CHEMISTRY 114.1 QUANTITATIVE AND QUALITATIVE ANALYSIS Fall 2008 SECTION 1. INSTRUCTOR AND COURSE INFORMATION Instructor: Mr. Xianbo Shi Office: Remsen 017 Office hours: Wednesday, 1:00-2:00 pm Telephone: (718) 997-4182 E-mail: xianbo.shi@qc.cuny.edu Laboratory: Tuesday, 1:40 4:30 PM; Remsen 209 Course Content: Goals/ Objectives: Webpage: Required Text: Quantiative and qualitative analysis of common household substances. The quantitative analysis of bottled and tap water. Qualitative analysis of heavy metal cations and common anions that act as contaminates in water. http://chem.qc.cuny.edu/~introchemlab/chem114/home.html Queens College Chemistry 114.1 Laboratory Manual. SECTION 2. POLICIES/RULES Attendance: Tardiness: Laboratories: Attendance in laboratory is mandatory. Missed laboratories cannot be made-up. An unexcused absence results in the loss of all points associated with that laboratory. If you have a university excused absence (such as illness, etc), you must show the excuse to the laboratory instructor the week after the absence. If you will miss a laboratory due to religious observance, you must inform the instructor the week BEFORE the absence, or the absence will not be excused. If the absence is excused, the instructor will make arrangements for you to make-up the laboratory during a different laboratory section (or a different laboratory period). It is your responsibility to attend class and to be punctual. Pre-laboratory quizzes will be given at the beginning of class and students will be given exactly 30 minutes to complete the quiz. Tardiness will not result in additional time being given for pre-laboratory quizzes. There are 14 laboratories, including the check-in/introductory laboratory and the final examination meeting. The grade will be determined from pre-laboratory quizzes, laboratory techniques, laboratory reports, and laboratory practical.
Pre-laboratory quizzes: Laboratory project: Laboratory worksheets: Laboratory reports: Academic dishonesty: All pre-laboratory quizzes will be 30 minutes in length and will be given at the beginning of the laboratory period. Additional time will not be given to students who are tardy. These quizzes will consist of three questions, which may contain multiple parts: Question 1. A post-laboratory question from the previous experiment. Question 2. A safety question about the current experiment. Question 3. A pre-laboratory question from the current experiment. A complete laboratory project consists of the following documents, which should be stapled together in the order listed: 1. Hard copy of the laboratory report 2. The laboratory report sheet 3. The post laboratory questions In addition to this package, an electronic version of the laboratory report should also be submitted to Blackboard. The completed project (i.e., Items 1-3) must be submitted at the beginning of the laboratory period on the date due (see the schedule). An electronic version (i.e., Word or WordPerfect) of the laboratory report must be submitted to Blackboard on the date due. Electronic versions of laboratory reports will be submitted to Turn-it-in software and checked for plagiarism. Turn-it-in checks both internet sources and previously submitted reports. Failure to submit an electronic version of the report will result in a zero on the laboratory report. For laboratories that are multi-part, worksheets accompany each day of the experiment. These worksheets must be submitted at the end of the laboratory in question. A style guide for the laboratory reports is attached to this document and can also be found on the course website. Unless specified by the laboratory instructor, laboratory reports are limited in length to 10 pages, excluding Figures and Tables. While it is natural to discuss experiments among each other, copying and/or plagiarism will not be tolerated on any assignment and will be treated in accordance with university policy. This policy can be found at: http://www1.cuny.edu/portal_ur/content/2004/policies/image/policy.pdf Electronic versions of laboratory reports will be submitted to Turn-it-in software and checked for plagiarism. Turn-it-in checks both internet sources and previously submitted reports. We should note here that copying and downloading graphs from the internet, without express permission of the instructor, constitutes plagiarism (even if referenced).
SECTION 3. GRADING Course grading: Laboratory projects (12@20 pts) 240 pts Laboratory quizzes (11@10 pts) 110 pts Quantitative practical (30 pts) 50 pts Qualitative practical: anions (25 pts) 25 pts Qualitative practical: cations (75 pts) 75 pts Total points: 500 pts Grading for Project: 1. Laboratory report * 7 pt a. Adherence to style guide (1 pt) b. Introduction and Experiment (2 pt) c. Results and Discussion (2 pts) d. Grammar and writing (2 pts) 2. Pre-laboratory questions 3 pt 3. Report sheet 4 pt 4. Post laboratory questions 3 pt Safety and laboratory technique 3 pt * Failure to submit an electronic version of the laboratory report will result in the loss of these 7 points.
Laboratory Schedule Lab Pre-Lab (60 min) Lab Work (120 min) 1 Safety exam. Syllabus and safety. 2 Pre-laboratory quiz. Buffers 3 Pre-laboratory quiz. ph of water 4 Pre-laboratory quiz. Water Hardness 5 Pre-laboratory quiz. Alkalinity Check-in and Techniques in titration. Equilibrium. Laboratory Project 1 due. Analysis of tap and bottled water: Part 1 and Part 2. Laboratory Project 2 due. Analysis of tap and bottled water: Part 2 and 3. Laboratory Project 3 due. Analysis of tap and bottled water: Part 3. Laboratory Project 4 due. 6 Pre-laboratory quiz. Finish the analysis of water. Laboratory Project 5 due. 7 Quantitative Analysis Practical. Laboratory Project 6 due. 8 Pre-laboratory quiz. Solution chemistry. Anion standards. 9 Qualitative Analysis Anion Practical. Laboratory Project 7 due. 10 Pre-laboratory quiz. Solution chemistry. 11 Pre-laboratory quiz. Solution chemistry. 12 Pre-laboratory quiz. Solution chemistry. 13 Pre-laboratory quiz. Solution chemistry. Group 1 and 2 cation standards. Laboratory Project 8 due. Group 2 and 3 cation standards. Laboratory Project 9 due. Group 3 and 4 cation standards. Laboratory Project 10 due. Group 4 and 5 cation standards. Laboratory Project 11 due. 14 Qualitative Analysis Cation Practical. Laboratory Project 12 due.
Laboratory Sessions 3-5, Experiment 2 Quantitative Analysis of Water Impurities I. INTRODUCTION Recently, Coca-Cola Company has been investigated for misleading consumers about the purity of Dasani bottled water [1]. As discovered by investigative reporting [2], Dasani water did not originated from natural springs, but from purified tap water. In fact, bottled water must meet less regulatory standards than major city tap water, since bottled water comes under the authority of the Food and Drug Administration and tap water is regulated by the Environmental Protection Agency. In this experiment, the acidity, calcium and magnesium concentration (water hardness), an the bicarbonate concentration (alkalinity) will be determined for bottled water and tap water. II. PROCEDURES A. Materials needed Chemicals ph 4.0, 7.0 and 10.0 buffers Ethylenediaminetetraacetic acid (EDTA) Hydrochloric acid Sodium hydroxide Sodium bicarbonate Phenophthalene indicator Methyl orange indicator Calmigate indicator Equipment burette burette stand and holder Pasco reader with appropriate adapter calcium selective electrode ph electrode In all procedures below, clean glassware implies that the glassware has been washed with soap and water and then rinsed with deionized water. If the glassware must be dried, then dry using paper towels. Four students will work in a group during this experiment. This group of four will be divided into pairs (Pair A and Pair B). The division of labor between the pairs is given in Table 1. B. Analysis of water: Part 1 ph of water The ph is a measure of the acidity of a solvent, with 0 being the most acidic (i.e., larger [H + ], where [X] is the concentration of species X) and 14 being consider basic (i.e., small [H + ]). The ph is determined using where a H+ is the activity of the H + ion. At low acid concentration, the activity a H+ of the hydrogen ion is approximately equal to the concentration [H + ] of the hydrogen ion. At ph < 3.0, only specialty plants and animals (known as acidophiles) can live in the water. The ph range that is optimal for most organisms is 6.5 8.2. Rainfall usually has a ph between 5 and 6.5 due to dissolve CO 2 and air pollutants (such as sulfur dioxide or nitrogen oxides). River water, on the other hand, has a ph between 7 and 8.5 because of the bicarbonate ions that result from dissolved limestone. Acceptable ph levels for most municipal drinking water standards is between 6.5 and 8.5. In this part of the experiment, the ph of deionized water, tap water and bottled water will be measured.
Table 1. Division of labor between Pair A and Pair B for this three week experiment. Day Experiments for Pair A Experiments for Pair B 1 B.1. ph electrode calibration C.1. Ca 2+ electrode calibration C.2. [Ca 2+ ] determination of B.2. ph determination of tap and bottled water tap and bottled water C.3i. EDTA standardization (2 runs) C.3i. EDTA standardization (2 runs) C.3ii. Total [Ca 2+ ] and [Mg 2+ ] in C.3ii. Total [Ca 2+ ] and [Mg 2+ ] in deionized water deionized water 2 C.3iii. Total [Ca 2+ ] and [Mg 2+ ] in C.3iii. Total [Ca 2+ ] and [Mg 2+ ] in tap water bottled water C.3iv. [Ca 2+ ] in bottled water C.3iv. [Ca 2+ ] in tap water 3 D.1. Calibration of HCl D.1. Calibration of HCl D.2. P alkalinity of tap water D.2. P alkalinity of bottled water D.3. T alkalinity of bottled water D.3. T alkalinity of tap water B.1. Calibration of ph electrode 1. Place 15 ml of the ph 4.0 buffer in one single clean dry 50 ml beaker and 15 ml of the ph 7.0 buffer in a second clean dry beaker. Calibrate the ph electrode following the instructions given in the techniques handout. 3. Once the ph electrode is calibrated, place the electrode in a beaker of deionized water until the electrode is ready for use. 4. Weigh 2.0 grams of KHP using weighing paper. (Be sure to record the exact weight to mg precision.) Transfer the KHP to a 100 ml clean volumetric flask (obtained from the stockroom). Using a pipette, rinse the weighing paper with a small amount of distilled water to ensure the transfer of all the solid. Add approximately 50 ml of deionized water. Cover the flask and shake until all of the solid KHP dissolves. Once the dissolution of KHP is complete, fill the volumetric flask to the 100 ml mark, cover and shake to obtain a homogeneous solution. 5. Pour 15 ml of this standard KHP solution into a clean and dry 50 ml beaker and measure the ph using the calibrated ph meter. Be sure to rinse the ph meter with deionized water after the measurement before returning the meter to the storage beaker of deionized water. 6. Using the same standard KHP solution, perform the ph measurement an additional four times. (Always rinse, clean and dry the 50 ml beaker before adding the next aliquot of standard KHP.) 7. Obtain the average and compare the measured ph with the calculated ph. Show the laboratory instructor your results to confirm the calibration. If the laboratory instructor verifies the calibration, your group can proceed to Step 8. If not, repeat Steps 1-3 to recalibrate the electrode and test the calibration using Steps 5 and 6. Once the laboratory instructor has confirmed the calibration, proceed
to Step 8. 8. Measure the ph of two separate aliquots of deionized water so that you have a basic standard. B.2. Determination of the ph of the water samples 1. Obtain from your instructor two water samples (one of city water and one of bottled water). Write down on your report sheet the name of the bottled water and the location where the tap water was obtained. 2. Using the calibrated ph electrode, determine the ph of the water samples. Use three separate aliquots of each sample for this study. C. Analysis of water: Part 2 Hardness of water Hardness is defined as the concentration of multivalent cations (e.g., Ca 2+, Mg 2+, Fe 2+, etc.) in water. Hard water cations form an insoluble compound with common soaps and detergents, leading to reduced cleansing ability and to the appearance of soap scum. The chemical reaction for the formation of soap scum is Similarly, boiler scale results in the equilibrium reaction of calcium carbonate and bicarbonate, namely This boiler scale, which requires heat to drive the reaction, forms on hot surfaces and, therefore, coats the inside of hot water pipe and forms a layer on heating elements. When boiler scale builds on the inside of hot water pipe, this scale reduces water flow. When it forms on heating elements, boiler scale reduces efficiency, since CaCO 3 (s) is not a strong heat conductor. Although hard water causes problems, it is not a health risk. The hardness of water is classified by the U.S. Department of Interior and the Water Quality Association as the concentration of ions in the water. The dominant ions in water are generally calcium, magnesium, lead and iron. However, since calcium ions predominate, the hardness is usually expressed in terms of calcium ion concentration with the values given in Table 2. The calcium concentration in deionized water will be determined, and the calcium and magnesium concentration in tap and bottle water will be measured. C.1. Calibration of calcium selective electrode Table 2. Hardness Classification of Water Ca 2+ concentration Classification 0 20 mg/l Soft 20 40 mg/l Moderately Soft 40 60 mg/l Slightly Hard 60 80 mg/l Moderately Hard 80 120 mg/l Hard > 120 mg/l Very Hard The calcium selective electrode will be stored in the hood in a beaker containing a CaCl 2 solution. Take one of these electrodes for your studies only after the standard solutions are prepared.
1. The calibration solutions for a calcium ion selective electrode are composed of dissolved calcium chloride. The high concentration standard is 1000 mg/l in Ca 2+, while the low concentration standard is 10 mg/l in Ca 2+. Determine the amount of CaCl 2 2H 2 O is needed to prepare 250 ml of the high concentration standard. Also calculate the dilution volume needed to prepare 100 ml of the low concentration standard. Show your calculations to your laboratory instructor. 2. Once the calculations are verified, prepare the high concentration standard solution using deionized water in a clean 250 ml volumetric flask, and the low concentration standard solution in a clean 100 ml volumetric flask. 3. Calibrate the calcium selective electrode using the two prepared standards. When not in use, the calcium selective electrode should be stored in a beaker containing the high concentration standard. 4. Determine the amount of CaCl 2 2H 2 O needed to prepare 100 ml of a 500 mg/l in Ca 2+ both by direct preparation and by dilution of the high concentration standard. Show your calculations to your laboratory instructor. 5. Prepare two separate 250 mg/l Ca 2+ solutions. Use direct preparation for Solution A and dilution of the high concentration standard for Solution B. Using three separate aliquots from each solution, measure the Ca 2+ concentration to verify the electrode calibration. Rinse the beaker used to hold the aliquots before each use. Rinse the electrode with deionized water before each measurement. 6. If the laboratory instructor confirms your results and, thereby, validates your calibration, your group can proceed to Step 7. Otherwise, redo Steps 1-3 and confirm the calibration using Step 5. Once the calibration has been verified by the laboratory instructor, proceed to Step 7. 7. Finally, to obtain a baseline, measure the calcium concentration in the deionized water used to prepare all samples in this experiment. This calcium concentration will need to be subtracted from the reading determined using the calcium selective electrode (potentiometric analysis). C.2. Potentiometric determination of the calcium concentration of the water samples 1. Using the calibrated calcium selective electrode, determine the calcium concentration in the water samples obtained in Section B. 2. Using Table 1 determine the hardness of the water. C.3. Determination of the hardness of the water samples using complexometric titration Disodium ethylenediaminetetraacetic acid (Na 2 EDTA or Na 2 H 2 Y, cf. Fig. 1) forms a one to one complex with divalent metal cations, which is one of the reasons why this compound is often used as a preservative for canned food. When aqueous H 2 Y 2- is added to a sample containing divalent cations X 2+, the following reaction occurs However, EDTA complexes with metals tend to lack color and, therefore, an indicator must be used to determine the end point of a titration using EDTA. Moreover, EDTA will complex all divalent cations. The
procedure here will determine the total concentration of both Mg 2+ and Ca 2+ and then, separately determine the Ca 2+ concentration. i. Standardization of the EDTA solution At ph 10, aqueous calmagite (HOC 10 H[N=NC 6 H 3 (OH)CH 3 ]SO 3 H) has a deep blue color. However, when complexed to a metal cation, the complex will change color to a red. However, calmagite does not form a strong complex and is displaced by EDTA. Thus, at the start of a titration, the solution will be reddish color. At the end point of the EDTA titration, indicating that all of the Ca 2+ and Mg 2+ are complexed with EDTA, the solution will again turn blue. The first step in the complexometric titration is the calibration of EDTA. 1. Calculate the mass of EDTA (molar mass = 372.24 g/mol) required to prepare 250 ml of a 0.01 M EDTA solution using deionized water. Show these calculations to your laboratory instructor. Once the instructor approves your calculations, prepare the solution in a 500 ml clean Erlenmeyer flask using a graduate cylinder to measure the necessary volume of water. 2. Prepare two burettes. Rinse the first burette several times with the EDTA solution and then fill. Rinse the second burette several times with the high concentration calcium ion standard. 3. Transfer 25 ml of the high concentration calcium ion standard to a 125 ml Erlenmeyer flask. Using a pipette, add 1.0 ml of ph 10 buffer and 2 drops of the calmagite indicator. Then place the Erlenmeyer flask under the burette containing the EDTA solution. 4. Titrate the standard calcium solution with the EDTA solution, swirling continuously. Near the end point, slow the rate of the addition EDTA. The last few drops should be added at 3 5 s intervals. The solution should change from wine-red to purple to sky blue at th end point. Record the final burette reading. 5. Repeat Steps 3 and 4 for one additional sample. Use these data to determine the exact concentration of EDTA. 6. The standard EDTA solutions should be stored in a clean, dry, labeled bottle provided by the laboratory instructor. ii. Determine the total Ca 2+ and Mg 2+ concentration in the distilled water 1. After the EDTA solution is standardized, empty the burette containing the calcium standard into the appropriate waste container. Then rinse the burette containing the calcium standard with deionized water several times and dry using a small amount of acetone and air. Place the deionized water into the clean and dry burette. 2. Using the burette containing the water sample, create a 25 ml aliquot in a cleaned and dry 125 ml Erlenmeyer flasks. To this aliquot, add 1 ml of the buffer and 2 drops of the calmagite indicator. 3. Titrate the aliquot with the standardized EDTA of Part A. Remember that as the end point is approached, the last few drops should be added at 3 5 s intervals to ensure that the end point is not missed. These data will be used to determine the total Ca 2+ and Mg 2+ concentration.
iii. Determination of the total Ca 2+ and Mg 2+ concentration in the water samples 1. After the EDTA solution is standardized, empty the burette containing the deionized water. Dry using a small amount of acetone and air. Place one of the water samples into the clean and dry burette. 2. Using the burette containing the water sample, create three 25 ml aliquots in three cleaned and dry 125 ml Erlenmeyer flasks. To each of these aliquots, add 1 ml of the buffer and 2 drops of the calmagite indicator. 3. Titrate each of the aliquots with the standardized EDTA of Part A. Remember that as the end point is approached, the last few drops should be added at 3 5 s intervals to ensure that the end point is not missed. These data will be used to determine the total Ca 2+ and Mg 2+ concentration. iv. Determination of the Ca 2+ concentration by complexometric titration 1. Using the burette containing the water sample, create three 25 ml aliquots in three cleaned and dry 125 ml Erlenmeyer flasks. To each of these aliquots, add 15 drops of 50% w/v NaOH solution and swirl to mix. 2. Wait 5 minutes (with intermittent stirring) to allow for the complete precipitation of Mg(OH) 2 (s). 3. Added three drops of hydroxynaphthol blue to act as an endpoint indicator. (Calmagite is not stable at the ph required to precipitate Mg 2+.) 4. Titrate each of the aliquots with the standardized EDTA of Part A. These data will confirm the calcium concentration determined using the calcium selective electrode. D. Analysis of Water: Part 3 Alkalinity of Water Alkalinity measures the ability of water to neutralize acids, at least to the equivalence point of carbonate or bicarbonate. When water passes over limestone (i.e., CaCO 3 ), the calcium carbonate dissolves increasing both the calcium concentration and the carbonate concentration. The carbonate anion reacts with water to form carbonic acid and bicarbonate, with equilibria equations given by where K a1 and K a2 are equilibrium constants for the reactions. The dominant species in water is determined by the ph of the water. If the ph is between 5.0 9.0, then the dominant species is bicarbonate (HCO 3 - ). At ph < 3.0, carbonic acid [H 2 CO 3 or CO 2 (aq)] dominates, while at ph > 11 carbonate (CO 3 2- ) is the primary species. For municipal and industrial water supplies, the alkalintiy ranges from 30 to 400 mg/l. Alkalinities of - 350 mg/l inhibit pipe corrosion, but alkalinities of > 500 mg/l and < 100 mg/l tend to increase pipe corrosion. Thus, high and low alkalinities tend to be avoided. The alkalinity of a sample is determined by titration with a strong acid. However, the value reported in the literature depends upon the indicator used to determine the end point. Two equivalence points exist in
the titration of carbonic acid with a concentration acid. The P alkalinity, otherwise known as the carbonate alkalinity, occurs around ph 8.2, which is close to the endpoint for phenophthalein. This endpoint indicates when all anions of the stronger bases have been mono-protonated. These anions, which are dominated by CO 3 2-, can also include OH - and PO 4 3-, which represent the strong bases also indicate the corrosiveness of water. The second equivalence point occurs when all of the bases contributing to alkalinity have been fully protonated. This point, called the T alkalinity (or bicarbonate alkalinity), occurs around ph 3.4 and gives the total alkalinity of the sample. Thus, the T alkalinity is a measure of the total buffer capacity of the water. The indicator used for this endpoint determination is methyl orange. If the only anions that exist in the water sample are the carbonate series of anion, then the T endpoint should occur at double the H + concentration as the P endpoint. If other anions exist, then the two endpoints will not be simply related. Generally, the alkalinity reported to state and national government water monitoring agencies is the T endpoint. In this experiment you will determine the alkalinity of deionized water, bottled water, and tap water. D.1. Standardization of HCl 1. Calculate the volume of 1.0 M HCl (l) required to create 500 ml of 0.015 M HCl (l). Show your results to your laboratory instructor for approval. 2. Once you calculations have been approved, prepare the 0.15 M HCl solution in deionized water using a graduate cylinder to measure volume. 3. Calculate the mass of Na 2 CO 3 that would be neutralized at the end point of methyl orange (i.e., approximately ph 3.4) with approximately 25 ml of the 0.015 M HCl. Show your calculations to your laboratory instructor for approval. 4. Once your calculations have been approved, use weighing paper to measure the required mass of Na 2 CO 3 to create the sample. A total of two samples will need to be prepared. Transfer each of the two samples to a separate 125 ml Erlenmeyer flask, and add 25 ml of deionized water to each sample and 3 drops of the methyl orange indicator. 5. Prepare a clean, dry 50 ml burette for titration. HCl will be the titrant. 6. Place a water sample below the burette and begin the titration. Slowly add the HCl titrant to the Na 2 CO 3 solution. As the color change from yellow to red slows with the addition of the titrant, gently warm the Erlenmeyer flask over an open flame for 1 2 minutes to evolve any of the CO 2 (aq) that has formed during the titration. 7. Continue to titrate the solution to the methyl orange endpoint. (We should note that the color change from yellow to red-orange is subtle and requires care for reproducibility.) 8. Repeat Steps 6 and 7 for the remaining sample. Use these data to calibrate the HCl solution. D.2. P alkalinity of tap water and bottled water 1. Label two clean, dry 125 ml Erlenmeyer flasks for the P alkalinity measurement. Pipette 25.0 ml aliquots of the water sample into each of the Erlenmeyer flasks. Added 2-3 drops of phenophthalene. 2. Ensure that at least 30 ml of standard HCl is in the burette for titration. If not, refill the burette.
3. Titrate the water sample to the phenophthalene endpoint. When the endpoint is reached (as signified by the change in color from pale pink to colorless), record the volume of the titrant. 4. Repeat the test for P alkalinity on the second water sample. D.3. T alkalinity of tap water and bottled water 1. Label two clean, dry 125 ml Erlenmeyer flasks for the T alkalinity measurement. Pipette 25.0 ml aliquots of the water sample into each of the Erlenmeyer flasks. Added 2-3 drops of methyl orange indicator. 2. Ensure that at least 30 ml of standard HCl is in the burette for titration. If not, refill the burette. 3. Titrate the water sample to the methyl orange indicator endpoint. When the endpoint is close (as indicated by a decreasing color change), gently warm the Erlenmeyer flask over an open flame for a minute. 4. Once the endpoint is reached (as signified by the change in color from yellow to red-orange), record the volume of the titrant. 5. Repeat the test for T alkalinity on the second water sample.
Name: Section: Grade Partner: Date: EXPERIMENT 2: PART 1. PRE-LABORATORY QUESTIONS 1. What is the acceptable ph range for municipal drinking water? 2. Why is the ph of rainfall usually slightly acidic? 3. Why is the ph of river water usually slightly basic? 4. If 2.0 grams of KHP, a monoprotic acid, is dissolved in 100 ml of water, what should the ph be? 5. What impurity in bottled water would lead to a ph of 7.3? 6. If KHP powder is dropped on your skin, what should you do?
Name: Section: Grade Partners: Date: EXPERIMENT 2: PART 1. REPORT SHEET Report all measurements and calculations to the correct number of significant figures and with the correct units. B.1. Results 1. Mass of KHP weighed using electronic balance: 2. What is the expected ph for the weighed KHP? 3. ph measurements of KHP: Aliquot 1 Aliquot 2 Aliquot 3 Aliquot 4 Aliquot 5 Average (with error) 4. ph measurement of deionized water: Aliquot 1 Aliquot 2 Average (with error) B.2. Results
1. ph measurement of tap water: Aliquot 1 Aliquot 2 Aliquot 3 Average (with error) 2. ph measurement of bottled water: Aliquot 1 Aliquot 2 Aliquot 3 Average (with error)
Name: Section: Grade Partner: Date: EXPERIMENT 2: PART 1. POST LABORATORY QUESTIONS Report all measurements and calculations to the correct number of significant figures and with the correct units. 1. What is the ph of a solution containing 1.25 g of KHP (204.22 g/mol) in 100 ml? 2. Describe, in detail, another method that can be used to determine the ph of KHP. 3. What is the ph for tap water? What is the ph of bottled water? Explain why the two are similar or different. 4. Why must the ph electrode be calibrated at two different ph values?
CHEMISTRY 114.1 PRELABORATORY QUIZ LABORATORY 2: PART 1 FALL 2008 Problem 1 (2 pts) If KHP is spilled on the counter top, what should you do? Problem 2 (4 pt) What is the ph of a solution made of 1.753 g of KHP (molar mass of 204.22 g/mol) in 100 ml of deionized water? Problem 3 (2 pts) What impurities in water would lead to a ph of 5.4? Problem 4 (2 pts) Why is the ph meter calibrated using two different ph buffers?