FURTHER EXPLORATIONS IN THE CHEMICAL WORLD CHM 103 LABORATORY MANUAL

Transcription

1 FURTHER EXPLORATIONS IN THE CHEMICAL WORLD CHM 103 LABORATORY MANUAL Scott A. Sinex Barbara A. Gage Department of Physical Sciences and Engineering Prince George's Community College 2005

2 ACKNOWLEDGMENTS The authors would like to credit our colleague Patricia Basili for showing them a better approach to utilizing laboratory activities to teach concepts in general chemistry. We also extend thanks to Indravadan Shah who has patiently prepared and troubleshot laboratory set-ups. We acknowledge our former colleague Margaret Tierney who was involved with CHM 103 for many years. ii

3 TABLE OF CONTENTS ACTIVITY Laboratory Protocol and Safety Using Excel for Graphing Data Use of the Analytical Balance Measurement Variation Sampling Variation The Accuracy and Precision of Volume Measuring Devices Analysis of Calcium by EDTA Titration to Assess Water Hardness Titrimetric Analysis of Antacid Tablets More Lights, Color, Absorption! Determination of the Rate Law for Crystal Violet/Sodium Hydroxide Reaction Investigating a Hypothetical Initial Rate Model of the Behavior of Crystal Violet Ion (CV + ) Absorbance Spectrophotometric Determination of the Equilibrium Constant for the Formation of a Complex Ion An Investigation of Electrochemical Reactions Conductometric Measurements Characterization of a Monoprotic Weak Acid by Potentiometric Titration Discovering Intramolecular Interactions Studying Vibrations in Molecules PAGE iii

4 NOTES TO STUDENTS Welcome to CHM 103 General Chemistry II Laboratory. You will be building on concepts uncovered in General Chemistry I using the approach encountered in CHM 101 at. Many of these activities are designed to introduce concepts before they are covered in lecture/discussion and so they may not be graded. However, it is important that you do all parts of each activity to get the full benefit. The course covers measurement and its error, spectrophotometry and its use as an analytical tool, and electrochemical measurements. Many of the methods involve non-visible means of tracking chemical reactions. We also venture into looking at molecular structure and IR spectroscopy as a tool for identifying functional groups and use later in organic chemistry. The Physical Sciences and Engineering Department provides numerous resources, such as guides to using ChemSketch, Chime, the TI-graphing calculator, Excel, and PowerPoint plus links to freeware to assist you in understanding topics and skills that are part of this course. We recommend the following web sites: Physical Sciences and Engineering Department homepage and click on Resources for Students Dr. Scott Sinex s CHM 103 page The manual, as individual pdf files or the entire manual in one file (go to bottom of page), plus interactive support materials that accompany activities are available at iv

5 LABORATORY PROTOCOL AND SAFETY The chemistry laboratory is probably safer than your own kitchen. We have a number of requirements for you to follow while in the laboratory. 1. Safety glasses must be worn over the eyes at all times (available in laboratory). 2. Closed-toed shoes must be worn on your feet at all times (NO sandals or open-toed shoes). 3. Your presence at the beginning of each laboratory is mandatory. Lateness will not be tolerated and as a result you may be blocked from the laboratory with loss of points. 4. NO food (including gum and candy) or drinks are allowed in the laboratory. 5. You are responsible for reading the instructions for each activity BEFORE coming to the laboratory. 6. We strongly suggest protective clothing be worn (lab coat or apron). 7. Any common equipment must be cleaned and placed back where found. 8. Your laboratory station must be left clean when you are finished working. 9. Failure to abide by the safety requirements and rules may result in your removal from the laboratory with loss of points. Here are some general safety rules to follow when handling chemicals, more specific instructions are typically given in each activity and will be reviewed by the instructor at the beginning of each laboratory period. 1. All solid waste is to be placed in the trash cans or appropriate labeled containers. 2. Liquid waste will be placed in labeled containers in the front of the laboratory unless otherwise instructed. 3. All broken glassware is to be placed in the glass disposal container. 4. Long hair must be tied back to avoid accidents with open flames and chemicals. 5. Open flames (Bunsen burners) must NOT be left unattended. Be sure gas valve is shutoff when you are finished using. 6. Flammable liquids must never be heated over open flames. 7. If you get any chemical on your skin, IMMEDIATELY start to rinse the chemical off 1

6 with cold water, then inform your instructor. 8. In the event of a chemical spill or thermometer breakage, IMMEDIATELY inform your instructor. 9. Know where the exits from the laboratory are located. Location of exits: 10. The laboratory is equipped with the following items listed in the table. You should know their purpose or function and their location in the laboratory. ITEM PURPOSE/FUNCTION LOCATION fire extinguisher fire blanket safety shower eye wash first aid kit flume hoods sodium hydrogen carbonate (powder and solution) chemical/paper/wood fires smother fire on burning human or small surface burning human or large chemical spill rinse chemical from eye cuts/burns use with toxic or noxious chemicals for acid or base burns on skin 11. Avoid taking excess amounts of chemicals as it is wasteful. Any excess chemicals removed from containers must never be placed back in the containers (avoids contamination). Share the excess with another student or place in appropriate waste container. 12. Never heat glassware that contains a crack. 13. Be extremely careful when handling hot glassware and apparatus (ring stands and burners). Use appropriate heat handling equipment. In the event of any laboratory accident (cut, burn, splash, spill, or fire) your instructor is the first person to inform. Your instructor will assess the need for any further actions. 2

7 Using Excel for Graphing Data Whether producing a graph using technology or by hand, it should contain the following: Ionic Conductance Ionic Conductance as a function of Temperature Temperature (Celsius) 1. axes with consistent scales and variables (with units) labeled; 2. data points plotted with a symbol; 3. title (written as Y vs X); 4. appropriate line or curve drawn through data; 5. graph should fill the page, especially hand-drawn graphs. For any graph in Excel, use a XY scatter plot that shows the data points. The scatter plot is the only option where the x-axis is plotted as a scaled variable. Chart sub-types: Scatter plot showing only data points Points connected using smoothing, (which uses a cubic spline) Points connected using lines (seldom applicable for scientific data) 3

8 The scatter plot using smoothing is very useful for absorption spectra. 1 Absorption Spectrum absorbance wavelength, nm For calibration curves, the regression should include the origin and the line should go through the origin (0,0). Absorbance Calibration Curve A = mc Concentration (M) To add a regression line to a plot: From the Chart menu, select Add Trendline; for Type, select a linear regression and then from Options select Set intercept = 0, Display equation on chart, and Display R-squared value on chart. The Forecast feature will extend (extrapolate) the line pass the data limits. For complete instructions for using Excel see 4

9 Name Partner(s) NONE Section Date USE OF THE ANALYTICAL BALANCE 1. What is mass? How is it determined? PRE-LAB QUERIES 2. How are mass and volume of a substance related? OBJECT This activity introduces you to the use of the analytical balance, an instrument that allows you to determine the mass of objects to g or 0.1 mg. PROCEDURE A. Determination of the mass of two objects (crucible and cover) individually and together 1. Mass each object individually three separate times on one analytical balance. Record the masses in the data table. 2. Mass both objects together three separate times on the same balance. Record the masses in the data table. 3. Calculate the sum of the masses of the two objects and record the result in the results table Compare the results of the sum of the two objects to the mass of them together; they should agree to within 0.3 mg. The replicate mass determinations of any object should also agree within this limit. One means to check this replication is by calculating the range, which is the difference between the high and low values. The range should be no larger than 0.3 mg. Repeat the massings two more times if needed to replicate the mass within the prescribed limit. Further Explorations of the Chemical World 5

10 B. Determination of the volume of a drop. 1. Mass a 10 ml beaker. 2. Place water in a clean buret making sure there is no air in the tip. Record the initial volume of the water in the data table. Deliver exactly 50 drops of water from the buret into the pre-massed beaker. Record the final volume of the water in the buret. 3. Mass the beaker with the water. 4. Measure the temperature of the water. 5. Calculate the mass and volume of one drop of water and record the values in the results table. 6. You can calculate the volume of a drop in another way using the specific volume (V s ) (the volume occupied by one gram of water corrected for buoyancy of air) Using the mass of a drop of water and table below, calculate the volume of a drop of water using the following formula: V drop (ml) = M drop (g) x V s (ml/g) Temperature ( o C) Specific Volume (ml/g) Further Explorations of the Chemical World 6

11 C. Mass of a fingerprint 1. Clean a watch glass with a kimwipe or microwipe. 2. Using a piece of paper or tongs, mass the watch glass. 3. Now touch the watch glass with your finger ten to twenty times, each time touching a clean part of the watch glass. Be sure to record the number of times you touch the glass. 4. Remass the watch glass after fingerprinting. 5. Calculate the mass of your fingerprint. D. Determining balance accuracy using a standard mass 1. Obtain an object of standard mass from your instructor. DO NOT TOUCH the mass with your fingers because substances in your finger oils may react with the object and modify its mass! Mass it three times on the analytical balance and record the masses in the appropriate data table. 2. Calculate the percentage error in your average mass using the formula below. experimental mass (average) - theoretical mass % error = x 100 theoretical mass The theoretical mass is the mass of the standard. For this lab, if the object is marked with a 10, the mass is g. Further Explorations of the Chemical World 7

12 DATA AND RESULTS A. Determination of the mass of two objects (crucible and cover) individually and together DATA RESULTS Trial Object 1 Object 2 Objects 1 and 2 Massed Together Sum of Object 1 and Object 2 Difference (mass sum - mass together) RESULTS Average Range B. Determination of the volume of a drop. DATA Mass of 10 ml beaker Mass of 10 ml beaker and water Volume reading on buret before dispensing water Volume reading on buret after dispensing water Number of drops Temperature of water Mass of water RESULTS Volume of water Mass of 1 drop Volume of 1 drop (from buret readings) Volume of drop (from V s calculation) Further Explorations of the Chemical World 8

13 How do the results for the volume of a drop compare for the two methods used? C. Mass of a fingerprint Mass of watch glass DATA Mass of watch glass and fingerprints Number of fingerprints Mass of fingerprints RESULTS Mass of 1 fingerprint Do your fingerprints contribute mass when determining the mass of an object to the nearest 0.1 mg? Explain. D. Determining balance accuracy using a standard mass Mass of standard - trial 1 DATA Mass of standard - trial 2 Mass of standard - trial 3 Average mass RESULTS % error Further Explorations of the Chemical World 9

14 CONCLUSIONS How do your results from the volume of a drop and single fingerprint mass (parts B and C) compare to three other students in the laboratory? Make sure you can support your statements! The percentage error allows you to access the accuracy or closeness to true value. It assumes we have a standard with a known mass. How does your balance error compare to three other students? POST-LABORATORY QUESTIONS 1. Explain the difference between data and results. 2. The percent error can be positive or negative. Explain what the sign of the error tells you about the error. 3. Why is it important not to touch the standard mass? Further Explorations of the Chemical World 10

15 Name Partner(s) Section Date MEASUREMENT VARIATION OBJECT This activity focuses on the variability in measurements of a property and explores methods of expressing the variation. Let's explore! PROCEDURE 1. Measure the height of five students in the laboratory. Use the heights of the four students in your group and the height of one member of another group. Record the data in the table below. 2. Have the same five students, independent of each other, measure the height of an object designated by the instructor. Record these data in the table below. 3. Calculate the average (mean) and range (high value - low value) for both sets of measurements. Student Student's Height Object's Height AVERAGE RANGE 4. Consider the two sets of data above. Describe and explain as many differences as you can between the two sets of measurements. What information does the range provide for each set? 11

16 We want to investigate the variability that comes from making multiple measurements of an object. For the purposes of this lab, we will simulate that variability by taking five similar samples of the same object and assume that the measurements are repeat measures of the same item. 5. Obtain three sets of washers, making sure that one set is labeled with an A, one with a B, and the last set with a C. Measure the diameter of the five washers in each set with a RULER. Record the data in the table below. Calculate the average and range for each set. Measurement Set A Set B Set C AVERAGE RANGE 6. Describe and explain any differences between the sets. 7. Graph the results above on a scatter plot with the set as the independent variable. See for information on how to set up the graph. Do the results provide any additional information about the three sets? 12

17 8. Calculate the deviation for each measurement (trial) in each set. Deviation = trial - average Calculate the sum of the deviations (Σdeviations) and the average of the deviations for each set. Σdeviations = dev 1 + dev 2 + dev dev n Σdeviations average of deviations = number of samples Measurement Set A Set B Set C Σdeviations AVERAGE Is there a difference between a positive and negative deviation? Explain. Remember that we are interested in how well we can reproduce a measurement (that is, how confident we are of the value). Does the average of the deviations tell you anything about the precision of the individual measurements? Explain. 13

18 9. As you can tell, the Σdeviations and average of deviations will be zero or close to zero if you include the sign of the deviation in the calculations. This does not give you any meaningful information about the precision of the measurements. How can we get more valuable information from the individual deviations? Removing the sign of the deviation and looking at the magnitude of the deviation without the direction will better represent the amount of deviation. We can accomplish this by taking the absolute value of each of the deviations; however, statisticians approach the problem by squaring the deviations. Square each of the deviations for the three sets to calculate the (deviation) 2. Record the values in the table below. Calculate the sum of the squared deviations for each set, Σ(deviation) 2. Measurement Set A Set B Set C Σ(deviation) 2 How does the Σ(deviation) 2 compare to the Σdeviations? 10. A useful statistic is the standard deviation, σ: σ = ( deviation) n-1 2 where n is the number of trials. 14

19 Calculate the standard deviation for each of the three sets. σ Set A Set B Set C Which set has the largest standard deviation? Why? Which set has the smallest σ? 11. Suppose you made a repeated measurement of an object that is 2255 cm long and determine that your standard deviation is 5.2 cm. Someone else making measurements of another object that is 543 cm has a σ = 2.2 cm. Is it fair to say that the second person s measurements are more precise than the first person s since the σ is smaller? A comparison of the standard deviation is only fair if the sets have approximately the same mean. To remove this problem we can normalize the standard deviation to the mean. This new parameter is called the coefficient of variation (CV) (with an older name of relative standard deviation (RSD). We often use the percent coefficient of variation expressed by the equation below: σ % CV = x 100 % mean (or average) The % CV tells you something about the measurement precision or how reproducible a measurement is. Calculate the % CV for each set. Use the average from question 5. %CV Set A Set B Set C Which set has the smallest % CV? Is it the one with the smallest σ? 15

20 12. Now, let s explore what happens to the deviations when we use a device that provides more information about a length. Repeat the measurement of diameter for set A with a pair of calipers. Ask the instructor for assistance if you have not used calipers before. Calculate the deviations, σ, and % CV. Trial Diameter deviation deviation squared σ = %CV= AVERAGE Σ= How do the σ and % CV for the caliper measurements compare to values for the ruler measurements? What does this tell you about precision using the different devices? 13. In the pharmaceutical industry, multiple measurement of a chemical parameter, such as %CaCO 3 in an antacid tablet, is often performed. Is variation expected? Is variation desirable or not? Explain. If large variations occur, what are possible causes? 16

21 Name Partner(s) Section Date SAMPLING VARIATION Object The object of this activity is to determine the variation and causes of variation that are introduced during sampling from a large mixture. Pre-Lab Queries In the previous lab you investigated variations in repetitive measurements of discrete objects. Often, scientists are required to sample from mixtures where the sampling might impact the results. To start an investigation of sampling variation, answer the following questions. Do you always get the same number of raisins in a scoop of Raisin Bran cereal or the same number of pretzels in a handful of party mix? Explain. Is obtaining a representative sample from a heterogeneous mixture different from or the same as obtaining a similar sample from a homogeneous mixture? Explain. Procedure You will work in pairs or individually for this activity. In the previous activity, Measurement Variation, you learned how to examine variation by determining deviations, standard deviation (σ), and percent coefficient of variation (%CV). We can apply the same statistics to determine how the actual sample and sampling technique contribute to variation. Using the statistics, we can also consider means of minimizing sample variation. Sand and gravel are the state of Maryland s most valuable mineral resources. These materials involve big money for Prince George s County and other regions of southern Maryland. Gravel is defined as any sedimentary material with a particle size greater than 2 mm. The category of 17

22 sand encompasses particles from 0.63 to 2 mm. Examine the sand and gravel mixture available in the laboratory but do not disturb the mixture. Your task in this activity is to determine the % by mass of sand in the sand and gravel mixture. % sand = mass of sand x 100 mass of sample Considering the laboratory mixture, is sampling the mixture going to present difficulties? Explain. What problems do you anticipate in obtaining accurate data to complete your task? By the way, your task requires you to make this analysis with great precision. Will you rely on a single sample? Explain. 1. Using the small scoop, obtain a sample of the sand/gravel mixture. Be sure you do not stir or significantly disturb the total mixture. Determine the mass of the mixture and record the data in the table provided. Use the balances in the lab (mass to nearest 0.01 g), not the analytical balances. 2. Separate the mixture using a sieve and mass the sand, recording the value. 3. Calculate the % sand in the mixture. 4. Repeat the sampling and separating with 4 additional samples. 5. Repeat steps 1-4 using the large scoop rather than the small scoop. 6. When you have completed all analyses, calculate the following statistics for your five percentages using the small scoop and the 5 using the large scoop. Record in the appropriate table. a. average (mean) b. standard deviation (σ) c. % CV 18

23 Using your TI-83 graphing calculator, enter your % sand results into L 1 (small scoop) and L 2 (large scoop). Use [STAT] and 1-Var Stats to get the mean and sample standard deviation, S x. 7. Enter the % sand results for your group in the Excel spreadsheet. Give your group a code name or use your initials to identify your group. Proofread the values entered! 8. Once the entire class has entered the percentages, the instructor will post the Excel file with the class data at a web address that will be written on the board. Write the address on the line below before leaving class. You will find that the posted file has the mean, standard deviation, and % CV calculated for each group and for the class as a whole. Use this information to answer the questions that follow. Check your group results! For the Group Results (n=5) Variation Analysis 1. In getting a scoop of sample, the mass of the sample collected varied. Does this influence the results? Would making sure that the scoop contents were level influence the results? Explain 2. Is there variation in the % CV among the different groups in the class using the small scoop? Explain. What does this tell you about the sampling techniques for each group? Do you find the same variation in the group results for the large scoop? Explain. 19

24 3. Which group had the largest % CV and which had the smallest % CV? Did the same groups have the largest and smallest values for the large scoop? What does this imply about the sampling technique of the group members? For the Pooled Results of the Class 4. What is the number of trials of this pool? n = 5. How does the class % CV for the small scoop compare to the % CV for the large scoop? Is this what you would have anticipated for the % CV values? Explain why or why not. 6. Discuss the factors that might lead to large group % CV values. 7. Does increasing the number of trials influence the precision of the analysis? Explain. 20

25 8. Suggest ways in which the precision of the results could be reduced for: a) each group; and, (b) for the class. 9. What is the cause(s) of the variation in this activity? Explain. 21

26 Small Scoop Data and Results Trial 1 Mass of container Mass of container and mixture Mass of container Mass of container and sand % sand Mean Standard deviation % CV 22

27 Large Scoop Data and Results Trial Mass of container 1 Mass of container and mixture Mass of container Mass of container and sand % sand Mean Standard deviation % CV 23

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29 Name Section _ Partner(s) NONE Date THE ACCURACY AND PRECISION OF VOLUME MEASURING DEVICES PRE-LAB QUERIES 1. Here are the masses of five paper clips: 5.123g; 5.117g; 5.124g; 5.120g; and 5.126g. Suppose you were asked, What is the mass of a paper clip? What is an appropriate statistical response? Include information on variation. 2. In the Use of the Analytical Balance activity you determined the volume of a drop of water by two different measurements methods. Three quantities were determined: mass of water; volume of water; and number of drops. Rank the three quantities in order of least to greatest error in their measurement and explain your ranking. OBJECT This activity explores the statistical variation in multiple measurements of volume and compares different volume measuring devices. The volume delivered by each device will be massed on the analytical balance. PROCEDURE All masses in this activity should be recorded to the nearest 0.1 mg. 1. Obtain and clean a buret, 10-mL pipet, and 25 ml graduated cylinder. 2. Label and pre-mass three plastic jars with lids. Be sure one jar is labeled with an A, one with a B, and one with a C. 3. Measure exactly 10 ml of water into jar A from the buret. Be sure to use the buret correctly. Mass the jar with the lid on to avoid evaporative loss of water. Repeat this procedure four more times adding the water to the same jar and obtain the mass after each additional of water. Record the masses (to the nearest 0.1 mg) in the table below. Record the temperature of the water to the nearest 0.1 o C 25

30 4. Repeat step 3 with jar B dispensing the water from a pipet. 5. Repeat step 3 using jar C and dispensing the water from the graduated cylinder. DATA T water = Trial Volume of water (ml) Mass (g) Jar A (buret) Mass (g) Jar B (pipet) Mass (g) Jar C (grad. cyl.) RESULTS 1. For each volume measuring device, calculate the mass of water for each 10 ml increment added in the five trials for each measuring device. Use the graphing calculator or spreadsheet program to compute the mean and standard deviation. Calculate the percent coefficient of variation (% CV). Record your values in the results table. 2. Using the density of water at the temperature recorded (available in the CRC Handbook of Chemistry and Physics) or the specific volume of water from the activity Use of the Analytical Balance, calculate the mass of ml of water. This will be the standard mass. Determine the percent error for each device using the formula below and record the results in the results table: %error = mean mass for device - standard mass x 100 standard mass Attach your calculations for standard mass and % error on separate sheets! RESULTS TABLE 26

31 Trial 1 Mass of Water (g) Buret Mass of Water (g) Pipet Mass of Water (g) Graduated Cylinder Mean s % CV % Error 27

32 CONCLUSIONS 1. Are there significant differences in the means of the three volume measuring devices? Explain using your data. What might account for these variations? 2. The percent coefficient of variation allows you to compare the amount of variability in the multiple measurements among several devices. This gives you an idea of the precision of each device since precision is the reproducibility of a measurement. Based on your data, rank the three volume measuring devices from most to least precise and justify your rank. 28

33 3. The calculation of % error provides you with information on the accuracy of each device. The accuracy is how close a measurement is to the true value. Based on your data, which device is most accurate? least accurate? Explain with justification. 4. Does the ability or lack of ability to use each device influence the error in any way? Explain. 29

34 POST-LAB QUESTIONS 1. A scatter plot of the masses for each volume measuring device is a method of determining any bias in the devices and their accuracy and precision. On graph paper, generate a scatter plot using the standard mass as your reference value. Show the three volume measuring devices as outlined in the sketch below. For each device, plot (in a line parallel to the mass axis) the five trials and the average. Measured Parameter Scatter Plot A B C Device or analyst Does the interpretation of this plot agree with your earlier conclusions? 2. The buret and pipet are volumetric glassware designed to deliver (TD) a specified volume of liquid. The graduate cylinder is designed to contain (TC) a specified volume. Does this difference in design affect the results? Explain why or why not. What type of volume measuring glassware is a volumetric flask? Explain. 30

35 Name Section Partner(s) NONE Date ANALYSIS OF CALCIUM BY EDTA TITRATION TO ASSESS WATER HARDNESS PRE-LAB QUERIES 1. What properties of a metal ion might be useful for the analysis of that ion? 2. We often speak of water samples as being hard or soft. What does this mean? OBJECT In this activity, you will perform a titrimetric analysis of Ca 2+ in a standard solution and water samples using the complexometric reaction of EDTA with metal ions. INTRODUCTION One measure of water quality is hardness which is defined by the amount of Mg 2+ and Ca 2+ (and sometimes Fe 3+ ) ions in a given amount of water. The presence of calcium and magnesium ions poses no health hazard but water hardness is of particular concern because the reaction of these ions with soap (sodium salt of a large fatty acid) produces an insoluble product we know as soap scum. This scum is abrasive and may weaken clothes fibers as they move against each other when worn. It will also be deposited on hair and skin when soap is used as a cleaner in hard water. Modern detergents do not produce the same degree of insoluble product and have effectively replaced soap for most products including hand cleansers, bars, shampoos, and laundry products. The most common source for Mg 2+ and Ca 2+ in water is carbonate compounds in Earth materials. As a result, hardness is often expressed as parts per million of CaCO 3 (ppm) by mass. A hardness of 100 ppm would be equivalent to 100 g of CaCO 3 in 1 million grams of water (10 3 L) or 0.1 g in 1.0 L or 100 mg in 1.0 L. 31

36 Titration procedures provide relatively inexpensive means for the analysis of different substances. These titrations are based on chemical reactions with completion points that can be monitored by some visible change in the reaction systems. As you are aware, the most common system involves acids and bases with the use of an indicator that changes color as the system moves from an acidic to basic composition. Calcium and magnesium ions can be measured through reaction with a chelating agent EDTA (ethylenediaminetetraacetic acid). This molecule has four carboxylic acid (~COOH) group sites and two nitrogens, all of which have lone pairs of electrons. The EDTA molecule can form a complex with as many as six sites on a particular cation like Ca 2+. These EDTA complexes are generally very stable are always in 1:1 (metal:edta) molar ratios. O OH H 2 C C OH O C CH 2 N CH 2 CH 2 N CH 2 C O Structure of EDTA in acid form HO C CH 2 OH O Structure of EDTA Structure of EDTA Complex In this activity you will be titrating Ca 2+ in a standard Ca 2+ solution or water samples with EDTA. Both Ca 2+ solutions and EDTA are colorless so an indicator is needed to signal the reaction completion. The indicator of choice is Eriochrome Black T which forms a wine-red complex with Mg 2+. A very small amount Mg 2+ will be bound to the indicator through most of the titration. When all of the Ca 2+ has reacted with EDTA, the Mg 2+ in the indicator will react with the EDTA. The indicator then returns to its acidic form which is a sky-blue and signals the end of the process. To maximize the EDTA-Ca (or Group IIA ion) complex formation and minimize formation of other metal complexes, the ph for the reaction system is set at 10 using an NH 3 -NH 4 + buffer. This keeps EDTA (H 4 Y) mostly in a half-neutralized form, H 2 Y 2-. Below are the reactions that occur during the titration where H 3 In is the general formula for the Eriochrome Black T. 32

37 During titration: H 2 Y 2- (aq) + Ca 2+ (aq) CaY 2- (aq) + 2 H + (aq) H 2 Y 2- (aq) + Mg 2+ (aq) MgY 2- (aq) + 2 H + (aq) At end point H 2 Y 2- (aq) + MgIn - (aq) MgY 2- (aq) + HIn 2- (aq) + H + Wine-red sky-blue The indicator reaction has to occur after the free Ca 2+ or any free Mg 2+ react. Why? What would happen if the indicator reaction occurred before this point? The equilibrium constant for the formation of Ca-EDTA has a larger value (~ x100) than for the production of Mg-EDTA. Therefore, the Ca 2+ is titrated first and then the Mg 2+ reacts. A small amount of magnesium ion will be added to the calcium solutions you use to generate the indicator changes. This amount is so small that it does not impact the analysis of the calcium. Standardization of EDTA PROCEDURE 1. Mass a sample of CaCO 3 between 0.30 and 0.32 g to the nearest milligram. Record the mass. Quantitatively transfer the CaCO 3 to a 250 ml beaker. Add 25 ml of distilled water. CAREFULLY and SLOWLY add 2 ml of 6 M HCl to the mixture in the beaker and mix. If the solution remains cloudy, add another drop of HCl and mix. Repeat if necessary until the solution is clear. Cover the beaker with a watch glass and heat the solution to boiling to remove CO 2. When cool, carefully transfer the solution, using a funnel, to a 250 ml volumetric flask. Rinse the beaker several times with small portions of distilled water and transfer the rinse to the volumetric flask. Rinse the funnel in the same manner several times. Fill the volumetric flask with distilled water until the meniscus just rests on the mark on the flask neck. Seal the flask and mix the contents by inverting times and shaking over a period of several minutes. 2. Obtain a ring stand with buret clamp and one buret. Clean the buret with a small amount of EDTA solution and drain. Fill the buret with EDTA Pipet a 25 ml aliquot of Ca 2+ solution into a 250 ml Erlenmeyer flask. Add 5 ml of ph 10 buffer and 2 drops of Eriochrome Black T indicator. Mix well. What color should the solution be at this point? Titrate the sample in the flask with the EDTA, slowing additions as you near the endpoint. Record the volume required to titrate the calcium sample. Keep the first sample as a color reference. 33

38 5. Prepare a second 25 ml portion of Ca 2+ solution and repeat the titration recording all data. If the volumes of EDTA agree to within 0.4 ml, proceed to the next section. If they do not agree, repeat the titration procedure until you have two sets of EDTA volumes that agree to within 0.4 ml. Determination of Water Hardness 1. Pipet exactly ml of a water sample provided into a 250 ml Erlenmeyer flask. Add 5 ml of ph 10 buffer and 2 drops of Eriochrome Black T indicator. Mix well. Titrate the sample with the same EDTA used in the previous section. Record all volumes. 2. Prepare and titrate 2 additional water samples recording the volume of EDTA used to just reach the end point. Calculations 1. Determine the moles of CaCO 3 (moles Ca 2+ ) in the volume tric flask and each ml aliquot titrated. 2. Using the volume of EDTA required to react with Ca 2+, calculate the molarity of EDTA and the average EDTA molarity. Use the best two trials to determine the average. 3. Using the volume and average molarity of EDTA required to titrate each water sample, determine the moles of Ca 2+ in each sample. Use these results to compute the moles of Ca 2+ per liter and grams of CaCO 3 per liter. From the g CaCO 3 /L, determine the hardness in ppm (mg CaCO 3 /L of water). 34

39 DATA AND RESULTS Preparation of Ca 2+ Solution Mass of CaCO 3 Moles of CaCO 3 Moles of Ca 2+ in ml aliquot Standardization of EDTA Initial buret reading Trial 1 Trial 2 Trial 3 Final buret reading Volume of EDTA Molarity of EDTA Average EDTA molarity Work: 35

40 Water Hardness Unknown Code Volume of water Trial 1 Trial 2 Trial 3 Initial buret reading Final buret reading Volume EDTA used Moles EDTA Moles Ca 2+ in sample Moles Ca 2+ / L Grams CaCO 3 / L Hardness (ppm) Average hardness % CV Work: CONCLUSION How do your results compare with at least two other students? Who has the best precision? Explain with justification. 36

41 POST-LAB QUESTIONS 1. What is the geometry of the Ca 2+ in its EDTA complex? 2. The CaCO 3 used to standardize the EDTA was a primary standard (very high purity). How would the results for the molarity of EDTA be influenced if the CaCO 3 was impure? Explain. 3. Hardness is traditionally expressed in mg CaCO 3 /L of water. Suppose that a water sample you analyzed contained mainly Mg 2+, how would the calculation of hardness be affected? Explain. 4. A g sample of CaCO 3 is dissolved in HCl and diluted to 250 ml in a volumetric flask. A ml aliquot of the sample requires ml of an EDTA solution for titration. What is the molarity of the EDTA? Show work. If a ml well water sample requires ml of the EDTA solution above, what is the water hardness of the well water? Show work. 5. Go to and describe a region in the United States with a low water hardness and one with a high water hardness. 37

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43 Name Section Partner(s) NONE Date TITRIMETRIC ANALYSIS OF ANTACID TABLETS PRE-LAB QUERIES 1. Predict the reaction products for the following antacids with stomach acid, HCl. Remember all neutralization reactions are double displacement reactions. NaHCO 3 (cheap antacid) CaCO 3 (Tums) Mg(OH) 2 (Milk of Magnesia) 2. What is the relative water solubility of each compound above? 3. How would you know when any of the above reactions were completed? OBJECT In this activity, antacid products will be analyzed by an indirect titrimetric procedure. From the analysis, the percentage of antacid in the tablet and percentage error compared to the label value will be calculated. The effectiveness of various brands will be compared. INTRODUCTION Antacids are compounds that act as bases to neutralize stomach acid, hydrochloric acid (~0.02M). Most antacids such as calcium carbonate and magnesium hydroxide are water insoluble and thus difficult to analyze by direct titration. You will want to review the background on acid-base titrations covered in the CHM 101 laboratory manual or your textbook. You should know the following terms: equivalence point; end point; titrant; standard solution; and, indicator. Because of the insolubility and any carbon dioxide production, antacids are analyzed by an 39

44 indirect method or what is known as a back titration. The antacid is allowed to react with a known amount of hydrochloric acid and then the excess hydrochloric acid is titrated with sodium hydroxide. The amount of acid reacted or neutralized by the antacid is determined by difference. moles neutralized = moles initial - moles excess If carbon dioxide is produced in the reaction of an antacid, the carbon dioxide must be removed before titration. The carbon dioxide gas can be removed by boiling, which shifts the equilibrium in the reaction below to the left CO 2 + HO 2 HCO 3 (aq) + H (aq) If CO 2 is left dissolved, the H + produced would be titrated by sodium hydroxide. To visualize the progress of the reaction, an indicator will be added. Thymol blue has the colors listed below at the given ph's. ph Color >2 pink/red 2-8 yellow >8 blue Since the titration is between a strong acid and strong base (equivalence point ph of 7), the equivalence point is located in a very steep region of the titration curve. The indicator changes a little above the equivalence point, but with no appreciable error in the volume of titrant required. To allow the comparison of the effectiveness of antacid brands, the amount of hydrochloric acid neutralized is normalized to the mass of tablet. The larger this normalized value, the more effective the antacid. PROCEDURE 1. Into a pre-massed massing boat add one crushed antacid tablet. Determine the mass of the tablet added to the nearest 0.1 mg. Quantitatively transfer the tablet to an Erlenmeyer flask. 2. Clean two burets. Fill one with the standardized HCl and the other with standardized NaOH. Be sure to record the concentrations of the acid and base. 3. To the crushed antacid, add exactly 25 ml of standard HCl and approximately 25 ml of distilled water. Gently swirl to dissolve the antacid. Heat the solution to boiling and allow it to boil for 5 minutes to insure CO 2 loss. Allow the flask to cool before adding 40

45 indicator and starting the titration. 4. Add 2-3 drops of thymol blue indicator to the cooled flask and titrate the excess HCl with standard NaOH. The end point is a pale blue color (yellow to blue) but will occur in the presence of suspended starch (cloudy) used as a tablet binder so careful observation is required. The blue color of thymol blue does not change once obtained (unlike phenolphthalein). Save the first titration of each brand as a color comparison for the endpoint. Repeat the preparation and titration on 2 additional tablets of the same brand. 5. Analyze a different brand of antacid tablet following the procedure above. 6. For all titrations complete the data and results table. For each brand of antacid, calculate the percentage of antacid compound in a tablet and percentage error using the formulas below: g antacid by titration % antacid = x 100 g tablet g antacid by titration - g antacid by label % error = x 100 g antacid by label 41

46 DATA AND RESULTS Concentration of standard HCl Concentration of standard NaOH DATA Brand: Antacid Compound: Amount on Label: Mass of Massing Boat Mass of Massing Boat and Antacid Tablet Volume of HCl Added Volume of NaOH Mass Antacid Tablet Moles HCl Initially Added Moles HCl Excess Moles HCl Neutralized Moles HCl Neutralized Per g Antacid Tablet Moles of Antacid in Tablet Mass of Antacid in Tablet % Antacid in Tablet % Error (compare to label amount) Trial 1 Trial 2 Trial 3 RESULTS 42

47 DATA Brand: Antacid Compound: Amount on Label: Mass of Massing Boat Mass of Massing Boat and Antacid Tablet Volume of HCl Added Volume of NaOH Mass Antacid Tablet Moles HCl Initially Added Moles HCl Excess Moles HCl Neutralized Moles HCl Neutralized Per g Antacid Tablet Moles of Antacid in Tablet Mass of Antacid in Tablet % Antacid in Tablet % Error (compare to label amount) Trial 1 Trial 2 Trial 3 RESULTS Show all calculations on a separate sheet. 43

48 CONCLUSIONS Discuss how the results for % antacid compare to the amount specified on the label. Using the average of the three trials, determine which brand is more effective and why. Suggest some possible sources of error in this determination and how they might affect results. 44

49 POST-LAB QUESTIONS 1. Calculate the volume (in milliliters) of stomach acid (0.020M HCl) that can be neutralized by the amount of antacid per tablet in each brand using the label amount. 2. For five different brands of antacids, complete the information below by making a trip to your local supermarket or drugstore. # Brand Antacid Amount/ Tablet Price per 100 Tablets Explain which brand is the best purchase. 3. Why do you chew the antacid tablet and not just swallow it whole? 4. Why does boiling remove the carbon dioxide? 45

50 5. If the carbon dioxide from a carbonate antacid was not removed by boiling, how might the final results be affected? Explain. 6. Explain why antacid tablets are less than 100% antacid. 46

51 Name Partner(s) Section Date MORE LIGHTS, COLOR, ABSORPTION! PRE-LAB QUERIES 1. The terms absorption and transmittance are often used when describing the interaction of light with matter. Explain what each term means and how absorption and transmittance of light are related. 2. What possible factors influence the amount of light absorbed by a solution and how might they affect it? OBJECT In this activity you will explore the variables that influence the absorbance of light by a solution. From the data collected you will derive Beer's Law. INTRODUCTION The spectrophotometer measures the ratio of the intensity of the outgoing light (I) to the intensity of the incoming light (I o ). The light impinging on the sample in a spectrophotometer is monochromatic (one specific wavelength, λ). 47

52 The spectrophotometer can display output as absorbance, which is (-log I/I o ), and as percent transmittance (%T), which is (I/I o x 100). Absorbance is NOT the same as absorption, which is (1 T). We will quantitatively explore three variables that influence the absorbance: wavelength; pathlength; and, concentration. Some conditions that influence absorbance will also be explored. Sample Preparation PROCEDURE 1. Place about 5 ml of copper (II) sulfate (0.015M) into a clean labeled spectrophotometer tube. Add 5 drops of concentrated ammonia (NH 3 ) to the tube and mix carefully. The following reaction occurs on mixing: Cu ++ (aq) + 4 NH 3 (aq) Cu(NH 3 ) 4 ++ (aq) If any white precipitate, Cu(OH) 2, remains add 1 or 2 drops more of ammonia to dissolve it. 2. Prepare a dilution of the 0.015M CuSO 4 by placing 5.0 ml of 0.015M CuSO 4 into a 25 ml graduated cylinder and adding water to bring the total volume to 15.0 ml. Mix the solution. Calculate the concentration of CuSO 4. Place 5 ml of this dilution into a clean labeled spectrophotometer tube and add 5 drops of concentrated ammonia. Mix well. Repeat using 10.0 ml of CuSO 4 and diluting to a final volume of 15.0 ml. Calculate the concentration of CuSO 4. Place 5 ml of this dilution into clean labeled spectrophotometer tubes and add 5 drops of concentrated ammonia. Mix well. Save the unused 10 ml of this solution for Part II and add 10 drops of concentrated ammonia to it. 3. Prepare a blank for calibration by placing 5 ml of distilled water into a clean labeled spectrophotometer tube and adding 5 drops of concentrated ammonia. The blank contains all the same reagents except the analyte. Any absorbance due to water (solvent), reagents, and/or the glass spectrophotometer tubes will be removed when the blank is set to zero absorbance (or 100%T) during calibration. Part I- Effect of Wavelength 4. Using the Spec-20D and the instructions given by the computer for the A vs λ function, scan the 0.015M Cu-ammonia complex solution from 400 to 650 nm at 20 nm intervals. Locate the wavelength of maximum absorbance, λ max, from the plot on the computer screen. Generate/print a graph of absorbance vs. wavelength from the computer. 48

53 Part II- Effect of Pathlength 5. One of the Spec-20D instruments will be set up to change the size of the sample tubes used. Set the wavelength control to λ max as determined in Part I. Measure the absorbance for the 0.010M Cu-ammonia complex solution in the three different sized tubes provided. Your instructor will show you how to set up the instrument and how to change the sample tube holders. Record the data in the section provided. Generate/print an absorbance versus pathlength graph. Determine the regression equation (trendline) and r 2 value and be sure these are printed on the graph. Part III- Effect of Concentration 6. Measure the absorbance of 0.005M, 0.010M, and 0.015M Cu-ammonia complex solutions at λ max using the A vs. C function. Record the data and generate/print an absorbance versus concentration graph (calibration curve). Determine the regression equation (trendline) and r value and be sure these are printed on the graph. Change the wavelength from λ max to 550 nm and repeat the measurements for the three Cu solutions. Record the wavelength and absorbances. Generate/print an absorbance versus concentration graph. Determine the regression equation (trendline) and r 2 value and be sure these are printed on the graph. Part IV- Potential Errors 7. Place the 0.015M Cu solution in the sample holder with the mark on the spectrophotometer tube aligned with the mark on the sample holder. Close the lid and read the absorbance. Consider each of the actions below. In the Data section, predict and record what might happen to the absorbance if this action was taken before making an actual measurement. Once you have recorded the prediction, read and record the absorbances at λ max for each situation: Leave the lid open; Rotate the tube 45 o (quarter turn); Remove the tube and place fingerprints on it; Clean the tube of fingerprints and add a small amount of Fuller's Earth and mix. 49

54 DATA Part I- Effect of Wavele ngth (using 0.015M Cu solution) Wavelength Absorbance Wavelength Absorbance 400 nm From computer plot λ max =. Part II- Effect of Pathlength (using 0.010M Cu solution) Pathlength Absorbance 0 cm cm 1.8 cm 2.4 cm Part III- Effect of Concentration Concentration 0 (blank) 0.005M 0.010M 0.015M Absorbance at λ max 50 Absorbance at 550 nm

55 Part IV- Potential Errors Potential Error Prediction Absorbance None (control) open lid rotated tube fingerprints Fuller's Earth RESULTS Generate the following graphs and attach them along with any calculations. For each graph generate the best fit curve or plotted regression equation (a smooth curve or straight line) through the data using the appropriate function in SpectroPro or Excel. Be sure that the regression equation and r 2 value are displayed on the graph in the form of the variables. absorption spectrum (absorbance vs. wavelength) locate λ max on the graph absorbance vs. pathlength determine the slope calibration curve (absorbance vs. concentration) (plot both wavelengths using different symbols) determine the slope for each wavelength CONCLUSIONS 1. Describe the type of relationship found for each of the three variables explored and absorbance. If the wavelength is held constant, write a simple mathematical statement of the incorporating absorbance, pathlength, and concentration. 51

56 2. Explain the cause of each of the errors examined in part IV. Error Possible Cause open lid rotated tube fingerprints Fuller's Earth POST-LAB QUESTIONS 1. On the absorption spectrum of the 0.015M Cu-ammonia complex solution, sketch as a dashed line the curve for a 0.005M Cu solution. 2. If the ammonia was not added to the copper (II) sulfate solutions, how would the results have been influenced? Explain. 52

57 3. What happens to the concentration when the cell size (pathlength) is doubled? Explain. 4. Beer's law, stated mathematically, is given below: A = abc where the terms are defined as A = absorbance; a = absorptivity; b = pathlength; and c = concentration (if c = molar concentration then a = molar absorptivity which is sometimes abbreviated with ε). The equation for a straight line is y = mx + b, where m is the slope and b the y-intercept. If the line passes through the origin, then b is zero and the equation simplifies. What type of plot did you get for A vs. b? What is the slope equal to? What type of plot did you get for A vs. c at λ max? What was the slope equal to? How did the slope for A vs. c at λ max compare to the slope for the same graph at 550 nm? Explain why. 5. Calculate the molar absorptivity at λ max for Cu(NH 3 ) 4 ++ using the slope of the calibration curve and Beer s Law. 53

58 6. What variable in Beer's Law is changing to cause the difference in slope for the two calibration curves determined in part III? 54

59 Name Partner(s) Section Date DETERMINATION OF THE RATE LAW FOR CRYSTAL VIOLET/SODIUM HYDROXIDE REACTION PRE-LAB QUERIES 1. Crystal violet has an intense violet or purple color when dissolved in aqueous solution. If the crystal violet was slowly converted to a colorless compound, how would the color intensity of the mixture change? 2. Based on laboratory techniques you have used in the past, how could we measure the color change above and relate it to the concentration of crystal violet? OBJECT This activity involves the measurement of the initial rate of reaction for crystal violet and sodium hydroxide by spectrophotometric determination of crystal violet concentration over time. From the data collected the rate law or rate expression and rate constant will be determined. INTRODUCTION Crystal violet (or methyl violet) is used as an acid-base indicator, biological stain (gram positive test in microbiology), textile dye, and as a topical antibacterial agent. The intense violet color is due to the resonance of electrons in the alternating single and double bonds. Hydroxide ion in high concentration will attack the crystal violet cation at the central carbon atom where the three rings are bonded. When the OH - bonds to the carbon, a colorless product is formed. CV + + OH - CV-OH purple colorless Since the crystal violet cation has an intense color, we can monitor its concentration by measuring the absorbance using spectrophotometry. 55

60 The crystal violet-sodium hydroxide reaction has the following general rate law: Rate = k(cv + ) x (OH - ) y By conducting three experiments where different concentrations of CV + and OH - are mixed, we can measure the initial rate of reaction and derive the rate law and rate constant, k. The absorbance of CV + as a function of time will be measured using the Spec-20D and computer. Beer's law allows us to relate the concentration to absorbance, hence a plot of A vs. time is essentially concentration vs. time, or the slope is the initial rate of reaction: Rate o = slope = - A/ time = - conc./ time PROCEDURE 1. Determine the λ max for CV + by scanning the 6 x 10-6 M solution of crystal violet from nm at 20 nm intervals using the Spec-20D and SpectroPro on the computer. Use distilled water as the blank. Print the absorption spectrum when completed. 2. Set λ max and standardize the Spec-20D. Have the computer set up to measure absorbance vs. time for 5 minutes with absorbance readings every 10 seconds. Three experiments will be run with reactant concentrations given in the table below: Expt. (CV + ) (OH - ) x 10-5 M M x 10-5 M M x 10-5 M M 3. Pipet 2 ml of each solution into a spectrophotometer tube. Mix well. Wait one minute and then place the tube in Spec-20D and start the computer run. Record the laboratory temperature. The one minute delay is to allow the reaction to get started. When the run is completed, check the output. If there are two or more readings that are identical at the beginning of the run, repeat the run (since this indicates that the reaction had not started when the run began) 4. At completion of each run or experiment, the computer will fit a linear regression line to the data plotted. The slope of the line is the initial rate of reaction. Record the slope (m) and the coefficient of determination (r 2 ) in the data table. 56

61 5. Set up Experiment 3 again as above; however, change the time of the run from 5 to 20 minutes. Fit the resulting data with a linear regression and record the r 2 value in the data table. 6. Transfer the data from the 20 minute experiment to an Excel file. Save the file in the CHM 103 folder in CV_Order using the name group_id_section #. group_id is some name that identifies your group. Generate plots of A vs time, log A vs time, and 1/A vs time. Fit each plot with a linear regression and record the r 2 values in the table provided. The plot with the highest r 2 value (best fit) indicates the reaction order. 7. When you are finished, place all spectrophotometer tubes in the solution of 1M HCl provided to remove any crystal violet stains. 57

62 DATA Absorption Spectrum Wavelength Absorbance Wavelength Absorbance 400 nm Laboratory temperature Rate Experiments Expt. Initial Rate Correlation Coefficient

63 RESULTS 1. Generate the absorption spectrum for the crystal violet cation. Locate ë max on the graph. 2. Calculate the concentration of each reactant after mixing and record the rate constant from each run. Compute the average k. Expt. (CV + ) o (OH - ) o Rate Constant, k AVERAGE 3. Derive the rate law (or rate expression) and evaluate the rate constant, k, using the initial concentrations for each experiment. Calculate the average rate constant. Attach a separate sheet with calculations. 4. Results from the 20 minute run data. Order Zero First Second Graph A vs t log A vs t 1/A vs t r 2 Based on the results above, the order of the reaction is: 5. Overall order from the 20 minute run n overall = What does the 20 minute run tell you about the overall order of the reaction? Does this agree with your result for the 5 minute runs? 59

64 CONCLUSION State the rate law found and your average rate constant at the temperature measured. Discuss any possible errors in your experimental process. POST-LAB QUESTIONS 1. After initially mixing the reactants, it takes time to completely mix the reactants and place the tube in the spectrophotometer (including the one minute delay). Would a two or three minute delay cause an error? How would it influence the graph of A vs. time? 2. If the temperature increased while the reaction was occurring in the spectrophotometer, how would the graphical results (A vs. time) be influenced? Explain. 3. What is the ph range and color change for crystal violet (methyl violet) as an acid-base indicator? Would this have been a consideration for this laboratory activity? Why or why not? 4. Why do we use initial rates to study a reaction? 60

65 Name Partner(s) Section Date INVESTIGATING A HYPOTHETICAL INITIAL RATE MODEL OF THE BEHAVIOR OF CRYSTAL VIOLET ION (CV + ) ABSORBANCE You have had the opportunity to investigate the kinetics of the crystal violet-sodium hydroxide reaction in the laboratory. This activity is an extension of that experience using a computer model that will allow you to change reaction conditions and explore the effects of variable changes on the reaction system. 1. Bring up the STELLA model CV_absorbance_kinetics_103" on the computer. This model permits variation in five variables that might influence the rate of reaction and/or the overall absorbance measured during a run. The influence can result from changes in Beer s Law parameters or from an error during the experiment that affects the kinetics. The five independent variables in the model are listed below: CV + concentration Temperature Wavelength NaOH concentration Pathlength 2. Select RUN and generate the graph for the model set at the initial conditions. Sketch the graph on the axes below. The variables can be changed by adjusting the sliders or clicking on the box with the numbers in it and typing the value in. To set the model back to the original, click on the U that appears on the slider when it is changed from the original value. Assume the rate law for the reaction is given by: Rate = k [CV + ] [OH - ] In general, when a variable changes in the above equation, how does the line responddifferent slope or intercept? Explain. For the absorbance versus time graph, what is the formula for the slope defined in terms of the two variables being measured? What does the slope represent in terms of the chemical reaction? Why is the slope negative? 61

66 3. What effect on the rate would you predict by changing the concentration of either the CV + concentration or NaOH concentration? Why? Keep NaOH constant and change [CV + ]. Sketch the resulting graph on the axes above and label it clearly. Include the change made. Keep CV + constant and change [OH - ]. Sketch the resulting graph on the axes above and label it clearly. Include the change made. Is the response the same for CV + and NaOH? Explain why or why not. 4. Predict how a decrease in temperature might influence the graph. Explain why. Further Investigations in the Chemical World 62

67 What does temperature influence in the rate law? Change the temperature and look at the response. Sketch the resulting graph on the axes below and label it clearly. Include the change made. Suppose the cell warmed up from instrument heat about half-way through the run. How would the run be influenced? See the model for a method to simulate a temperature change mid-run. Sketch the resulting graph on the axes below and label it clearly. Include the change made. 5. How might the system respond to a change in pathlength? Why? Further Investigations in the Chemical World 63

68 Make an adjustment in the pathlength. Sketch the resulting graph on the axes below and label it clearly. Include the change made. Explain the results of the graph when the pathlength was changed. 6. How might a run respond to a change in wavelength (where λ max = 500 nm)? Explain your prediction. Sketch the resulting graph on the axes below and label it clearly. Include the change made. Further Investigations in the Chemical World 64

69 Was your prediction correct? Explain why or why not. Changing the wavelength, λ, influences the absorbance; yet, λ does not appear in Beer s Law. Explain the variation. 7. For all the runs in this model, the measurements were taken with the normal one minute delay. Suppose that you waited another 1.5 minutes to start the measurements, would you get the same rate? Explain. Suppose you waited 10 minutes, would you get the same rate? For each situation, sketch the predicted graph on the axes below and label it clearly. Suppose the reaction in a particular run is slow starting. On the axes above, sketch what the graph would look like with and without time delay. Does the slow start without time delay effect the regression line? Explain. Further Investigations in the Chemical World 65

70 8. What happens if a student puts CV + in the spectrophotometer tube but forgets to add the NaOH? Sketch and label the result on the axes below. What happens if a student puts NaOH in the spectrophotometer tube but forgets to add the CV +? Sketch and label the result on the axes below. If the (CV + ) is cut in half and the (NaOH) is doubled, does anything change? Explain. Sketch and label this situation on the axes below. 9. Suppose when the CV + was made up, it gave an absorbance greater that 1.00 (A > 1 has too large an error in measurement). What variable(s), other than diluting the CV +, could be adjusted to allow the run with a starting absorbance at 1.00 or less? Further Investigations in the Chemical World 66

71 Name Partner(s) Section Date SPECTROPHOTOMETRIC DETERMINATION OF THE EQUILIBRIUM CONSTANT FOR THE FORMATION OF A COMPLEX ION PRE-LAB QUERIES 1. What is the oxidation number of iron in FeSCN +2? 2. Consider the equilibrium for this activity given in the introduction below. How could you shift the reaction to change the concentration of the complex ion? OBJECT In this activity you will determine the equilibrium constant for the formation of the thiocyanatoiron (III) ion, FeSCN +2. Since this ion has a reddish-orange color, its concentration can be determined by spectrophotometry. INTRODUCTION An orange colored complex ion is formed by the addition of iron (III) ion and thiocyanate ion as long as the iron (III) ion concentration is well above the thiocyanate concentration. Fe +3 (aq) + SCN - (aq) FeSCN +2 (aq) (1) Since the thiocyanatoiron (III) ion, FeSCN +2, is the only colored ion in the reaction above, its concentration can be determined by spectrophotometry. If excess SCN - is present the following reaction occurs forming a blood-red complex, Fe(SCN) 2 + (dithiocyantoiron (III)), an old movie method of making blood. FeSCN +2 (aq) + SCN - (aq) Fe(SCN) 2 + (aq) (2) In this activity we will maintain the SCN - at low values compared to the iron (III) ion to avoid reaction (2) from occurring. 67

72 The equilibrium constant for reaction (1) is given below: K = [FeSCN +2 ] [Fe +3 ] [SCN - ] If we start out with a very large iron (III) concentration, reaction (1) will be shifted to the products such that the [SCN - ] o = [FeSCN +2 ] and the absorbance of the complex ion can be measured. This will serve as the standard since both concentration and absorbance are known. We will assume that Beer's Law holds and that we can determine the concentration of the complex by measuring absorbances of the complex ion while lowering the concentration of Fe +3 in solution. The complex ion concentration in various equilibrium mixtures can be found from the following equation. c unkn = (A unkn /A std )c std PROCEDURE 1. Working in pairs, each group will need to clean and dry six test tubes and label them 1 to 6. Obtain seven spectrophotometer tubes and two 5 ml pipets. 2. Pipet 5 ml of 2.0 x 10-4 M NaSCN into each of the six test tubes. Pipet 5 ml of 0.20 M Fe(NO 3 ) 3 into test tube #1. This is your standard for the analysis. The concentration of Fe +3 is so large as to shift essentially all the SCN - into the complex. The remaining solutions will be made by diluting the Fe +3 in a graduated cylinder, allowing the reaction to shift to the left and be at equilibrium. Mix all the solutions thoroughly. 3. Place 10 ml of 0.20 M Fe(NO 3 ) 3 into a 25 ml graduated cylinder and add distilled water to bring the volume to 25 ml. You have 25 ml of M Fe(NO 3 ) 3. Pipet 5 ml of this solution into test tube #2. 4. Place 10 ml of M Fe(NO 3 ) 3 into another 25 ml graduated cylinder and add distilled water to bring the volume to 25 ml. You now have 25 ml of M Fe(NO 3 ) 3. Pipet 5 ml into test tube #3. 5. Place 10 ml of M Fe(NO 3 ) 3 into another 25 ml graduated cylinder and add distilled water to bring the volume to 25 ml. You have 25 ml of M Fe(NO 3 ) 3. Pipet 5 ml into test tube #4. 6. Place 10 ml of M Fe(NO 3 ) 3 into another 25 ml graduated cylinder and add distilled water to bring the volume to 25 ml. You have 25 ml of M Fe(NO 3 ) 3. Pipet 5 ml into test tube #5. 7. Place 10 ml of M Fe(NO 3 ) 3 into another 25 ml graduated cylinder and add distilled water to bring the volume to 25 ml. You have 25 ml of M Fe(NO 3 ) 3. Pipet 5 ml 68

73 into test tube #6. 8. Using the standard (test tube #1), perform a wavelength scan from 400 to 650 nm at 20 nm intervals to determine λ max. Record in the table and generate/print an absorbance vs. wavelength graph (absorption spectrum). If you use Excel, select a smooth, connected scatter plot format. 9. Set the spectrophotometer to λ max. Measure the absorbances of the solutions in all six test tubes after setting the spectrophotometer using distilled water as a blank. Record the data. DATA/RESULTS Absorption Spectrum Wavelength Absorbance Wavelength Absorbance 400 nm λ max = Attach the absorption spectrum of Fe(SCN) +2. Label λ max on the graph. 69

74 Tube Number [Fe 3+ ] o Absorbance 1 (standard) For all tubes: [SCN - ] o = Open the following file from the CHM 103 Folder: Fe(SCN)++_equil_const.xls. Record the absorbance and Fe 3+ concentration data in the Excel spreadsheet. The spreadsheet will calculate the equilibrium values for all reactants and products and the equilibrium constant (K) using the following equations: [FeSCN +2 ] eq = [Fe +3 ] reacted = [SCN - ] reacted [SCN - ] eq = [SCN - ] o [FeSCN +2 ] eq [Fe +3 ] eq = [Fe +3 ] o [FeSCN +2 ] eq K = [FeSCN +2 ] [Fe +3 ] [SCN - ] 70

75 The spreadsheet will compute the average K, %CV, and % error. Be sure that you complete the spreadsheet with the group member names and save the file on the desktop with a different file name. Print a copy of the spreadsheet for the instructor and submit this before you leave the lab. The instructor will post an address where you will be able to obtain the class data. Use these to analyze results and make conclusions. ANALYSIS/CONCLUSIONS Be sure to attach a copy of the spreadsheet and any graphs to this activity if it is being submitted for credit. 1. The literature value of the equilibrium constant is 140. How does your average value compare to this value? Rank the groups based on the degree of error. Discuss factors that might contribute to a larger % error and how they affect the value. 2. How does your precision compare to other groups of students? Which group was most precise? Least precise? Support your statements! POST-LAB QUESTIONS 1. Using a mathematical equation, show how the equilibrium concentration of the complex ion, [Fe(SCN) 2+ ] eq, is determined? Define all known and unknown variables. 71

76 2. Complete the table below using ONLY the quantities already given in the table or the number zero if appropriate. Reaction Fe 3+ + SCN - Fe(SCN) ++ initial [Fe 3+ ] o [SCN - ] o equilibrium [Fe(SCN) 2+ ] eq 3. Write the equilibrium constant for the reaction. 4. The calculation of the equilibrium constant will be done by supplying your data for the six test tubes into an active Excel spreadsheet. The Excel spreadsheet is set up with all the necessary calculations stored and protected in the appropriate cells, you just type in your data and the calculations will be done for you. You must still understand how the calculations are done! Attach the printout of the Excel spreadsheet results. 5. For test tube #1, which is the standard due to its very high (Fe 3+ ), how is the reaction influenced by the following additions? Consider both how the reaction shifts and how the absorbance might change. Addition Fe 3+ + SCN - Fe(SCN) 2+ Shift Absorbance case 1 case 2 add more add more 72

77 Name Partner(s) Section Date AN INVESTIGATION OF ELECTROCHEMICAL REACTIONS PRE-LAB QUERIES 1. How does electricity or electrical current get from the power plant to your house? Can you get electrical current from any source other than an electrical outlet? If so, what? 2. Dissociate the following soluble salts into ions: NaCl MgSO 4 Fe(NO 3 ) 3 3. If the following gases were produced during a reaction, how would you identify them? Consider chemical tests and any physical properties to aid in identification. hydrogen oxygen chlorine sulfur dioxide 4. For the spontaneous reactions below, predict the products and list evidence that would indicate reaction occurred. Mg(s) + HCl(aq) Fe(s) + CuSO 4 (aq) 73

78 PROBING THE CURRENT SITUATION In this activity we will explore some electrochemical properties of solutions and reactions. Careful observation is critical in this activity. You may want to review An Investigation of Chemical Reactions from your CHM 101 manual. The conductivity of a solution is a measure of its ability to carry electrical current. Obtain a conductivity probe, battery, and voltmeter and set them up as shown in the illustration. Your instructor will help with the wiring. The conductivity probe functions by measuring the relative ability of a solution to conduct electricity. A voltage is read on the voltmeter that is mathematically related to conductivity. (They are NOT directly proportional.). voltmeter - + battery probe beaker with test solution Do NOT change the spacing of the electrodes after you start making measurements with the probe. 1. To show that voltage on the voltmeter is related to the conductivity of the solution, hold a piece of pencil lead (graphite with a clay binder, a good conductor) across the two metal electrodes (paper clips) on the probe using plastic forceps. What happens? 2. What type of particles, ions or molecules, are required for a solution to be conductive? Explain your answer. 74

79 Predict whether each of the solutions in the table below will be conductive. Dip the probe into distilled water and then into each of the 0.01 M solutions listed below and record each reading on the voltmeter. Do NOT leave the probe sitting in any solutions. Rinse the probe in distilled water after each solution. Classify each of the solutes as non-electrolyte, weak electrolyte, or strong electrolyte. Solution distilled water sugar NaCl CH 3 COOH HCl tap water Will Conduct? (Y or N) Voltmeter Reading Type of Electrolyte 3. Now let's see what factors influence the conductivity of a solution. Dip the probe into the solutions of strong electrolytes listed in the table that follows and record the readings on the voltmeter. Remember to rinse the probe with distilled water after testing each solution. Look for trends or patterns in the readings. 75

80 Possible Factor Solution Reading Concentration Type of Salt M NaCl 0.01 M NaCl 0.1 M NaCl 0.01 M NaCl 0.01 M CaCl 2 Cation Anion 0.01 M HCl 0.01 M NaCl 0.01 M KCl 0.01 M NaCl 0.01 M NaOH 0.01 M NaNO 3 Temperature 0.01 M 25 o C 0.01 M 60 o C Summarize your results for each factor and explain how each factor influences conductivity. Does any one ion seem to have an unusually high conductivity? Explain. 4. Conductivity measurements can be used to quickly determine total dissolved solids (TDS) in stream water, since the proportion of ions are fairly constant. TDS is usually measured by massing the residue remaining after evaporating a quantity of filtered stream water (suspended solids removed). Using the handheld meter, measure the TDS content of the stream water provided and tap water. Multiply the scale reading by 10 to get TDS in mg solids/l water. 76

81 Solution TDS stream water tap water Here are some typical TDS levels for various water sources. Source Rainwater municipal water systems rivers and streams seawater TDS < 10 mg/l < 500 mg/l mg/l 35,000 mg/l How do the stream water and tap water you measured above compare to the sources in the table? 5. Next we will pass an electrical current through a solution and observe if any chemical reaction occurs. What types of evidence would indicate that a chemical reaction is occurring? 77

82 Obtain the small-scale Hoffman apparatus, illustrated below, and a battery. The electrodes in this version are made of pencil lead. Be careful they are fragile! Reservoir Electrodes Place distilled water into the reservoir, making sure the electrodes are submerged, and apply a current by connecting the battery to the electrodes. What happens and why? Empty the apparatus and place an aqueous solution of sulfuric acid into the apparatus. Apply the current. What happens and why? Empty the apparatus, rinse with distilled water, and place an aqueous solution of NaNO 3 in the apparatus. Apply the current. What happen and why? Rinse the small-scale Hoffman apparatus with distilled water when finished. Observe the reaction again using the large Hoffman apparatus containing aqueous sulfuric acid. Do NOT change the solution in this apparatus! Test the products and record your results. What type of chemical reaction is occurring? Write an overall chemical reaction. This is called the electrolysis of water. What is the purpose of the sulfuric acid or NaNO 3 in the apparatus? 78

83 6. In any electrochemical reaction, both oxidation (loss of electrons) and reduction (gain in electrons) occur simultaneously. The two electrodes, or half cells, are given names depending on which process occurs. Oxidation occurs at the anode and reduction occurs at the cathode. Here is an example of a half-cell reaction which occurs at the positive electrode of the conductivity probe. Fe(s) Fe +2 (aq) + 2e - This is the anode (Why?) and the process slowly corrodes the paper clip. You may have noticed a rusty appearance of some of the solution after testing them with the conductivity probe. Set up the electrolysis cell for aqueous copper (II) chloride. Obtain a 50-mL beaker, two 3-cm pieces of pencil lead, and wires with a battery as shown below. Fill the beaker with copper (II) chloride. battery beaker Apply the electrical current using a battery and observe both electrodes. Do NOT allow this cell to operate more than 2 minutes. Note the polarity (sign) of each electrode. Identify the anode and cathode and write the half-cell reactions occurring at each. 79

84 7. In the reactions above, we supplied the electrical current by means of a battery. Can a chemical reaction generate electrical current on its own? Explain. Obtain a voltmeter, galvanized nail, and thick piece of copper wire. Attach the nail to the negative lead on the voltmeter and the copper wire to the positive lead. Push the nail and wire into an apple or orange. Record what happens. 8. Now let's construct another voltaic or galvanic cell. We will consider some single displacement reactions using a metal and another metal salt. Obtain a strip of copper and zinc, a porous cup and a 250-mL beaker. Set up the arrangement shown below: voltmeter Zn Cu porous cup beaker The porous cup (made of fragile porcelain) allows electrical contact between the two solutions without allowing mixing. Place 25 ml of 1.0 M ZnSO 4 into the beaker and 25 ml of 1.0 M CuSO 4 in the porous cup. Be sure the metal electrodes are dipping into the solutions. The following overall reaction or the reverse of it will occur in this cell. Zn(s) + CuSO 4 (aq) ZnSO 4 (aq) + Cu(s) Observe the electrodes for evidence of reaction and determine the anode and cathode half-cell reactions. Record any voltage. Be aware, metals in a finely powdered state will look black! 80

85 In galvanic cells, the source of electrons is the anode, which is negative in polarity. The cathode is the positive electrode. Allow the reaction to run for a period of time. Does the voltage change? Explain. Clean the cell up. 9. Describe and attach an illustration of how to construct a cell using Mg and Zn and their 1.0 M sulfate solutions. Set the cell up and measure the voltage, identify the anode and cathode, and determine the half-cell reactions and polarity. POST-LAB QUESTIONS 1. Based on your laboratory observations, rank the three metals (copper, magnesium, zinc) in order of reactivity and explain. > > MOST LEAST 2. Lead storage batteries (a collection of galvanic cells) in automobiles commonly fail (do not produce enough voltage to start car) in very cold weather. Why might this be the case? 3. How would the "total dissolved solids" measurement by conductivity be influenced if the following materials were dumped into a stream? 81

86 road salt, CaCl 2 urea, CO(NH 2 ) 2 -a fertilizer and non-electrolyte 4. If you made a cell of magnesium metal in 1.0 M MgSO 4 and copper metal in 1.0 M CuSO 4, what would the anode and cathode half-cell reactions be? Which metal electrode is negative? 82

87 Name Partner(s) Section Date CONDUCTOMETRIC MEASUREMENTS PRE-LAB QUERIES 1. Why did you get a difference in the voltage readings for the 0.01M acetic acid and 0.01M hydrochloric acid solutions using the home-made conductivity probe in the previous lab activity An Investigation of Electrochemical Reactions? 2. Complete the following neutralization reactions: HCl + NaOH CH 3 COOH + NH 3 OBJECT In this activity, the conductivity of solutions will be measured for the dilutions of strong and weak electrolytes, and a titration of a weak acid with a weak base will be performed. The measurement of the equivalence point of the weak acid-weak base titration, considered infeasible by potentiometric means, will be accomplished by graphical analysis. This activity will use Vernier Software conductivity probe along with the Texas Instrument (TI) Calculator-Based Laboratory (CBL) and TI-83 or TI-83 Plus calculator. INTRODUCTION The conductivity of a solution depends on the following factors in the specified manner: 1. directly on the surface area of the electrodes; 2. inversely on the distance between the electrodes; 3. directly on concentration of the ions in solution; 4. directly on the mobility of the ions; and 5. directly on the temperature. Graphite electrode Voltage converted to conductivity by CBL 1 cm 83 1 cm 1 cm

88 Conductivity, in units of microsiemens per centimeter (µs/cm), is measured with a probe using graphite electrodes (fairly inert) with an alternating current used to prevent electrolysis. Earlier activities used steel paper clips. The paper clips corrode (oxidize) due to the applied direct current. With the Vernier conductivity probe, the first two factors listed above are held constant for all measurements. The conductivity probe can be calibrated with a conductivity standard for exact measurements or can be used to determine relative differences as applied in this activity. The probe also has automatic temperature compensation that corrects for any temperatures between 5-35 o C and references conductivities to 25 o C Why is it critical not to change the spacing between the electrodes in a conductivity cell? PROCEDURE Set up the conductivity probe with the CBL and TI-83 calculator. The conductivity probe requires a DIN adapter, which is to be plugged into CH1 on the CBL. The calculator will need the Vernier Software program CHEMBIO loaded on it. Your instructor will provide the program. Use the link cable to link the CBL to your TI-83 calculator. General CBL and TI-83 Settings Use the [PRGM] on the TI-83 or [APPS] on the TI-83 Plus to select the program CHEMBIO. Select the Set up probes from the menu and follow the steps. Select the use stored calibration. After the probe is ready for use, select the Collect data option and choose the trigger/prompt. When the CBL has stabilized push the trigger button on CBL and then enter the x-variable (concentration or volume). Experiment Setting of range switch on probe Set up choice in program Part I - hydrochloric acid 0-20,000 µs/cm 5 Part I - acetic acid µs/cm 3 Part II - titration 0-20,000 µs/cm 5 How will conductivity change on dilution? Will the change be the same for HCl and CH 3 COOH? Now let s explore the behavior. 84

89 Part I - Dilution Behavior of Electrolytes 1. Obtain ml of 0.010M HCl and measure its conductivity using the CHEMBIO program. Be sure that the opening for the graphite electrodes is completely covered by the solution. Dilute the 0.010M HCl by placing 25 ml of M HCl into a graduated cylinder and bring the volume up to 50 ml with distilled water. Measure the conductivity. Repeat the dilution process four more times. After the last datum point, choose graph data to see the plot. Then press [ENTER] NO and QUIT. Use [STAT] EDIT to retrieve the lists of data (L 1 - concentration and L 2 - conductivity). Record the data in the table. Rinse the probe with distilled water. 2. Obtain ml of vinegar, 0.83M acetic acid, and measure its conductivity using the CHEMBIO program. Based on the behavior of HCl, predict the acetic acid conductivity when it is diluted in half (25 ml acid + 25 ml water). Perform the dilutions five times measuring the conductivity after each dilution. After the last datum point, choose graph data to see the plot. Then press [ENTER] NO and QUIT. Use [STAT] EDIT to retrieve the lists of data (L 1 - concentration and L 2 - conductivity). Record the data in the table. Rinse the probe with distilled water. Part II - Titration of Acetic Acid in Vinegar with Ammonia 3. Place 15.0 ml vinegar in a 150 ml beaker and add 25 ml distilled water. Place the conductivity probe in the beaker and have the CHEMBIO program ready for measurement. Record the initial conductivity. From a buret, add the volumes of ammonia shown in the data table. Mix before triggering the conductivity reading! When finished, rinse probe with distilled water. Retrieve data as described above (L 1 - volume and L 2 - conductivity). Press 2nd HALT on the CBL to turn it off! Part I - Dilution Behavior of Electrolytes HYDROCHLORIC ACID DATA ACETIC ACID Concentration Conductivity Concentration Conductivity 0.010M 0.83M

90 Prediction for 0.42 M CH 3 COOH conductivity: Part II - Titration of Acetic Acid with Ammonia Volume Conductivity Volume Conductivity RESULTS 1. Generate the following graphs of the data: a. conductivity vs. concentration of HCl b. conductivity vs. concentration of CH 3 COOH c. conductivity vs. volume of NH 3 2. Explain the difference in conductivity behavior between the HCl and CH 3 COOH. 86

91 3. Determine the equivalence point of the CH 3 COOH - NH 3 titration. Calculate the concentration of ammonia assuming the vinegar is 0.83 M CH 3 COOH. CONCLUSIONS Explain the behavior occurring during each section of the piece-wise function of the curve for the conductometric titration of CH 3 COOH with NH 3. POST-LAB QUESTIONS 1. On your conductivity vs. concentration graph for acetic acid, sketch a dashed line/curve for the behavior of sugar, a non-electrolyte. 87

92 2. Suppose a spectrophotometric titration was performed as outlined by the general reaction below: ANALYTE + TITRANT PRODUCTS intense color colorless colorless Sketch a graph of absorbance vs. volume of titrant. VOLUME OF TITRANT 3. Why is the potentiometric titration of acetic acid with ammonia considered infeasible? 4. The old soda-acid fire extinguishers were pressurized by carbon dioxide to force the aqueous solution out onto the fire by the reaction below: H 2 SO 4 (aq) + 2NaHCO 3 (aq) Na 2 SO 4 (aq) + 2H 2 O + 2CO 2 (g) Explain the warning found on the extinguisher: DO NOT USE ON ELECTRICAL FIRES! 88

93 Name Partner(s) Section Date CHARACTERIZATION OF A MONOPROTIC WEAK ACID BY POTENTIOMETRIC TITRATION PRE-LAB QUERIES 1. Complete the neutralization reactions given below: HNO 3 + NaOH H 2 SO 4 + NaOH CH 3 COOH + NaOH 2. Circle the monoprotic acids in the reactions above. 3. Explain the difference in behavior in aqueous solutions of hydrochloric acid and acetic acid. OBJECT In this activity you will perform an acid-base titration and follow the neutralization reaction by making an electrochemical measurement using a ph meter. From the graphical analysis of the data collected you will determine some characteristics of a weak acid, such as its dissociation constant and molar mass, and how to choose an indicator for visual detection of the end point. INTRODUCTION An electrochemical measurement of the amount of hydrogen ion (H + ) in solution will be accomplished using a combination glass ph electrode and a ph meter or potentiometer. The combination glass ph electrode serves as both the indicator and reference electrodes in one unit. The indicator electrode is the fragile glass membrane (50 µm thick) at the end, which responds to changes in H + concentration by a change in the potential across the membrane. The glass membrane is referenced against a silver-silver chloride electrode of constant potential, connected through a small porous plug on the side of the electrode. To prevent any electrochemical reaction during the measurement (thus a change in concentration of species due to measurement) a potentiometer is used in place of a direct reading voltmeter. The potentiometer measures the cell voltage without drawing current. 89

94 In very simple terms, the cell potential (E) is related to the ph of the solution by the following equation: E = E' - (2.303RT/F)pH where R is the gas constant, T is absolute temperature, F is the Faraday constant, and E' depends on the membrane and internal filling solution. E' is normally constant but will change over long periods of time. The combination glass electrode must be calibrated against a solution of known ph, usually a buffer. The ph meter is set to room temperature to eliminate this as a variable. As a result, the measured E can be read directly in ph units. We will use a computer with a ph probe to follow the titration of a weak acid. This will allow us to measure how the [H + ] changes as the neutralization reaction proceeds. From a plot of ph vs. volume of titrant (NaOH) called a titration curve, we can determine the equivalence point or point of complete neutralization. To accurately locate the equivalence point, we will determine the first derivative of the titration curve. The titration curve will provide data to allow us to determine the dissociation constant, K a, of the acid and its molar mass. PROCEDURE 1. Obtain an unknown solid monoprotic acid sample and record the unknown number. Mass a sample of the solid acid in the range of gram to the nearest tenth of a milligram. Place the sample into a 250-mL beaker, add 50 ml of distilled water and stir with a stirring rod. Since most of the solid acid unknowns do not readily dissolve, warm the beaker on a hot plate and, only if needed, add 5 ml increments of ethanol until the sample is completely dissolved. 2. Clean a buret and fill with standard NaOH solution. Record the concentration of the NaOH. 3. On the computer, start the LoggerPro program called titr.103 (found in the CHM 103 folder). Follow the instruction for a two point calibration using two buffers with set ph values. Never touch or wipe the glass membrane!!! Be sure the electrode remains in some solution at all times unless you are transferring or cleaning it. 4. Rinse the electrode with distilled water and place into the dissolved solid acid solution. Record the initial ph of the solution. 5. Place the buret over the beaker to allow smooth delivery of the NaOH. Use a stirring rod to mix the reaction. 6. Start adding titrant in 2-3 ml increments until the ph reaches 4.5, then start adding in smaller increments as the ph changes. Near the equivalence point you will need to decrease the increment to 0.1 ml. Remember you do not want to blow by the equivalence point or you will have poor results! Data should be collected until the curve seems to have leveled off (very little rise with each increment added). 7. Using the program, display the first and second derivatives. These can be used to 90

95 accurately determine the equivalence point and, subsequently, the K a. Generate graphs which have the original titration curve and each derivative displayed together and print them. Print a complete set of data. 8. Rinse the electrode with distilled water and store as you found it upon completion DATA/RESULTS Unknown solid acid number Molarity of NaOH Mass of solid acid Table of titration data --- attach Generate and attach the following graphs along with any calculations: ph vs. volume of titrant (titration curve.) ph/ volume vs. average volume (first derivative of the titration curve) ( ph/ävolume)/ volume vs average volume (second derivative of the titration curve) On the titration curve, the equivalence point is where the slope of the curve changes from increasing values to decreasing values. This is called an inflection point by mathematicians. On some titration curves the equivalence point can be difficult to locate due to the gentle change in slope. The first derivative allows easy location of an inflection point (sign of slope changes from positive to negative), since the first derivative will be at a maximum for the equivalence point on a titration curve. Determine the volume of the equivalence point from the first derivative plot. Locate this volume on the titration curve and label the equivalence point. Record the equivalence point volume and ph in the table that follows. The second derivative reflects how much change there is in the ph as a function of change in the volume. The second derivative will cross the x-axis (y = 0) at equivalence point. Locate this volume on the titration curve and label the equivalence point. Record the equivalence point volume and corresponding ph in the table that follows. 91

96 Source V eq.pt. ph eq.pt. Titration Curve First Derivative Second Derivative The dissociation constant, K a, can be obtained from the titration curve. Since the pk a is the ph where the (acid) = (salt) which occurs at V eq.pt. /2. Locate it on the titration curve. Calculate the K a. Calculate the molar mass of the weak acid using the equivalence point volume and the sample mass. For an indicator to work, its color change range (pk ind ± 1) must bracket the equivalence point ph. Find two indicators that would correctly work for your solid acid unknown. List the reference source used. Put all information in the table below. Indicator Acid color Base color ph range Citation(s) CONCLUSIONS Summarize what you have determined about the solid acid. Discuss possible sources of error in your process. 92

97 POST-LAB QUESTIONS 1. Compare the shape of a strong acid-strong base titration curve such as HCl with NaOH to the shape of your weak acid-naoh curve. What general differences in shape occur and why? 2. List three indicators that would work for a strong acid-strong base titration. Cite your information source. 3. A diprotic acid is titrated with NaOH and the first equivalence point volume occurs at ml. At what volume will the second equivalence point occur? Why? 93

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99 Name Section Partner(s) Date DISCOVERING INTRAMOLECULAR INTERACTIONS To complete this activity go to the following Internet site: Images will be interactive at this site. Use these sheets to record requested information. Methane and iodomethane have four atoms surrounding the central carbon. Based on your experience with VSEPR, what would the bond angles be? With the cursor on the image click and hold while moving the mouse to rotate the molecule. To check if you are correct, compare the H-C-H bond angles in methane (left) and iodomethane (right) shown above. Right click, go to select and then mouse click action, and then choose angle. You can also determine bond distance and torsion angle from the same Chime menu. The table below summarizes the actions you can take on the images. Property Measured bond angle distance torsion angle Action Taken on Chime Image select and click on three consecutive connected atoms. place the cursor on the first atom, click and then select a second atom and click. click on four adjacent connected atoms in succession. The torsion angle is the angle between the first and fourth atom while sighting down the axis between the second and third. Watch the lower left of your browser screen to get the measured property results. The distance x 100 will give you picometers (pm). See the Chime Guide for help. Methane Iodomethane Why are they different? 95

100 A large constituent atom, such as iodine, in a molecule can push other atoms away and hence, change the bond angles. This is a type of INTRAmolecular interaction or an interaction WITHIN the molecule. This activity will explore intramolecular interactions in a variety of molecules. The bond between atoms in a molecule is a type of intramolecular interaction; however, we are not going examine this; we are going to look at the interaction between nonbonded constituents, such as the iodine affecting the location of the hydrogens in iodomethane. The importance of intramolecular interactions in biological materials will be demonstrated. The interactions BETWEEN two different molecules (or molecules and ions) are INTERmolecular forces. Intramolecular interactions are different from the intermolecular forces (IMFs) that you studied previously to help explain differences in properties of liquids (such as boiling points). IMFs are short-range attractive forces that may be dipole-dipole, hydrogen bonds, ion-dipole, or dispersion forces. Using the wooden ball and stick model kit, build a molecule of ethane, C 2 H 6. What motions within the molecule are possible? Feel free to manipulate the model you built. Look at the animation of the ethane molecule, C 2 H 6, and describe the motion. Which of the two possible arrangements of ethane shown below would be the most stable (that is, have the lowest energy)? Think about interactions between the hydrogen atoms in the ethane. View the molecule in the space-filled mode and decide if your answer is reasonable. (Right click, select display, and select spacefill and then Van der Waals radii) The van der Waals radius is the estimated radius to the edge of the electron cloud in a non-bonding situation. eclipsed staggered 96

101 The two different arrangements of ethane above are the result of rotation around the carboncarbon bond. This is called internal rotation. The repulsive interactions of the hydrogens can be minimized by staggering the configuration of the hydrogens in the molecule. Now suppose we substituted one chlorine atom in place of one hydrogen atom on each carbon. Using the wooden ball and stick model kit, build a molecule of 1,2-dichloroethane (C 2 H 4 Cl 2 ) by replacing a hydrogen on each carbon with a chlorine. Consider the four structures of 1,2-dichloroethane given below. What motion in the molecule yields the four structures? You may want to align the molecules and with your eyes to look down the C-C bond. Measure the Cl-C-C-Cl torsion angle in each of the dichloroethanes and record on the lines above. View the molecules in the space-filled display mode, which more closely illustrates the true size of the electron clouds around the atoms. How would you describe the chlorine atoms on the molecules? Change the display back to ball and stick and then generate the electrostatic potential surfaces for all four molecules. Use the red-white-blue color scheme, where red is negative and blue is positive. How does the charge on the chlorine atoms compare to the remaining part of the molecule? 97

102 How would the chlorine atoms interact with each other? Why? Which of the four structures would be the most stable (lowest energy)? Why? Would 1,2-dichloroethane have a permanent dipole moment? Explain why or why not. These four structures differ only because of the internal rotation around the C-C single bond. No bonds were broken and/or changed. These rotationally different structures of a compound are isomers called conformers. Usually the interconversion of one conformer to another is so fast the isomers cannot be isolated. The conformer with the Cl-C-C-Cl torsion angle of 180 o would have the lowest energy. The highest energy conformer would have a torsion angle of 0 o due to the interaction of the chlorines. The molecule would not have a permanent dipole moment since these structures interconvert. The energy of the conformer can be calculated using computational software. The relative energy (y-axis) of each conformer can be graphed as a function of torsion (or dihedral) angle (x-axis) and is presented in the image below. This image is linked to a movie that shows the structure of the conformer at each point on the graph. Click on the image to bring up the Windows media player and view the animation (it's a big file, so it will take a while to load). Produced by Spartan '02 Explain the variation of the relative energy as the torsion angle changes. 98

103 Does the molecule respond in any other way? Using Excel, plot the data in the table below where the Cl-C-C bond angle is given as the torsion angle varies, as well as, two bond lengths as a function of torsion angle. Produce two graphs- one with bond angle and the other with the bond lengths. Explain any trends in the data. You can copy the data and paste it into an Excel spreadsheet. Torsion or Dihedral Angle Cl-C-C-Cl Bond Angle Cl-C-C Bond Length Cl-C (pm) Bond Length C-C (pm)

104 Let's consider the total amount of variation in the bond angle and the bond lengths as the torsion angle changed for the data above. Calculate the percentage change and compare *. % Change = 100 (range)/(average value) where range = high value - low value * A better way to do this would be to calculate the standard deviation and the coefficient of variation as a percent (%CV). quantity Bond Angle Cl-C-C Bond Length Cl-C Bond Length C-C average range % change How does the variation in bond angle compare to the variation in the bond length? Now construct a molecule of dimethyl peroxide, H 3 C-O-O-CH 3, and predict the most and least stable conformers by considering rotation around the O-O bond. Sketch the conformers. One of the conformers is shown below. Are the possible interactions within dimethyl peroxide as it rotates around the O-O bond the same or different than those in 1,2-dichloroethane? Compare your ball and stick models. 100

105 Using your model of dimethyl peroxide, move the groups through the various conformers from a torsion angle of 0 o to 180 o. At each stage, think of the relative energy based on interactions. Sketch a graph of relative energy as a function of torsion angle. Once you have finished the sketch click here to see an animation of the graph of relative energy as a function of torsion angle. How close was your predicted graph to the actual? How is the graph for the peroxide different from the 1,2-dichloroethane? Explain how and why this might be. What will happen if the methyl groups [-CH 3 ] are changed to the bulkier t-butyl groups [- C(CH 3 ) 3 ]? Build a molecule of di-t-butyl peroxide, (CH 3 ) 3 C-O-O-C(CH 3 ) 3. One of the conformers is shown below. Using your model, move the groups through the various torsion angles. Did you have any problems with the wooden ball and stick model? How is the real molecule going to overcome the problems of the wooden model? 101

106 How is the C-O-O bond angle going to compare for eclipsed and staggered configuration of the two peroxides? Let's examine some data for the two peroxides. Using Excel, plot a graph of bond angle as a function of torsion angle for both peroxides. Print and attach a copy of this graph. Torsion or Dihedral Angle C-O-O-C dimethyl peroxide Bond Angle C-O-O Explain the reason for the difference between the two peroxides. di-t-butyl peroxide Bond Angle C-O-O We have been looking at interactions that are repulsive. Are all intramolecular interactions repulsive? Explain. One of the strongest IMFs is the hydrogen bond. Look at the structure of salicylic acid given below. 102

107 Can you find a position in the molecule where an intramolecular hydrogen bond can form? Indicate where this might be. The DNA molecule is shown below. What is holding the two strands together? Right click, select options, and then display hydrogen bonds. Look for the dashed lines in the molecule. You may want to zoom in by holding the shift key down and moving the mouse. Intramolecular hydrogen bonds play an important role in maintaining the structure of biological materials, such as secondary (coiled) and tertiary (folded) structures of proteins and the double helix of DNA. Other types of intramolecular attractions, such as electrostatic interactions between -NH 3 + and -COO -, also occur in proteins. Post-laboratory Questions 1. How do you think the graph of the C-O-O bond angle as a function of torsion angle for diphenyl peroxide, C 6 H 5 -O-O-C 6 H 5, will compare to the other two peroxide examined earlier? Is the phenyl [-C 6 H 5 ] group (MM = 77 g/mole) as bulky as the t-butyl [- C(CH 3 ) 3 ] group (MM = 57 g/mole)? Ball and stick model is available! The phenyl group is derived from benzene, a planar six member carbon ring with alternating single and double bonds. Sketch your prediction as a dashed line on your earlier graph. Attach your graph. 103

108 2. Hydrogen peroxide, H 2 O 2 or H-O-O-H, has a permanent dipole of 2.0D, which is comparable to water at 1.9D. What does this infer about internal rotation around the O-O bond? Click here to see how the relative energy varies with the torsion angle for H 2 O 2. How is this graph different form pervious relative energy as a function of the torsion angle plots? What is the torsion angle for the lowest energy structure? Sketch and label the graph for H 2 O In general, the overall shape of a molecule can change. How does this occur? Which conformer of hexane, C 6 H 14, has the lower energy? Why? 4. Is the flexibility of a molecule shown by larger changes in bond lengths or bond angles? 5. In a molecule, are bond angles and bond lengths constant? Explain. 104

109 Name Partner(s) Section Date STUDYING VIBRATIONS IN MOLECULES To complete this activity go to the following Internet site: Images will be interactive at this site. Use these sheets to record requested information. Consider the formaldehyde molecule given below. You may want to build a ball-and-stick model of it. What is a vibration in a molecule? You may want to review the Molecules in Motion activity. There are six possible vibrational modes of motion in formaldehyde. Draw the structure, sketch arrows, and explain the modes in the table below. View a shockwave animation of modes and the terminology to describe them. mode: mode: mode: wavenumber: wavenumber: wavenumber: mode: mode: mode: wavenumber: wavenumber: wavenumber: Click here to go to the Purdue University chemistry site to verify your modes of vibration you predicted above. Record the wavenumber (energies) for each vibration in the table above. 105

110 Look at the structure of phosgene given below. Does phosgene have the same or different modes of vibration? Explain. Here are the vibration energies (wavenumbers) as calculated using Spartan '02 (AM1 minimization) for formaldehyde and phosgene. mode of vibration CH 2 O (X = H) COCl 2 (X = Cl) C-X symmetric stretching 3121 cm cm -1 C-X asymmetric stretching 3085 cm cm -1 C-O stretching (C=O) 2053 cm cm -1 X-C-X in-plane scissoring 1444 cm cm -1 X-C-X in-plane rocking 1176 cm cm -1 X-C-X out-of-plane wagging 1165 cm cm -1 When the H's are replaced by Cl's, what happens to the energy of the vibration? Why? How does the carbonyl group (C=O) behave from CH 2 O to COCl 2? Now the wavenumber of the vibrations, ν, is given in reciprocal centimeters, cm -1. Calculate the range of wavelengths, λ, in nanometers (nm) of electromagnetic radiation for the typical energies of cm -1. Remember that 1 m = 100 cm = 10 9 nm. ν = 1/ λ 106

111 Based on your answer above, what part of the electromagnetic spectrum is involved in vibrations? View the video of a single vibration in the amino acid cystine (produced in Spartan '02). The vibration is the stretching of the S-S bond (light blue in center of molecule) at 409 cm -1. How does it affect the molecule as a whole? How does the S-S stretching energy compare to the C=O stretching energy? Let's examine the behavior of a vertical spring with a mass hanging on it. Click here to get a STELLA model to examine this. The vertical spring and mass are similar to the bond as a spring between two atoms, the masses. How do mass and the force constant (spring constant) influence the frequency? Click here to get an interactive Excel spreadsheet for estimating the wave number, which is the energy or frequency of a stretching vibration for a specific bond. You need to use it to address the following three items: 1. How does wave number change as reduced mass increases? Reduced mass is calculated by this equation: Reduced Mass = M 1 M 2 /(M 1 + M 2 ) Support with data and explanation. For this comparison M 1 = 12 for carbon. M 1 M 2 Energy C H C D ( 2 H) C C C F C Cl C Br C I 107

112 2. How does the wave number vary for the H-X bond, where X is a halogen? Support with data and explanation. Here M 1 = 1 for hydrogen. M 1 M 2 Energy H F H Cl H Br H I 3. How does the wave number vary for carbon-carbon single, double, and triple bonds? Support with data and explanation. Set M 1 = 12 for carbon, look only at the row for carbon, and change the setting of k to vary the bond order. Force constant, k Bond order 5 x x x Energy Compare your predicted energy for the C-D bond to the measured value which is 2253 cm -1. Generate the IR spectra for one of the four compounds that will be assigned to your group. Your instructor will help you. Compare and contrast the spectra with other groups (attach analysis as a separate sheet of paper). benzene cyclohexane 108