression line). From the equation for the line of best fit, calculate the Rydberg constant (this is likely how Rydberg would have determined this value). Compare the different values you calculate with the literature value for the Rydberg constant (1.0973 x105 cm-1). Show a sample calculation. 62 UOIT Chemistry 1800, W12; Exp. 2-9.13 Wavelength / nm n1 n2 Rydberg Constant / cm-1 Percentage Error Rydberg Constant from Graph 4. If you could record the absorption spectrum of hydrogen gas, what would it look like? Justify your answer. 5. In Part III what is your unknown cation? Justify your answer. 6. When you observe the spectra of the metal salts, how do you know the emission comes from the cation and not the anion? Suggest an experiment to verify this. 7. What is the vapour (or gas) present in fluorescent lights? What experimental evidence do you have to support your assertion? 63 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 3. MEASUREMENT OF THE IDEAL GAS CONSTANT, ‘R’ Objectives 1) To apply Dalton’s Law of Partial Pressures to a mixture of gases (water vapour and hydrogen); 2) To measure the ideal gas constant and compare the measured value with the accepted value (8.314 L kPa mol-1 K-1. Introduction The behaviour of gases as the pressure, volume, temperature and amount of gas (i.e., the number of moles of gas) are changed was one of the earliest things to be studied systematically in science. Over 300 years ago Boyle observed that (to a very good approximation) the volume of a gas was inversely proportional to the pressure of the gas (at a constant temperature): (3.1) Charles expanded (no pun intended!) on this and reported that at constant pressure the volume of a gas was directly proportional to its temperature: (3.2) Charles’s results showed that when the V vs. T data was extrapolated to a volume of 0 L the corresponding temperature was ~ -2730C. This temperature was then defined as “absolute” zero and is the basis of the Kelvin scale for temperature. At roughly the same time Gay-Lussac showed (not surprisingly!) that the pressure of a gas was proportional to its temperature: (3.3) It is also fairly clear that the volume of a gas depends on how much of it (the number of moles) is 64 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 present: (3.4) (This was first stated in a slightly different way by Avogadro). These four laws can be combined into a single statement: (3.5) (3.6) and so: (3.7) where ‘b’ is a proportionality constant. The proportionality constant is more commonly called ‘R’, the gas constant and this equation is written in its more recognizable form as: (3.8) which is known as the ideal gas equation. It should be noted this is an “ideal” case. For real gases this equation is an approximation and better models are used for more advanced applications of physical chemistry. However, it is a very good approximation for many gases (including hydrogen) at moderate pressures. The failures of the ideal gas law will be examined in higher level physical chemistry courses. Dalton’s Law of Partial Pressures is also very important in the study of gases. This law states that each gas in a mixture exerts a pressure proportional to its concentration. The total pressure is the sum of the individual “partial” pressures: (3.9) (3.10) 65 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 (3.11) (3.12) In this experiment you will react an accurately known mass of magnesium (measured on the analytical balance) with an excess of hydrochloric acid according to: Mg(s) + 2HCl(aq) v MgCl (aq) + H (g) 2 2 (3.13) From the stoichiometry of the reaction the number of moles of hydrogen gas can be calculated (assuming all the magnesium reacts). The volume of gas will be measured using a “gas burette”. Now, if the temperature is known and the pressure of the hydrogen gas can be determined, the value of the gas constant ‘R’ can be calculated from the ideal gas equation: (3.14) The measurement of the pressure of the hydrogen is non-trivial but straightforward. Because the gas is collected over water, the gas collected in the burette is a mixture of hydrogen gas and water vapour. The total pressure of gases (Pgases) in the burette is given by Dalton’s Law: (3.15) (3.16) where PH2 is the pressure of the hydrogen gas and PH20 is the pressure of the water vapour (the vapour pressure). The vapour pressure of water as a function of temperature is well known and given in the table at the end of the experiment. How is the total pressure in the gas burette measured? Initially, the gas burette is filled with water. The open end (actually the end with the one hole stopper) of the burette is submerged in a 600 mL beaker filled with water. As the reaction proceeds the water is displaced by the generated 66 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 H2 gas. After the reaction between the magnesium and the hydrochloric acid is complete the burette should still be partially filled with water (more accurately, a dilute solution of HCl). If the pressure of the gases in the burette were equal to the atmospheric pressure, the level of the water in the burette would be the same as the liquid level in the 600 mL beaker. The water level in the burette will be higher than the level in the beaker so the pressure in the burette must be less than atmospheric pressure: (3.17) (3.18) Plevel difference is found by measuring the difference (in millimetres) between the water levels in the burette and in the beaker. The atmospheric pressure will be provided, so to calculate Pgases, the height of the column of water must be converted to pressure (in Pa) to find Plevel difference. The downward forces on the water column are the weight of the water plus the force exerted by the gases in the burette. The upward force on the water is the pressure exerted by the atmosphere (Pascal’s Principle). The column of water is not moving so the forces on the water must be balanced. Therefore: (3.19) (3.20) (3.21) h = height of the water column A = cross sectional area of the water column (hA = volume of water in the column) ñ = density of water g = acceleration due to gravity Dividing by A (and noting that P = F / A): 67 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 (3.22) (3.23) If Pgases = 0 (i.e., a vacuum at the top of the burette), Patmosphere = ñhg and the height of the water column would be given by: (3.24) If Patmosphere = 101 325 Pa (101.325 kPa), then h = 10 339 mm (assuming ñ(H2O) = 1000 kg m-3). Therefore, (3.25) From equations (3.23) and (3.24) (with Pgases = 0): (3.26) Dividing equation (3.26) by equation (3.25) gives: (3.27) (3.28) Substituting equation (3.28) into equation (3.23): (3.29) Therefore, if the height of water in the column (h) is measured in millimetres and Patmosphere is reported in Pascals, Pgases (in Pa) can be calculated. The atmospheric pressure (in mm Hg) will be 68 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 reported in the lab. Note: 1 mm Hg = 133.32 Pa. Apparatus 1) 600 mL beaker; 2) copper wire; 3) 00 one hole stopper; 4) gas burette; 5) retort stand; 6) burette clamp; 7) metric ruler; 8) magnesium ribbon; 9) 6 mol L-1 hydrochloric acid (HCl). Safety Magnesium 1. Magnesium is a combustible metal that will burn with an intensely white light that is damaging to the eyes. The chances of ignition in this laboratory are extremely remote. If magnesium does it ignite, the fire can be smothered with sand from the bucket in the laboratory. 6 M HCl 2. Spills should be treated with a small amount of sodium bicarbonate to neutralize the acid (the fizzing will stop). Wear gloves and wipe up the spill with a damp sponge or paper towel. Thoroughly rinse the sponge and wash hands with soap and water. 3. Skin exposure: immediately flush skin with cool running water. Continue for 20 - 30 minutes. Wash area thoroughly with soap and water. If irritation develops, seek medical attention. 69 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 4. Eye exposure: Immediately flush with running water for at least 30 minutes. Seek medical attention. Notify teaching assistant, technician or senior laboratory instructor. Waste 5. The waste from this experiment is fairly innocuous: a dilute solution of magnesium chloride in ~0.3 mol L-1 HCl. It can be disposed of in the sinks with plenty of water. Procedure Note: Chemistry 1010 students must submit the laboratory report for this experiment at the end of the laboratory period. Part I: Calibration of Gas Burette At the top of the gas burette is small volume (the “dead volume”) which is uncalibrated. In order to get an accurate value for the gas constant, this dead volume needs to be determined. 1. Put roughly 9 mL of water into a 10.0 mL graduated cylinder and record the volume in the cylinder. 2. Using a dropping pipette transfer water to the gas burette until the level of the water is at the first graduation on the burette. 3. Record the volume remaining in cylinder. The difference between the volume recorded in step 1 and this volume is the “dead” volume of the burette. 70 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 Part II Determination of the Ideal Gas Constant 1. Half fill the 600 mL beaker with deionized water. 2. Obtain a strip of magnesium ribbon roughly 2 - 3 cm long. Using a small beaker, weigh the magnesium ribbon using the analytical balance. 3. Form the magnesium strip into a coil. 4. Obtain a 25 cm length of copper wire and twist roughly 20 cm around the end of a stirring rod to make a coil. Leave the remaining 5 cm straight to use as a “handle”. 5. Put the magnesium in the copper coil so that the copper wire forms a “cage” around the magnesium. The magnesium must be held securely so that the magnesium does not come free. Do not flex the magnesium strip too much or it may break. 6. Thread the straight portion of the copper wire through the hole of the 00 stopper. 7. Rinse the gas burette with tap water followed by deionized water. Add the water to the burette using a beaker. 8. Carefully pour ~15 mL of 6 mol L-1 hydrochloric acid into the gas burette. 6 mol L-1 hydrochloric acid is corrosive. If you spill any on your skin, immediately rinse with plenty of cold water and wash your hands thoroughly with soap and water. 9. Fill the burette with deionized water. Be careful not to cause undue mixing of the acid and the water. This may be easier if the water is gently poured down the side of the burette. Be sure to completely fill the burette - there must not be any air remaining in the burette! 71 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 10. Put the stopper in the burette and position the magnesium so that it is about 4 cm into the burette; it must not block the hole. It will be held in place by the copper cage / wire. Again, make sure there is no air left in the burette. 11. Cover the stopper hole with your finger and invert the burette. Quickly submerge the stoppered end of the burette in the water in the 600 mL beaker and remove your finger. Do not remove your finger until the end of the burette is submerged. Clamp the burette in place. 12. The acid is more dense than the water and so quickly falls towards the bottom of the burette. Once the acid reaches the magnesium the reaction will begin and proceed rapidly, producing hydrogen gas. 13. Once the reaction appears to have finished, check to make sure there is no unreacted magnesium on the sides of the burette or trapped in the copper coil. 14. Free any hydrogen gas adhering to the copper coil or to the sides of the burette by gently tapping the burette with a pencil. 15. The reaction is exothermic. Wait about 5 minutes and then record the room temperature to the nearest 0.10C. It is reasonable to assume the temperature of the gases in the burette is equal to the room temperature. 16. Record the volume of gas in the burette to the nearest 0.01 mL. Add the “dead” volume to this volume to account for the uncalibrated volume at the top of the burette. 17. Finally, measure the difference in water level between the water in the beaker and the water in the burette. Measure this in millimetres. Note the atmospheric pressure, Patmosphere. It will be written on the board. 72 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 18. Drain the gas burette completely and rinse several times with tap water, followed by deionized water. Flush the contents of the beaker down the drain and clean the beaker with plenty of tap water followed by deionized water. 73 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 Experiment 3 - Measurement of R Name: ____________________________ Student Number ____________________________ Day ____________________________ Time: ____________________________ TA Name ____________________________ 74 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 Results, Discussion and Questions Initial volume in the 10.0 mL graduated cylinder _______________________ Final volume in the 10.0 mL graduated cylinder _______________________ “Dead” volume of the gas burette _______________________ mass of beaker + magnesium strip / g _______________________ mass of beaker / g _______________________ mass of magnesium strip / g _______________________ volume reading from the burette / mL _______________________ correction for dead volume / mL +_______________________ volume of gases in the burette / mL =_______________________ height of the water column / mm _______________________ Temperature / 0C _______________________ atmospheric pressure / mm Hg _______________________ atmospheric pressure / Pa _______________________ 75 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 Show all calculations: 1. From the measured height of the water column, calculate Pgases in kPa. (Equation (3.29)) 2. Calculate the pressure of the water vapour in the gas burette, PH2O. Use the table at the end of the lab. You may have to use a linear interpolation between points in the table (see explanation after the table). 3. Calculate the pressure of hydrogen in the gas burette (PH2) 4. Convert your measured temperature from the Celsius to the Kelvin scale. 5. Based on the chemical equation for the reaction between magnesium and hydrochloric acid (equation (3.13)), calculate the number of moles of hydrogen gas that were produced in the reaction. 6. Using the experimental data and the ideal gas equation, calculate an experimental value for the gas constant in L kPa mol-1 K-1. 7. Given the accepted value for ‘R’ is 8.314 L kPa mol-1 K-1, determine the percentage error in this experiment: (3.30) 8. Using “dimensional analysis” show that 1 L kPa = 1 J. Hints: P = F / A; 1 Pa = 1 N / m2; energy = F x d. 9. Why must you wait after the end of the reaction before taking your measurements? 76 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 10. If some magnesium remained after the reaction finished, how would this affect your measured value of ‘R’? Why? 11. If you did not consider the partial pressure of the water in the gas burette (PH2O), how would this affect the measured value of ‘R’? Why? 77 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 THE VAPOUR PRESSURE OF WATER Temperature / T/K Pressure / kPa 15 288.15 1.71 16 289.15 1.81 17 290.15 1.93 18 291.15 2.07 19 292.15 2.20 20 293.15 2.33 21 294.15 2.49 22 295.15 2.64 23 296.15 2.81 24 297.15 2.99 25 298.15 3.17 26 299.15 3.36 27 300.15 3.56 28 301.15 3.77 29 302.15 4.00 30 303.15 4.24 0 C Linear Interpolation It is very likely that the temperature you record for this experiment will lie somewhere between the values given in this table. To estimate the vapour pressure a “linear interpolation” will be used. This method assumes that between two values in the table the dependent variable (here, vapour pressure) 78 UOIT Chem. 1800U, W12; Exp. 3 - 9.14 depends linearly on the independent variable (temperature). The following equation can then be used: (T2,P2) and (T1,P1) represent the two known pairs of data from the table and (Tinterp, Pinterp) is the interpolated point. So: As an example, if Tinterp = 17.60C (the measured temperature) then Pinterp = 2.01 kPa (the interpolated water vapour pressure). The astute reader will notice from the data in the table that the vapour pressure is not a linear function of temperature. Therefore a linear interpolation is not strictly valid. However, when the ‘x’ and ‘y’ intervals are “small” it is a reasonable approximation. 79 UOIT Chem. 1800U, W12; Exp 2- 9.13 4. EFFECT OF CONCENTRATION AND TEMPERATURE ON REACTION RATE Objectives 1) To observe the effect of reactant concentration on reaction rate; 2) to observe the effect of temperature on reaction rate; 3) to determine the order of a simple reaction with respect to one of the reactants; 4) to calculate the activation energy for the same reaction. Introduction Physical chemistry is a huge branch of chemistry and concerns itself with understanding chemical phenomena by applying mathematical and physical models. In the world of chemistry physical chemistry is the branch closest to physics. Indeed, the lines between physical chemistry, chemical physics and physics are very blurry in modern research. In very broad terms “classical” physical chemistry can be divided into two categories: thermodynamics (including equilibria) and kinetics. Crudely speaking, thermodynamics is concerned with whether a reaction will occur and what the relative amounts of reactants and products will be (“where” the reaction is going). Kinetics is concerned with how quickly the reaction gets there. Detailed kinetics studies can also lead to conclusions about the reaction pathway or mechanism. Ultimately, a thorough understanding of a chemical process requires knowledge of both the kinetics and the thermodynamics. For example, thermodynamics predicts that all diamonds will transform into graphite. However, kinetics indicates that under “normal” conditions the transformation will take a phenomenally long time. For this reason chemists refer to substances as either “thermodynamically” or “kinetically” stable. Diamond is kinetically stable. 80 UOIT Chem. 1800U, W12; Exp 2- 9.13 Rate Laws and the Effect of Concentration Consider the generic chemical reaction: aA + bB v cC + dD (4.1) where the lower case letters are used to represent the stoichiometric coefficients. The rate of the reaction can be defined as: (4.2) The negative signs appear for the first two terms because ‘A’ and ‘B’ are disappearing (i.e., Ä[A] and Ä[B] < 0). A “rate law” for the process can be written as: (4.3) where ‘k’ is the rate constant and ‘x’ and ‘y’ are empirically derived constants. The values of ‘x’ and ‘y’ may be either integral or non-integral and have no necessary relationship to the stoichiometric ratios in the chemical equation (although, by coincidence, they may be the same). The value of ‘x’ is the order of the reaction with respect to ‘A’ and ‘y’ is the order of the reaction with respect to ‘B’. The sum x + y is called the overall order of the reaction. In this experiment the following reaction will be studied: 2H+(aq) + S2O32-(aq) v S(s) + SO2(aq) + H2O(l) (4.4) The rate of this reaction can be determined by measuring how long it takes for a given amount of sulphur to be produced: (4.5) 81 UOIT Chem. 1800U, W12; Exp 2- 9.13 (4.6) If [H+] is constant (as will be the case in this experiment): (4.7) where (4.8) and ‘m’ is the order of the reaction with respect to S2O32-. In this experiment the amount of sulphur will not be explicitly measured, but the time to produce a constant (but unknown) amount of sulphur will be measured. Consequently, the rate constant for the reaction cannot be measured, but the order of the reaction can be determined by plotting -ln Ät versus ln [S2O32-] (see question 1). It should be fairly clear from equation (4.7) that as long as Ä[S] is constant the relationship between the rate and Ät is a simple one. The rate is inversely proportional to time: the shorter the time the faster the reaction. Effect of Temperature on Rate Rate constants vary with temperature according to the Arrhenius equation: (4.9) where A is a constant, R is the ideal gas constant, T is the absolute temperature (in Kelvin) and Ea is the activation energy (a constant) of the reaction. If the rate constant of the reaction is measured 82 UOIT Chem. 1800U, W12; Exp 2- 9.13 at several different temperatures, the activation energy of the reaction can be found by re-casting the Arrhenius equation: (4.10) (4.11) A plot of ln k (on the y-axis) versus 1 / T (on the x-axis), then, would give a straight line with a slope of -Ea / R. For this experiment: (4.12) (4.13) If Ä[S] and [S2O32-] are kept constant we can re-write this equation: (4.14) where b is a constant. Therefore: (4.15) (4.16) As ln A and ln b are both constants their difference must also be constant and can be designated as 83 UOIT Chem. 1800U, W12; Exp 2- 9.13 AN: (4.17) Thus, if the time required to precipitate a known (constant) amount of sulphur is measured as a function of temperature (using a constant initial [S2O32-]), the activation energy for the reaction can be determined. As before, a plot of -ln Ät vs. 1 / T should be linear with a slope of -Ea / R. Determining how long it takes to precipitate a fixed amount of sulphur is a fairly straightforward task. As the sulphur precipitates from this reaction, the solution will become more and more cloudy. At some time the solution will be completely opaque. If the same vessel and solution volume are used for each run, it will require the same amount of sulphur to render the solution opaque. So, if the time required to make the solution opaque is measured for different concentrations of reactants, the time to precipitate a constant amount of sulphur will be known and the dependence of reaction rate on concentration can be deduced. How can the solution be judged to be “opaque”? A cross is made on a piece of paper and placed beneath the beaker. Once the cross becomes invisible (while looking down on the beaker) the solution is judged to be “opaque” and the time to reach this point is recorded. If the same (or at least very similar) beaker, solution volume and cross are used for each run, the disappearance of the cross will represent the precipitation of the same amount of sulphur. The astute student will notice that the reaction rate depends on the concentration of the reactants and that the concentration of the reactants must, by definition, decrease as the reaction proceeds. Therefore, the reaction rate cannot be constant. This is true and in this experiment the initial rate is measured. In any event, the rate constant and the order of the reaction do not change as the reaction proceeds. Consider a simple reaction such as: A v B+C 84 UOIT Chem. 1800U, W12; Exp 2- 9.13 where: (4.18) It is a reasonably simple problem in calculus to show that for a first order reaction: (4.19) If the initial concentration of A ([A]t=0) is relatively large, ‘t’ is short and ‘k’ is small, then the concentration of A will remain nearly constant and so will the reaction rate. This is the situation in this experiment. Apparatus and Materials 1) 2 x 150 mL beakers; 2) 100 mL graduated cylinder; 3) 10 mL graduated cylinder; 4) stopwatch; 5) 20 mm x 150 mm test tube; 6) hot plate; 7) thermometer; 8) sheet of paper marked with a large X; 8) 0.20 mol L-1 sodium thiosulphate (Na2S2O3); 9) 2.0 mol L-1 hydrochloric acid (HCl) Safety The main safety concern in this experiment is the evolution of SO2 as a product of the reaction. Once the reaction is “complete” make sure the solutions are immediately disposed of in the waste buckets in the fume hood. • Spills of HCl can be neutralized with a small amount of bicarbonate and the resulting solution mopped up with a sponge. Wear gloves to avoid skin irritation. Wash hands thoroughly. 85 UOIT Chem. 1800U, W12; Exp 2- 9.13 • Large spills of Na2S2O3 can be contained with sand or vermiculite and then disposed of as chemical waste. Small spills can be mopped up with a sponge. Wear gloves to avoid irritation. Rinse sponge well after clean up. Wash hands thoroughly. • Skin exposure to either reagent should be treated with copious amounts of cool, running water. • Eye exposure should be treated with the eye-wash station for at least 30 minutes followed by medical attention. • Do not use cracked or damaged glassware, especially in part II when solutions must be heated. • Students with allergies to bisulphites should be aware that this experiment generates SO2(g) which forms bisulphites when dissolved in water. Procedure You will work with a partner in this experiment. One partner should work on part I and the other should work on part II. After each kinetics measurement be sure to thoroughly clean the beaker with tap water followed by deionized water. Any cross-contamination of the thiosulphate solution with hydrochloric acid will cause erroneous results; SO2 will be generated prematurely and will stink! At the beginning of the experiment collect ~200 mL of 0.20 mol L-1 Na2S2O3 and ~50 mL of 2.0 mol L-1 HCl in separate, clean, dry, and labelled beakers. These should be sufficient quantities 86 UOIT Chem. 1800U, W12; Exp 2- 9.13 for both parts of the experiment. Do not waste your solutions. Part I: Determination of the Effect of Concentration on Reaction Rate 1. On a white piece of paper mark a large ‘X’. This paper will have to be used throughout the experiment so do not lose it! 2. To the first clean and dry beaker add 50.0 mL (use a 100 mL graduated cylinder) of 0.20 mol L-1 Na2S2O3. 3. Measure and record the temperature of the thiosulphate solution to at least one decimal place. For Part I, you may assume the temperatures of all the solutions are the same. 4. Place the beaker over the X and look down on the cross. 5. Add 5.0 mL (use a 10 mL graduated cylinder) of 2.0 mol L-1 HCl all at once. Immediately start the timer (if you are using a wristwatch, note the start time). Swirl the beaker once. 6. When the ‘X’ completely disappears, stop the timer (or note the end time). Record the time for the reaction. 7. Immediately dispose of the contents of the beaker in the fume hood. The gas produced is SO2 which is foul smelling and mildly toxic. 8. Repeat steps 2 - 7 with the quantities of reagents given in Table 1. After each reaction is complete thoroughly rinse the beaker with tap water followed by deionized water. If the beaker is to be re-used, dry it well with paper towel. 87 UOIT Chem. 1800U, W12; Exp 2- 9.13 Table 1: Reagent Volumes for Part I Beaker: I II III IV V 0.20 mol L-1 Na2S2O3 / mL 50 40 30 20 10 deionized water / mL 0 10 20 30 40 2 mol L-1 HCl / mL 5 5 5 5 5 Note that in each case the total volume is constant (55 mL) so [S2O32-] changes in each beaker, but [HCl] does not. Part II: The Effect of Temperature on Reaction Rate 1. Use a second set of clean, dry beakers for part II of the experiment. 2. Pour 10.0 mL of 0.20 mol L-1 Na2S2O3 and 40 mL of deionized water into one of the beakers. Note, these are the same quantities used in run V in Part I above. 3. Put 5.0 mL of 2.0 mol L-1 HCl in a large (20 mm x 150 mm) test tube. Do not get any HCl on the outside of the test tube. If you do, clean and dry the outside of the test tube thoroughly. 4. Place the large test tube in the beaker containing the Na2S2O3 + water solution. 5. Using the hot plate heat the thiosulphate and hydrochloric acid solutions to ~300C. Measure the temperature using a thermometer immersed in the Na2S2O3 + water solution. Note, the temperature does not have to be exactly 300C. 6. Remove the beaker from the hot plate and place it over the cross as in Part I. Wait 30 s and 88 UOIT Chem. 1800U, W12; Exp 2- 9.13 then measure the temperature to 0.10C. At this point it is important to know the temperature precisely. The temperature of the HCl can be assumed to be the same as the Na2S2O3 + water solution. 7. Immediately pour the acid into the thiosulphate solution. Record the time it takes for the cross to become invisible. Rinse and dry the outside of the test tube used to hold the HCl. 8. Repeat steps 2 - 7 using the same quantities of Na2S2O3, H2O and HCl but use the following, approximate, temperatures: 400C, 500C, 600C. The data from Part II can be combined with the data collected at room temperature (from Part I). 89 UOIT Chem. 1800U, W12; Exp 2- 9.13 Experiment 4 - Effect of Concentration and Temperature on Reaction Rate Name: ____________________________ Partner ____________________________ Day: ____________________________ Time: ____________________________ TA Name: ____________________________ 90 UOIT Chem. 1800U, W12; Exp 2- 9.13 Results, Discussion and Questions Part I: Determination of the Effect of Concentration on Reaction Rate Temperature: _______________________ [HCl] [S2O32-] Time, Ät / mol L-1 Beaker / mol L-1 /s -ln (Ä t) ln [S2O32-] I II III IV V Show sample calculations for the concentrations in this table. 1. Beginning with equation (4.7) show that the relationship between -ln Ät and ln [S2O32-] is linear and that the slope of the line gives the order of the reaction with respect to S2O32-. 2. Using SigmaPlot or Excel, create a graph of -ln Ät (y-axis) versus ln [S2O32-] (x-axis). Use the linear regression feature to find the equation of the line relating -ln Ät (y-axis) and ln [S2O32-]. From this equation determine the order of the reaction with respect to S2O32-. 91 UOIT Chem. 1800U, W12; Exp 2- 9.13 Part II: The Effect of Temperature on Reaction Rate Concentration of S2O32- Beaker Temp / 0C _________________________ T/K / K-1 Ät / s -ln Ät part I I II III IV 3. Plot a graph of -ln Ät (y-axis) vs. 1 / T (x-axis) and use a linear regression to find the slope of the line. 4. From the slope of the line and using equation (4.17) determine the activation energy for the reaction between HCl and Na2S2O3. 5. If [HCl] and [Na2S2O3] are kept constant, by how many degrees would this reaction have to be heated (from 250C) to precisely double the reaction rate? 92 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 5. THERMOCHEMISTRY AND THE MEASUREMENT OF THE ENTHALPY OF DISSOCIATION FOR ACETIC ACID Objectives: 1) To measure the enthalpy of neutralization for the reaction between hydrochloric acid (HCl) and sodium hydroxide (NaOH); 2) To measure the enthalpy of reaction between acetic acid (CH3COOH) and sodium hydroxide (NaOH); 3) To calculate the enthalpy of dissociation for aqueous acetic acid. Introduction Classical physical chemistry can be crudely divided into two very broad categories: thermodynamics and kinetics. This experiment is about thermochemistry which is a small part of the vast field of thermodynamics. Because of its huge range of application, thermodynamics is quite possibly the most important area of study in science and technology. It plays a central role in everything from biochemical regulation to running nuclear power plants. Like it or not, thermodynamics rules our lives! Yet, it is based on three simple laws. The first of these is well known to anyone with even the most modest understanding of science: energy can neither be created nor destroyed (the Law of Conservation of Energy). The other two laws concern entropy (broadly, the amount of “randomness” in a system). If the truth of these laws in accepted (and all experimental evidence points to their validity) four fundamental equations (the “Four Fundamental Equations of Gibbs”) can be derived and from these all the remaining equations of chemical thermodynamics can be derived using relatively simple mathematics. The derivations are mathematically rigorous. In this experiment the enthalpy of reaction ÄHrxn will be determined for the reaction between hydrochloric acid and sodium hydroxide: 93 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 HCl(aq) + NaOH(aq) v NaCl(aq) + H2O(l) (5.1) and for the reaction between acetic acid and sodium hydroxide: CH3COOH(aq) + NaOH(aq) v NaCH3COO(aq) + H2O(l) (5.2) Once these two enthalpies of reaction are known the enthalpy of reaction for the dissociation of acetic acid: CH3COOH(aq) v CH3COO-(aq) + H+(aq) (5.3) can be calculated using Hess’s Law (see below). The enthalpies of reaction for reactions (5.1) and (5.2) are found by measuring the heat evolved (or absorbed) when known quantities of the reactants are combined. The heat transferred in a chemical reaction is measured using a technique known as calorimetry. Calorimetry Chemists are most interested in thermodynamics as it relates to chemical reactions. As such, they study the heat released (or absorbed) in the course of a chemical reaction. This is a field of study known as “thermochemistry”. At the heart of thermochemistry is the experimental technique known as “calorimetry”. The device used to perform calorimetry is called (not surprisingly!) a “calorimeter”. An ideal calorimeter is one that does not allow the transmission of heat to or from the “system” (the reaction or process under study). Calorimeters can range from the very sophisticated (costing many thousands of dollars) to the very simple - the well-known polystyrene coffee cup, for example. To begin, some definitions are required: qsystem: the heat released (or absorbed) by the reaction or process (in some books this 94 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 is called qreaction or qrxn). qsurroundings: the heat absorbed (or released) by the calorimeter (qcalorimeter), its contents (qcontents) and the rest of the universe (quniverse): (5.4) The First Law of Thermodynamics states that all the heat must be accounted for, so: (5.5) In other words, what is lost (or gained) by the system must be exactly balanced by what is gained (or lost) by the surroundings. In a calorimetry experiment qsurroundings is measured by measuring the temperature change of the surroundings. In this experiment qcalorimeter and quniverse are assumed to be zero. Physically, this is equivalent to assuming the calorimeter absorbs no heat and no heat is “lost” from the calorimeter to the “universe”. Thus: (5.6) (5.7) (5.8) The heat absorbed or (released) by the contents can be found from: (5.9) where mcontents = mass of the contents cp, contents = specific heat capacity at constant pressure ÄTcontents = temperature change of the contents. The specific heat capacity of the contents must be measured as a separate experiment. This is usually done by adding a known amount of heat (often electrically) and measuring the temperature change for a known mass of the contents. In a calorimetry experiment, then, the temperature change 95 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 of a known mass is measured and qcontents is calculated. This procedure will be followed in this experiment. The molar enthalpy change of the reaction, ÄHrxn, can be related to the heat of the reaction (qsystem). In this experiment the reaction is carried out at constant pressure. The calorimeter is not tightly sealed (as would be the case in constant volume calorimetry). If any gases evolve, they escape to the atmosphere, so the reaction remains at constant pressure. If the pressure remains constant, then the molar enthalpy change of the reaction is equal to: (5.10) where nrxn is the number of moles of a product or reactant involved in the reaction (see below). Note that ÄHrxn can be either positive or negative. When ÄHrxn < 0, the reaction is said to be “exothermic” (heat is evolved) and when ÄHrxn > 0 the reaction is said to be “endothermic” (heat is absorbed). Standard Molar Enthalpy Change and Hess’s Law A thermochemical equation is a chemical equation that includes the molar enthalpy for that reaction, e.g.: H2(g) + ½ O2(g) v H O(l); 2 ÄH0rxn = -285.8 kJ mol-1 (5.11) The superscript ‘0' is used to indicate the enthalpy change is a standard” enthalpy; the pressure of the system is 1 bar. The molar enthalpy change is for the reaction as written. Thus, in this example, 285.8 kJ are evolved for every mole of water produced (or equivalently, for every mole of H2(g) consumed). If the chemical equation is multiplied by 2 to give: 2H2(g) + O2(g) v 2H O(l) (5.12) 2 the molar enthalpy change must also be multiplied by 2 to give -571.6 kJ mol-1. In a thermochemical equation it is also understood that the reagents and products are at the same temperature. Commonly 96 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 enthalpy changes are reported for systems at 298.15 K. The states of all the species in the reaction must also be specified. If a thermochemical equation is reversed, the sign of the molar enthalpy change is reversed. This is because enthalpy is a “state” function. The enthalpy depends only on the “state” of the system and not on how the system reaches (the “path”) that state. The state is defined by the pressure, temperature and composition of the system. Because enthalpy is a state function the change in enthalpy, ÄH, is also a state function and does not depend on the path. Hess’s Law: The overall enthalpy change for a reaction is equal to the sum of the enthalpy changes for the individual steps in the reaction. is based on this property of enthalpy changes. As a consequence, a set of reactions (with known ÄH’s) can be combined to determine the ÄH for a reaction with an unknown molar enthalpy change. This is tremendously useful if the enthalpy change for a reaction is difficult to measure. Consider, for example, the reaction: C(s) + 2H2(g) v ÄH05.13 CH4(g); (5.13) This reaction is difficult to carry out “cleanly”; a mixture of products is likely to result. However, this reaction can be “constructed” from a sequence of reactions where each reaction is “clean” and the enthalpy change is comparatively easy to measure. Reaction (5.13) can be generated from the following reactions: C(s) + O2(g) v CO (g); ÄH05.14 = -393.5 kJ mol-1 2 H2(g) + ½ O2(g) v H2O(l); CH4(g) + 2O2(g) v CO2(g) + 2H2O(l); ÄH05.15 = -285.8 kJ mol-1 ÄH05.16 = -890.3 kJ mol-1 in the following way: 97 (5.14) (5.15) (5.16) UOIT Chem. 1800U, W12; Exp. 5 - 1.1 C(s) + O2(g) (5.14) 2H2(g) + O2(g) 2 x (5.15) (5.16), reversed CO2(g) + 2H2O(l) v CO2(g) v v 2H2O(l) CH4(g) + 2O2(g) ÄH05.14 = -393.5 kJ mol-1 2 x ÄH05.15 = -571.6 kJ mol-1 -ÄH05.16 = 890.3 kJ mol-1 (5.14) + 2 x (5.15) + (5.16) reversed gives: C(s) + 2H2(g) v CH (g); ÄH05.13= ÄH05.14 + 2 x ÄH05.15 - ÄH05.16 = -74.8 kJ mol-1 4 Reversing a chemical equation is the same as “subtracting” it. Standard Molar Enthalpies of Formation and Standard Molar Enthalpy Changes The standard molar enthalpy of formation, ÄH0f, is the enthalpy change associated with the formation of 1 mole of compound at 1 bar from its elements (in their most stable forms) at 1 bar. The temperature must also be specified, but most commonly it is 298.15 K. The standard molar enthalpies of formation for elements in their most stable forms are defined as zero. Standard molar enthalpies of formation can be used to calculate standard molar enthalpy changes of reactions. This is also a consequence of the fact that enthalpy change is a state function. Essentially, the path is: reactants at 1 bar, 298 K ! elements in most stable ! products at 1 bar, 298 K forms at 1 bar, 298 K The standard molar enthalpy for the reaction will be given by: (5.17) where n = the stoichiometric coefficient for each of the products m = the stoichiometric coefficient for each of the reactants 98 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 When it comes to the measurement of standard enthalpies of formation of ions, a complication arises. Ions must come in pairs (to maintain electrical neutrality). Consider the dissolution of HCl(g) in water: HCl(g) v H+(aq) + Cl-(aq) ÄH0rxn = -74.14 kJ mol-1 (5.18) (5.19) ÄH0f(HCl(g)) = -92.30 kJ mol-1. Therefore: (5.20) Because it is impossible to measure the individual enthalpies of formation of ions (since ions cannot occur separately), it is impossible to know how much of the -167.44 kJ mol-1 “belongs” to the H+(aq) and how much to the Cl-(aq). To get around this problem the ÄH0f (H+(aq)) is defined as zero. In other words: (5.21) Once this is done the ÄH0f (Cl-(aq)) has a value (from equation 5.20). Once this is assigned the ÄH0f of other ions can be calculated, based on the value of ÄH0f (Cl-(aq)). Through several iterations,values of ÄH0f can be assigned to all ions. The measurements carried out in these experiments will be made at room temperature as opposed to 298.15 K. Thermodynamic data is normally tabulated at 298.15 K and standard enthalpy changes do depend on temperature. However, for small temperature variations the effect of temperature on enthalpy changes is very small. Therefore, in this experiment the enthalpy change is assumed to be independent of temperature. 99 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 Apparatus 1) 50 mL beaker; 2) two coffee cups + 1 lid; 3) 250 mL beaker; 4) magnetic stir bar; 5) thermometer; 6) rubber band; 7) magnetic stirrer; 8) 100 mL graduated cylinder; 9) stopwatch; 10) 1.0 mol L-1 NaOH; 11) 1.0 mol L-1 CH3COOH ; 12) 1.0 mol L-1 HCl. Safety • 1.0 mol L-1 HCl, NaOH and CH3COOH will each cause irritation of the skin with prolonged exposure. Rinse with plenty of cool running water. • If HCl, NaOH or CH3COOH splashes in the eye, flush for at least 30 minutes using the eyewash station; seek medical attention. Notify the teaching assistant, technician or senior laboratory instructor. • Spills of HCl and CH3COOH can be neutralized with baking soda and then mopped up with a sponge. Rinse the sponge thoroughly and clean hands with soap and water afterwards. • Spills of NaOH can be neutralized with vinegar and then mopped up with a sponge. Rinse the sponge thoroughly and clean hands with soap and water afterwards. Procedure You will perform this experiment with a partner. Part I: Reaction of Sodium Hydroxide with Hydrochloric Acid 1. Construct a simple coffee cup calorimeter as shown in Figure 5.1. The magnetic stir bar goes in the inner cup and is considered a part of the calorimeter. 100 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 Figure 5.1: Coffee Cup Calorimeter 101 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 2. Weigh the calorimeter assembly (excluding the thermometer and the elastic) on an open pan balance and record the mass. The calorimeter should be dry for this weighing. 3. Collect ~125 mL of 1.0 mol L-1 NaOH (note the actual concentration) in a clean and dry beaker. Rinse a graduated cylinder with tap water, then deionized water and finally rinse the cylinder with small portions of the NaOH. Pour 50.0 mL the NaOH into the inner cup of the calorimeter. Replace the lid of the calorimeter. Measure the temperature of the NaOH. 4. Collect ~60 - 70 mL of 1.0 mol L-1 HCl (note the actual concentration) in a clean, dry beaker. Measure the temperature of the HCl. If you use the same thermometer as in step 3, rinse the thermometer with deionized water and dry it thoroughly before transferring it. The temperature of the NaOH and the HCl should be the same (within 0.50C). If the temperatures are not, then either gently warm (on a hot plate) or cool (using an ice bath) the HCl until the two temperatures are within 0.50C. 5. Obtain a different graduated cylinder (from that used in step 3) and rinse it with tap water, then deionized water and finally rinse the cylinder with small portions of the HCl. Pour 50.0 mL of HCl into the graduated cylinder. 6. Place the calorimeter on a magnetic stirrer. Adjust the speed of the stirrer so that the magnetic stir bar spins rapidly but does not splash the sides of the cup with the NaOH. 7. Wrap an elastic band around the thermometer several times to support it in the lid of the coffee cup. Adjust the height of the thermometer so that the bulb is immersed in the NaOH but the stir bar is not striking the thermometer. Be sure you can read the temperature on the thermometer. 8. One complete run will take approximately 15 minutes (3 minutes before the addition of the 102 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 HCl and 12 minutes after). Temperatures must be recorded to 0.10C. Once the timer has started do not stop it until the end of the run. If the timer is stopped, the run must be repeated. 9. Set the stirrer to a relatively slow speed and reset the stopwatch to zero. 10. Measure and record the temperature and start the stopwatch. This temperature will be the temperature at 0 s. 11. Measure and record the temperature and time every 30 s for at least 3 minutes (180 s). The temperature is not likely to change very much. 12. After ~3 minutes remove the lid and quickly pour in the HCl (from the graduated cylinder). Avoid getting HCl on the sides of the cup; it should all be added to the NaOH solution. Carefully note the time at which the HCl is added. Once the HCl is added quickly replace the lid of the calorimeter (to prevent heat loss through the top of the cup). Do not attempt to measure the temperature at this point. 13. Record the temperature and time every 30 s for at least 12 more minutes. The temperature should rise to a maximum fairly quickly and then stabilize or fall very slowly. If it falls quickly there is a problem with the calorimeter (i.e., the lid is not properly sealed). 14. At the end of the run (~15 minutes after the initial reading), re-weigh the calorimeter (and the contents) on an open-pan balance. This is a crucial step and without it the experiment is meaningless. 15. Retrieve the magnetic stir bar and empty the contents of the calorimeter. Rinse the calorimeter thoroughly with deionized water. Dry the inner cup of the calorimeter. 103 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 16. Repeat Part I from from the beginning. Re-weigh the “empty” calorimeter. Chemistry 1800 students need only do one run. Part II: Reaction of NaOH with CH3COOH This procedure is almost identical to the reaction of NaOH with HCl. 1. Weigh the dry calorimeter with the stir bar on an open pan balance. Record this mass. 2. Add 50.0 mL of 1.0 mol L-1 NaOH to the inner cup of the calorimeter and measure its temperature. 3. Obtain 60 - 70 mL of CH3COOH in a clean, dry beaker. As in part I, measure its temperature and ensure that the temperature is within 0.50C of the NaOH solution in the calorimeter. 4. Once the temperatures are the same pour 50.0 mL of the CH3COOH solution into a clean, properly rinsed graduated cylinder. 5. Place the calorimeter on the magnetic stirrer. 6. Start the stirrer so that the stir bar is spinning but does not cause solution to splash on the sides of the cup. 7. Record the time and temperature of the NaOH in the calorimeter every 30 s for 3 minutes. 8. Remove the calorimeter lid and add the CH3COOH from the graduated cylinder. Again, avoid getting CH3COOH on the sides of the cup. Record the time at which the CH3COOH 104 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 is added. 9. Replace the calorimeter lid. Continue to record the time and temperature at 30 s intervals. 10. Continue to record the time and temperature until 12 minutes after adding the CH3COOH (15 minutes in total). 11. After the run in complete, re-weigh the calorimeter on the open-pan balance. 12. Empty, clean and dry the calorimeter and repeat Part II from the beginning. Chemistry 1800 students need only do one run. At the end of the experiment, retrieve the magnetic stir bar, empty the calorimeter and rinse it thoroughly with tap water followed by deionized water. Do NOT use soap to clean the calorimeter. Unplug the stirrer and wipe it with a damp J-cloth or sponge. 105 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 Experiment 5 - Thermochemistry Name: ____________________________ Partner ____________________________ Day: ____________________________ Time: ____________________________ TA Name ____________________________ 106 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 Results, Discussion and Questions NOTE: This is a very long report to prepare. Do not leave it to the last minute! Part I: NaOH + HCl Run I mass of empty calorimeter / g ________________________ mass of calorimeter + solution (end of run) / g ________________________ mass of solution / g ________________________ Time (s) Temp. / 0C Time (s) 107 Temp. / 0C UOIT Chem. 1800U, W12; Exp. 5 - 1.1 Run II (not required for Chemistry 1800) mass of empty calorimeter / g ________________________ mass of calorimeter + solution (end of run) / g ________________________ mass of solution / g ________________________ Time (s) Temp. / 0C Time (s) 108 Temp. / 0C UOIT Chem. 1800U, W12; Exp. 5 - 1.1 Part II: NaOH + CH3COOH Run I mass of empty calorimeter / g ________________________ mass of calorimeter + solution (end of run) / g ________________________ mass of solution / g ________________________ Time (s) Temp. / 0C Time (s) 109 Temp. / 0C UOIT Chem. 1800U, W12; Exp. 5 - 1.1 Run II (not required for Chemistry 1800) mass of empty calorimeter / g ________________________ mass of calorimeter + solution (end of run) / g ________________________ mass of solution / g ________________________ Time (s) Temp. / 0C Time (s) 110 Temp. / 0C UOIT Chem. 1800U, W12; Exp. 5 - 1.1 Determination of ÄTcontents For each run (for each part) plot a graph of temperature in Kelvin (on the ordinate or y-axis) versus time in seconds (on the abscissa or x-axis). Each graph should be on a separate piece of paper and should follow the guidelines given at the beginning of the laboratory manual. You may use either Excel of SigmaPlot to prepare the graphs. See the example (from a similar experiment from a previous year) on the following page. On each graph clearly mark the time of mixing. Using the linear regression function in SigmaPlot (or Excel), determine the “line of best fit” for the data before the time of mixing. Using this equation, calculate the temperature at the time of mixing. This temperature will be Ti. Next, apply a linear regression to the data after the time of mixing to find the “line of best fit”. Use only the data after maximum temperature has been attained. After the reaction is complete the temperature should either remain the same or decrease slightly. Using the equation calculate the temperature of the solution at t = tmixing. This temperature is Tf. The change in temperature, ÄT, is given by: (5.22) The extrapolations are used because the thermometer does not respond instantaneously to temperature. The extrapolations can also be used to account for heat losses from the calorimeter. Performing the regression analysis is best done by placing the data before mixing and after the maximum temperature into two sets of columns (see below). Regression analysis can then be easily performed on these two separate sets of data. The resulting regression plots can be added to the original graph. 111 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 112 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 113 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 Calculation of ÄH0rxn for Part I 1. Assume the heat capacity of the final solution is 4.184 J K-1 g-1. Using the final mass of the solution in the calorimeter, calculate qcontents from equation (5.9). Calculate qsystem from equation (5.8). 2. Calculate the moles of H2O (aq.) produced, nrxn. Be sure to identify the limiting reagent in the reaction. 3. Calculate the standard molar enthalpy change, ÄH0rxn, for reaction (5.1). Calculate the average ÄH0rxn from the two runs. Calculation of ÄH0rxn for Part II 1. Again, assume the heat capacity of the final solution is 4.184 J K-1 g-1. Calculate qcontents and qsystem. 2. Calculate the number of moles of NaCH3COO (aq.) produced, nrxn. Again, identify the limiting reagent. 3. Calculate the standard molar enthalpy change for reaction (5.2) for each of the two runs. 4. Calculate the average ÄH0rxn from the two runs. Questions 1. Using Hess’s Law and your results from parts I and II, calculate the enthalpy of reaction, ÄH0rxn, for reaction (5.3) 114 UOIT Chem. 1800U, W12; Exp. 5 - 1.1 2. Calculate the ÄH0f of the OH- ion given that the ÄH0f (H2O) = -285.8 kJ mol-1. 3. In this experiment it was assumed that the calorimeter neither absorbed nor released heat (i.e., qcalorimeter = 0), but this is almost certainly not the case. Would the measured value of ÄH0rxn for reaction (5.1) be more or less exothermic if the calorimeter absorbed heat (that is, if qcalorimeter > 0)? Explain why. 4. Without doing any calculations, predict whether the observed ÄT would increase, decrease or remain the same if the following procedural changes were made. Justify your predictions: a. 50.0 mL of 2.0 mol L-1 NaOH mixed with 50.0 mL of 2.0 mol L-1 HCl. b. 25.0 mL of 2.0 mol L-1 NaOH + 50.0 mL of water mixed with 25.0 mL of 2.0 mol L-1 HCl. 5. Use tabulated standard molar enthalpies of formation (from the textbook) to calculate ÄH0rxn for the following (unbalanced) reactions: a. b. H2S(g) + O2(g) ÿ SO2(g) + H2O(g) c. 6. N2O4(g) + H2(g) ÿ N2(g) + H2O(g) Fe2O3(s) + HCl(g) ÿ FeCl3(s) + H2O(g) One of the cleanest burning components of gasoline is 2,3,4-trimethylpentane which has the chemical formula C8H18. a. Write a balanced chemical reaction for the complete combustion of 2,3,4trimethylpentane to CO2(g) and H2O(g) such that the stoichiometric coefficient for C8H18 is 1. b. When 50.0 g of 2,3,4-trimethylpentane are completely burned at 298.15 K and 1 bar of pressure, 2.218 x 103 kJ are released. trimethylpentane. 115 Calculate the ÄH0f of 2,3,4- UOIT Chem. 1800U, W12; Exp. 6-9.13 6. ELECTROCHEMISTRY AND VOLTAIC CELLS Objectives 1) To measure the voltages generated by several common galvanic cells; 2) to observe the effect of changing concentration on cell potential; 3) to estimate the formation constant (Kf) for Cu(NH3)42+. Introduction Electrochemistry is an incredibly important area of chemistry. In its broadest definition electrochemistry is the study of the transfer of electrical charge (or energy) between chemical species or the generation of electrical energy (electrical work) from chemical systems. Originally electrochemistry was considered a part of physical chemistry, but now it plays important roles in almost all areas of chemistry: analytical, synthetic, energy science, and materials science. In this experiment the generation of electricity from chemical reactions will be studied. Devices that generate electricity from chemical reactions are usually referred to as voltaic (or sometimes galvanic) cells. Batteries are one common example of this phenomenon in everyday life. Reduction - Oxidation Reactions If a small amount of lead metal (Pb) is added to a solution of copper sulphate (which is blue) a reaction is observed. The colour of the solution will fade, copper metal will precipitate from the solution and lead will dissolve. The net ionic equation for this reaction is: Pb(s) + Cu2+(aq) v Pb2+(aq) + Cu(s) In a formal sense this reaction can be broken into two “half-reactions”: 116 (6.1) UOIT Chem. 1800U, W12; Exp. 6-9.13 Pb(s) v Cu2+ + 2e- Pb2+ + 2e- v Cu (6.2) (6.3) The first of these involves the loss of electrons and is called “oxidation” (the oxidation halfreaction). The second is the gain of electrons and is called “reduction” (the reduction half-reaction). In the overall reaction Pb is called the “reducing agent” (it causes Cu2+ to be reduced) and Cu2+ is the “oxidizing agent” (it causes Pb to be oxidized). The overall reaction is called a “redox” (an abbreviation of reduction-oxidation) reaction. Voltaic Cells If the electrons that are transferred from the reducing agent to the oxidizing agent can be forced to flow through an external circuit, the electrons can be used to do electrical work. This can be accomplished by physically separating the two half-reactions. The oxidation half-reaction takes place at the “anode” and the electrons flow through a wire to the “cathode” where the reduction reaction takes place. To maintain overall electrical neutrality, anions must flow to the anode and cations must flow to the cathode. This is often done using a salt bridge or, as in this experiment, a porous cup. Without the flow of anions and cations the flow of electrons stops. The combination of the anode half-cell, cathode half-cell and porous cup is called a voltaic cell and the overall chemical reaction is called the cell reaction. Voltaic cells are often described using a short hand notation: anode * anode solution 2 cathode solution * cathode The single vertical lines represent a change of phase (solid to solution, for example) and the double vertical line represents the porous barrier. As an example, the voltaic cell constructed in Figure 6.1 using lead and copper solutions would be written as: Pb * Pb2+ (aq, 1 mol L-1) 2 Cu2+ (aq, 1 mol L-1) * Cu 117 UOIT Chem. 1800U, W12; Exp. 6-9.13 Figure 6.1 - Electrochemical Cell 118 UOIT Chem. 1800U, W12; Exp. 6-9.13 Standard Reduction Potentials and Overall Cell Potentials Once the voltaic cell is constructed a voltage (or cell potential) can be measured in the external circuit. The voltage measured depends on the reactants, products and their concentrations. A predicted voltage for a voltaic cell can be calculated using a table of standard reduction potentials. The overall cell potential can be regarded as the sum of two half-cell potentials, one for oxidation and one for reduction. The half-cell potentials are tabulated (by convention) as reduction potentials. The potential for a single half-cell can never be measured; only overall cell potentials can be measured. Therefore, half-cell reactions are only measured with respect to some reference electrode which is arbitrarily assigned a potential of 0 V. The reference chosen is the standard hydrogen electrode (SHE): H+ (aq, 1 mol L-1) + e- v½H 2 E0 = 0.00 V (g, 1 bar) (6.4) The superscript ‘0’ indicates this is a “standard” potential. That is, it is “measured” under standard conditions: pure liquids and solids in their most stable forms at 1 bar, 1 mol L-1 for dissolved species, and 1 bar for all gases. Standard reduction potentials are tabulated and can be used to calculate overall cell potentials simply by adding the two standard half-cell potentials involved. If the reaction in the half-cell is an oxidation reaction, the sign of the reduction potential is simply reversed. For reaction (6.1): Pb2+ + 2ePb v v E0red = -0.126 V Cu2+ + 2e- v Cu Pb(s) + Cu2+(aq) v (6.6) E0red = 0.3419 V Pb2+ + 2e- (6.5) E0ox = 0.126 V Pb (6.7) E0cell = 0.468 V Pb2+(aq) + Cu(s) E0cell is always the sum: (6.8) 119 UOIT Chem. 1800U, W12; Exp. 6-9.13 When adding half-reactions, it is important to ensure that the number of electrons is the same in each. This ensures the electrons “cancel” in the overall cell reaction. The electrons can be balanced by multiplying the half-reactions by suitable factors. When a half-reaction is multiplied by a constant, the half-cell potential remains the same. If the overall cell potential is positive (as in this case), the reaction proceeds spontaneously in the direction as written. If it is negative, the reaction is spontaneous in the opposite direction. Effect of Concentration on Cell Potential The concentration of the species in the voltaic cell also affects the measured voltage. The relationship between concentration and voltage is given by the Nernst Equation: (6.9) E = observed voltage E0 = standard cell potential R = gas constant, 8.314 J mol-1 K-1 T = absolute temperature, K n = number of moles of electrons transferred in the reaction as it is written F = Faraday’s Constant, 96 485 C mol-1 Q = the reaction quotient. This has the same form as the equilibrium constant but the concentrations are those of the cell as opposed to the equilibrium concentrations. For reaction (6.1): (6.10) If the concentrations of Pb2+ and Cu2+ are both 1 mol L-1, the ratio is 1 and E = E0 (as it should be). 120 UOIT Chem. 1800U, W12; Exp. 6-9.13 The E and E0 values can also be used to calculate unknown concentrations: (6.11) If the concentration of Cu2+ were known, the [Pb2+] could be calculated from: (6.12) (6.13) In this experiment the cell potentials for a number of reactions will be measured and compared to calculated values. The influence of concentration on voltage will be examined and finally an equilibrium constant for the reaction: Cu2+(aq) + 4NH3(aq) v Cu(NH3)42+(aq) (6.14) will be estimated. Safety • All the metal solutions should be collected in the waste bottles. • Small spills of the metal solutions can be cleaned up with paper towel or a sponge. Rinse the sponge thoroughly. Thoroughly wash hands afterwards. For larger spills - use an absorbent such as sand and collect the adsorbent as hazardous waste. • Wash hands after handling the electrodes. 121 UOIT Chem. 1800U, W12; Exp. 6-9.13 • The most dangerous reagent is the 6 M NH3. Only small quantities of this will be provided, so any spills should be very small. Treat with a small amount of vinegar and clean up with a sponge. Rinse sponge thoroughly. Wear gloves. Wash hands thoroughly afterwards. • For all reagents treat eye exposure at the eye-wash station for at least 15 minutes (30 minutes for the ammonia) - seek medical attention. • For skin exposure: treat with cool running water for several minutes. If skin irritation develops, seek medical attention. • Magnesium is combustible. Magnesium fires cannot be extinguished with an ABC extinguisher - they must be put out with a type D fire extinguisher. They can also be extinguished with sand or sodium chloride. Procedure You will perform this experiment with a partner. You must complete and submit the laboratory report before the end of the laboratory period. Note that most of the data required for the table in part I (anode and cathode half reactions and the theoretical voltage) can and should be filled in before coming to the laboratory. Part I: Voltaic Cells 1. Obtain ~10 mL of 0.2 mol L-1 MgSO4 and pour it into the porous cup. The porous cup contains the anode half-cell. 2. Insert the magnesium electrode (the thin ribbon) into the cup. Using the alligator clip, 122 UOIT Chem. 1800U, W12; Exp. 6-9.13 connect the magnesium strip to the voltmeter using the black wire. The other end of the wire should be connected to the “-” terminal of the voltmeter. The electrode in the porous cup should always be connected to the “-” terminal of the voltmeter. 3. Obtain ~10 mL of 0.2 mol L-1 CuSO4 and pour it into the glass electrode jar. Clamp the copper electrode in the holder provided. Ensure the tip of the electrode is in the solution. Using the alligator clip connect the red wire to the copper electrode. The other end of the red wire should be connected to the red terminal of the voltmeter. The glass electrode jar contains the cathode half-cell. 4. Measure and record the voltage. 5. Place the porous cup inside the glass electrode jar and record the voltage. If no voltage is observed, reverse the connections at the voltmeter. Do not leave the external circuit connected for too long as the flow of electrons will alter the concentrations of the ions in the cell leading to a change in the voltage. 6. Remove the porous cup, rinse the outside with deionized water, dry it and set it aside. The MgSO4 solution can be re-used. 7. The copper solution will also be re-used. Decant it into a clean, dry beaker. 8. Rinse the glass electrode jar with deionized water and dry it. 9. Obtain 10 mL of 0.2 mol L-1 ZnSO4 solution and pour it into the glass electrode jar. As with the copper half-cell, use the zinc strip as an electrode and connect it to the voltmeter. 10. Repeat steps 5 - 8. The zinc solution can be saved and re-used. 123 UOIT Chem. 1800U, W12; Exp. 6-9.13 Obtain ~10 mL of Fe2+ solution and an iron electrode and use them to construct the cathode 11. half-cell. Repeat steps 5 - 8. The iron solution can also be re-used. 12. Repeat the measurements with the following combinations: Porous cup (anode half-cell) connected to “-“ connected to the red terminal 1 Mg / Mg2+ Cu / Cu2+ 2 Mg / Mg2+ Zn / Zn2+ 3 Mg / Mg2+ Fe / Fe2+ 4 Zn / Zn2+ Cu / Cu2+ 5 Fe / Fe2+ Cu / Cu2+ 6 13. Glass Electrode Jar (cathode half-cell) Zn / Zn2+ Fe / Fe2+ After all the measurements have been made thoroughly rinse the porous cup and glass electrode jar. Part II: Approximating Kf of Cu(NH3 )42+ 1. Put ~10 mL of 0.2 mol L-1 CuSO4 into the porous cup. Insert a copper electrode and connect it to the voltmeter (to “-“). 2. Put ~10 mL of 0.2 mol L-1 CuSO4 in the glass electrode jar. Insert a copper electrode and connect it to the voltmeter. 3. Place the porous cup in the glass electrode jar. The voltage should be zero. 4. Add 4 mL of 6 mol L-1 NH3 to the porous cup and swirl the cup. If any precipitate forms, add 124 UOIT Chem. 1800U, W12; Exp. 6-9.13 a little more NH3 and swirl again. 5. Record the voltage. 6. Thoroughly rinse the porous cup (inside and out) and the glass electrode jar. 125 UOIT Chem. 1800U, W12; Exp. 6-9.13 Experiment 6 - Electrochemistry Name: ____________________________ Partner ____________________________ Day: ____________________________ Time: ____________________________ TA Name: ____________________________ 126 UOIT Chem. 1800U, W12; Exp. 6-9.13 Results, Discussion and Questions Part I Complete the following table. Anode half-reaction Cathode half-reaction Exp. Theor. Voltage Voltage 1 2 3 4 5 6 Part II Voltage before addition of NH3 ____________________ Voltage after addition of NH3 ____________________ 127 UOIT Chem. 1800U, W12; Exp. 6-9.13 Questions 1. Why is no voltage observed before the porous cup is immersed in the glass electrode jar? 2. If a voltaic cell were constructed with two copper half cells and the concentrations of Cu2+ were the same in each half-cell, what would the measured voltage be? If the Cu2+ concentration in the anode half cell were 0.0020 mol L-1 and Cu2+ concentration in the cathode half cell were 0.20 mol L-1, what would the observed voltage be (assume T= 298.15 K)? Why is it not practical to perform this experiment with the voltmeters used in this experiment? 3. The reaction between Cu2+ and NH3 is: Cu2+(aq) + 4NH3(aq) v Cu(NH3)42+(aq) (6.14) For which an equilibrium constant can be written: (6.15) From the observed voltage in part II (after adding NH3) and the Nernst Equation: (6.16) calculate the concentration of “free” Cu2+(aq) ([Cu2+]cup) in the porous cup (assume T = 298.15 K). This quantity should be very small. Estimate the moles of Cu(NH3)42+ formed, assuming all the original Cu2+ in the porous cup reacted to form the Cu(NH3)42+. Because [Cu2+]cup should be very small this is a reasonable approximation. Calculate the [Cu(NH3)42+] in the porous cup. 128 UOIT Chem. 1800U, W12; Exp. 6-9.13 Assuming the concentration of ammonia in the porous cup is 1.1 mol L-1, calculate the equilibrium constant, Kf. 129 ...
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