CHEM 210: Introduction to Analytical Chemistry

Laboratory Manual

Dr. Beth Kautzman, Dr. John Sivey, Dr. Ryan Sours, and Dr. Shannon Stitzel

Spring 2014

1

Table of Contents

Table of Contents ........................................................................................................................................... 1

Laboratory Notebook Guidelines ................................................................................................................... 2

Example Completed Note Card for Experiment 1 .......................................................................................... 3

Guidelines for Writing Reports ....................................................................................................................... 4

Formal Report Evaluation Criteria .................................................................................................................. 6

The Analytical Balance .................................................................................................................................... 7

1 An Introduction to Acid-Base Titrations ................................................................................................... 8

2 Quantitative Volumetric Techniques ...................................................................................................... 13

2A - Precision and Accuracy of Common Analytical Glassware ............................................................ 13

2B Quantitative Serial Dilution ........................................................................................................... 16

3 Determination of Copper in a Penny by Flame Atomic Absorption Spectrophotometry (FAAS) ........... 18

4 Buffer Preparation .................................................................................................................................. 20

5 Introduction to Instrument Calibration and Method Validation ............................................................ 22

6 Determination of Ca2+ by Ion Chromatography (IC) and Flame Atomic Absorption

Spectrophotometry (FAAS) .................................................................................................................... 24

7 Potentiometric Titration of Sodium Carbonate and an Unknown Soda Ash Sample ........................... 273

8 Determination of Caffeine by High Performance Liquid Chromatography (HPLC) and UV-Visible

Spectrophotometry ................................................................................................................................ 30

2

Laboratory Notebook Guidelines

The proper recording of measurements and observations is central to any laboratory science. A proper

laboratory notebook provides a permanent record (often the only permanent record) of the conduct

and results of an experiment. A suitable laboratory notebook should be permanently bound with

consecutively numbered pages (hand numbering is permissible). Most importantly, the notebook

should be used in the laboratory. Notes written on scrap paper, paper towels, etc. for later recording,

does not constitute an acceptable technique!

Rules for Maintaining a Laboratory Notebook

1. Record all data, calculations, and observations in ink (preferably black). No Pencil!

2. Write your name, the course title and number, and the initial and final dates of data entry on the

cover. You may also wish to write your e-mail address on the cover so that you can be contacted if

you misplace your notebook.

3. Starting with the first page, consecutively number each page (front and back) in the upper right-or

left-hand corner.

4. Set aside pages 1-5 for a table of contents.

5. Use headings or labels to identify each series of measurements. Be certain that the heading is

detailed enough to be understood, even by one who has never seen the actual experiment.

6. The date must be given on each page of the notebook.

7. Never erase an entry or use white-out. Rather, use a single line to neatly cross it out (don't

obliterate it!) and briefly explain, in the margin, why it is to be voided.

8. Do not rip pages from the notebook. This will negate the authenticity of the notebook.

Format Guidelines for the Laboratory Notebook

1. Start each new experiment on a new page. Never put information from more than one experiment

on a page. Do not skip pages. The entries for an experiment may not necessarily be contiguous

(e.g.: Exp 2 results, pg 6; Exp 3 results, pg 7; Exp 2 calculations, pg 8 thats the reason for the table

of contents)

2. Entries should be neat, organized, and detailed enough to be clear to you or someone else who

wishes to duplicate your work at a later time.

3. Title each experiment, for example: "Standardization of HCl".

4. Provide a brief (typically one-sentence) purpose statement for each experiment.

3

5. Provide a reference to the procedure wherever appropriate. Be sure to note any changes in the

referenced procedure.

6. List or tabulate, with headings, all relevant data (with units!) and observations.

7. Summarize the results and any pertinent statistical analyses.

8. Discuss briefly any apparent anomalies or discrepancies.

9. Show essential reactions, calculations (no arithmetic), and statistical analyses. Remember to

include all units. Segregate calculations from data and observations.

10. Computer printouts (graphs, tables, etc.) should be taped (not stapled) onto separate pages (dont

cover any other writing). All raw data and calibration curves must be included in your notebook.

4

Example Completed Note Card for Experiment 1

Front

J. Doe 4/1/07

Section 2

Average = 30.11 %

Exp # 1 Standard Dev. = 0.09 %

Rel. Std. Dev. (RSD) = 0.28 %

% composition Conf. Interval = 0.14 % at 95%

30.18 %

30.16 % Unknown Composition =

29.99 % 30.11 +/- 0.14 %

30.11 %

Back

Show a sample calculation for one of the % composition entries (choose one from the front of the card, left column).

5

6

Guidelines for Writing Reports

The formal report is the culmination of all the work you as the experimenter have performed. The

report is generally the only material that the outside world (your professor, supervisor, colleagues etc.)

see of your work and, consequently, judgments about you and your work are generally based on the

report. It is therefore critical that the report, in its presentation, organization, and content, be of the

highest standard of which you are capable. All reports are an individual effort, even for experiments

performed with a lab partner.

A report should be typed, written in complete sentences, and free of spelling, grammatical, and

mathematical errors. The data presented should be clearly stated and be a compilation of information

that is directly relevant to the particular conclusion(s) and/or result(s) that you are presenting to the

reader. The following is a list of the section headings and important information that should be present

in a formal report:

Title Page

Include your name, section, the name of your lab partner (if applicable), experiment number

and title, and date submitted.

Experimental (if required)

A brief summary of the method that was followed and instrumentation/techniques used.

Results and Discussion

This section must begin with text (in which you introduce your results). Do not begin this

section with a table or with a figure.

All data collected while performing the experiment must be presented (as tables/graphs).

The final value for a given measurement should be an average, including uncertainties and units

(2.34 0.05 mg/L). For all values less than 1.0, include a leading zero before the decimal point

(CORRECT: 0.23 mL, INCORRECT: .23 mL).

Perform statistical tests whenever appropriate (Grubbs test, t-test, etc.).

Describe your data and discuss its meaning. How does it relate to what you were expecting?

Include a consideration of errors and any conclusions based on interpretation of the results.

Discuss the implications of the results.

Discuss any problems you had and how they were addressed. How could the experiment be

changed to improve the data?

Provide answers to any questions posed in the experiment handout.

7

References

Cite any outside sources of information used in the report.

Examples of reference formats:

Journal

Authors. Title of Article, Abbreviated Title of Journal. Year of publication, Volume number, page

numbers.

1. Ho, M.; Pemberton, J. E. Alkyl Chain Conformation of Octadecylsilane Stationary Phases by

Raman Spectroscopy, Anal. Chem. 1998, 70, 49154920.

Book

Authors. Book title, edition (if listed); Publisher: City of Publication; Year of Publication.

2. Bard, A. J.; Faulker, L. R. Electrochemical Methods, 2nd ed.; Wiley: New York; 2001.

Edited Book

Authors. Chapter title. In Book Title; Editors, Eds.; Publisher: City of Publication; Year of

Publication; page range.

3. Mirzabekov, A. D. Properties, Manufacturing, and Applications of MicroArrays of Gel-Immobilized

Compounds and Cells on a Chip. In Microfabricated Sensors; Kordal, R., Usmani, A. Law, W. T. ,

Eds.; American Chemical Society: Washington, DC, 2002; pp 23-41.

Website

Title of website. Web address. (accessed date).

4. Analytical Chemistry (ACS Publications). http://pubs.acs.org/journals/ancham/index.html

(accessed Aug 23, 2008).

Note: Carefully check to ensure there are no typographical errors in the web address. An

interested reader should be able to access the webpage(s) you cite simply by copying and pasting

the web address. Even a minor typo will cause major headaches (and possible score reductions).

Example text demonstrating the correct use of in-text citations:

This is an example of how references should be cited within a report. Please use the

American Chemical Society (ACS) guidelines for referencing. Do not use MLA or APA formatting! If

you need to reference a previously published observation or value, insert a superscripted number1

8

next to the text to be cited. If you write an entire sentence which summarizes previously published

information, the citation number should appear at the end of the sentence immediately following

the period.2 Directly quoting material from a reference is extremely uncommon in technical writing;

therefore, you should avoid the use of direct quotes. If material is quoted directly from a reference,

this material must appear in quotation marks, followed immediately by the reference number.

According to the syllabus, It is far better to rephrase ideas in your own words rather than to

include direct quotes.3

References should be listed at the end of your report, and should be numbered in the order

that they appear in your text, not listed alphabetically.4 If you need to cite the same reference later

in the text, use the same number it was originally given.1 Sometimes more than one reference

applies to a single sentence or idea. For ideas requiring two references, separate the superscripts

with a comma.4,5 For three or more references cited in numerical order, use a dash.2-5 For additional

examples of how to properly format in-text citations and reference lists, please review recent

articles published in the journal Analytical Chemistry (http://pubs.acs.org/journal/ancham).

9

Formal Report Checklist

Be sure your report contains the items below and that your report is correctly formatted before

submitting it for a grade. The sections in the checklist are NOT the sections you need to have in your

report. Your report should be formatted according to the headings listed in "Guidelines for Writing

Reports" above.

Description of Experimental Method (Lab 8 only)

Contains concisely organized information that allows the experiment to be replicated; identifies

the sources of data; identifies sequential information in an appropriate chronology; does not

contain unnecessary, wordy descriptions of procedures.

Qualitative Evaluation of Results

All data collected for the experiment is displayed in appropriately labeled and titled tables

and/or graphs; data are presented in text as well as graphical form; final values for

measurements are reported as averages including uncertainties and units; appropriate

uncertainty propagation and statistical tests are utilized.

Quantitative Evaluation of Results

Evaluation of required data, calculations, significant figures, and if present, a check standard

comparison.

Discussion of Results and Conclusions

Summarizes the data and statistics; explains expected results and offers explanations and/ or

suggestions for unexpected results; discusses the implications and/ or draws conclusions that

are consistent with the data and scientific reasoning; tone of explanation is clear and does not

overstate the results.

Scientific Format Criteria

All material placed in correct sections; organized logically within each section; appropriate

referencing; free of grammar and spelling mistakes.

10

The Analytical Balance

An analytical balance is one which weighs to 0.1 mg. A number of companies produce reliable, rugged

and reproducible analytical balances. The balances in the laboratory are accurate to 0.1 mg and have a

maximum capacity of 110 g.

General rules for use of the analytical balances:

1. Use only your assigned balance!

2. Keep the balance clean. Remove dust, etc. from the pans with a camel-hair brush.

3. Learn the capacity of your balance, and never exceed this capacity.

4. Objects to be weighed should be at room temperature.

5. Strategies must be developed to ensure that moisture/fingerprints are not transferred to the

object being weighed during handling. You should, however, avoid wearing gloves. Gloves can

cause electrostatic charges to accumulate on glassware. This charge can cause systematic errors

when using an analytical balance.

6. Do not rub or polish objects before weighing (again, to avoid the accumulation of electrostatic

charges).

7. Chemicals are never placed directly on the balance pan. Use a weighing bottle, beaker, watch

glass, etc.

8. After you have completed weighing, check the following:

a) You have recorded your results correctly.

b) There are no objects left on the pan.

c) The balance pan is completely clean.

9. Report and record anything unusual.

11

1 An Introduction to Acid-Base Titrations

Introduction

This experiment will introduce you to the analytical method of volumetric titration. Volumetric titration

is one of the two important classical or wet analytical methods, the other being gravimetric analysis.

Titrimetric methods are based on the ability to prepare a standard solution, that is, a solution where the

concentration is accurately known. The standard solution is then used to determine, either directly or

indirectly, the composition of a sample.

Standard solutions may be prepared by two methods. One method involves the direct weighing and

dissolution of a high purity standard material to form a solution of known concentration.

Alternatively, a known amount of standard material is titrated by a previously prepared solution. The

concentration of the previously prepared solution is then determined by the volume that reacts with a

given amount of standard material. In this experiment you will prepare a standard acid solution which

will be used to determine the amount of sodium carbonate in an unknown base sample.

Procedure

Preparation of 1.0 liter 0.12 M Hydrochloric Acid (HCl)

Into a clean 1.5-liter plastic bottle place approximately 1.5 liters of deionized water. Add a sufficient

volume of concentrated hydrochloric acid to the deionized water to produce an approximately 0.12 M

solution. Cap, mix thoroughly, and label with your name, date and description of contents.

Standardization of 0.12 M HCl

Obtain from your instructor about 2.5 to 3 g of high-purity sodium carbonate in a clean, dry weighing

bottle and dry in an oven at 140C for 2 hours. Accurately weigh three samples, 0.17-0.23 g, of the pure

sodium carbonate into three Erlenmeyer flasks. As always, record every digit from the balance. YOU DO

NOT HAVE TO WEIGH OUT THE SAME AMOUNT OF SODIUM CARBONATE FOR EVERY TRIAL. Dissolve

each sample with about 50 mL of deionized water and add 4 drops of bromocresol green indicator.

Titrate with the 0.12 M HCl to an intermediate blue-green color. At this point, stop the titration and

heat the solution to boiling on a hot plate to expel CO2 (the solution should return to its original bright

blue color), then complete the titration. If you have not overrun the endpoint prior to boiling, the

solution, as indicated previously, will be blue and a sharp blue to yellow color transition will be observed

at the endpoint. Calculate a molarity for the hydrochloric acid solution from the data obtained in each

titration. Average these values. If the relative standard deviation is > 0.5%, consult with your instructor

for further directions.

12

Analysis of an Unknown Base

To obtain your unknown sample of soda ash, submit a clean, dry weighing bottle to your instructor. Dry

in an oven at 140 C for 2 hours. To determine a suitable sample size for analysis, perform a trial

titration: weigh out a known amount of sample, titrate, record the volume delivered and adjust the

sample size in subsequent titrations so that you use a minimum of 25 mL of titrant. Weigh (by

difference) an appropriate amount of sample into an Erlenmeyer flask, dissolve in approximately 50 mL

of deionized water, add 4 drops of bromocresol green indicator, and titrate to the end point with your

standard HCl.

NOTE: Keep the HCl standard solution. It will be used in Experiment 7.

Calculation of Results

Calculate the % sodium carbonate in your soda ash sample, the standard deviation, the percent relative

standard deviation, and the confidence interval at a suitable confidence level. A relative standard

deviation of > 5.0 (i.e. 5 parts per thousand)should be considered unsatisfactory.

13

2 Quantitative Volumetric Techniques

Introduction

Pipets and volumetric flasks are basic laboratory equipment designed to accurately deliver (the pipet) or

contain (the volumetric flask) a desired volume of pure liquid or liquid solution. For this reason,

knowledge of the precision (measured as reproducibility) and the error (measured in terms of accuracy

or tolerance) associated with volumetric glassware is important. Remember, each measurement and

measuring device contributes to the total uncertainty of the final result of an analytical determination.

Accuracy is defined as a measure of how close a measured value is to the real true value. Accuracy is

often described in terms of its absence, error. Tolerance is the manufacturers stated uncertainty in the

accuracy of a measuring device. Precision is a measure of the reproducibility of a measurement. It is

very important to remember that the skill of the operator, as well as the quality of the measuring

device, determine the quality of the final result.

Your objective in this experiment is to become proficient in the use of volumetric measuring devices.

You should perform a sufficient number of replicate determinations to ensure that you are both

competent and confident in your ability to use these devices to their full potential. Proper use of these

critical tools is vital to the success of your future laboratory experiments.

2A - Precision and Accuracy of Common Analytical Labware

The purpose of this part of the experiment is to learn the correct technique for operation of volumetric

pipets and digital pipetters. The precision and accuracy of these tools is determined through a mass

measurement. Successive delivered aliquots of deionized water are weighed and converted to a volume

using the density of water at the temperature of the water. Performing many replicate determinations

allows the estimation of the accuracy and precision associated with the measurement.

Procedure

You are responsible for repeating the following measurements until satisfactory precision and accuracy

are attained for a 10-mL volumetric pipet and a 5-mL digital pipetter.

The 10-mL Volumetric Pipet (Acceptable accuracy 10.00 0.02 mL)

1. Before starting the experiment, ensure that your 10-mL volumetric pipet is clean (no beading of

water is observed when water is delivered from the pipet). Clean the pipet with an appropriate

cleaning solution if necessary (see the instructor for information). Rinse the pipet well using

deionized water.

14

2. Practice filling the pipet and adjusting the meniscus to the calibration mark until you become

proficient with the technique.

3. Weigh a clean, capped 125 mL plastic bottle to the nearest 0.1 mg.

4. Collect approximately 250 mL of equilibrated deionized water from the carboy in lab into a clean

beaker. Record the water temperature and look up the density of water at that temperature.

5. Transfer 10.00 mL of the equilibrated distilled water to the plastic bottle using the volumetric

pipet (remember this type of pipet is drained by gravity only). Be careful and do not allow the

water to splash out of the container. Reweigh the capped bottle plus water to the nearest 0.1

mg. Using the mass difference and the known density, calculate the volume of water delivered

by the pipet. Repeat at least two more times; use as many replicates as necessary to obtain

results that approximate the expectations of the manufacturer. You must have a set of at least

three data points. If any change in technique is made, a new data set should be started. (Note:

It is not necessary to empty the plastic bottle and cap between each addition. Just continue to

add successive aliquots to the bottle, but be sure that you weigh the bottle and cap after each

addition.)

6. Store the pipet in the appropriate pipet jar near the prep room.

The 5-mL Digital Pipetter (Acceptable accuracy 5.00 0.05 mL)

You will follow essentially the same procedure as outlined above for the volumetric pipet. Of course,

you will not need to clean the device as it has replaceable and disposable tips. Your instructor will

describe and demonstrate the proper way to select the delivery volume, fill, and dispense liquid

samples. You will transfer 5.00 mL of the equilibrated distilled water to the plastic bottle using the

digital pipetter for each trial.

Calculation of Results

You will need to look up the density of water based on the temperature of the water used to collect

your data. Convert every mass to a volume using the appropriate density of water. Then, calculate the

mean, standard deviation, and relative standard deviation of the volumes delivered with each device.

For this part of the experiment, an RSD 0.3% indicates acceptable precision. Acceptable accuracy is

achieved if the mean volume is in the range 9.98 10.02 mL for the 10-mL pipet and 4.95 5.05 mL for

the digital pipetter. If your results do not have acceptable precision or accuracy, repeat another set of

measurements for that tool. If your results still do not meet expectations, consult the instructor.

You must show the instructor your notebook with acceptable results prior to moving on to the next

part of the experiment.

15

16

2B Quantitative Serial Dilution

The purpose of this part of the experiment is to learn how to perform accurate dilutions of concentrated

solutions. A solution of copper nitrate will be diluted in multiple steps using volumetric glassware. The

resulting solution will then be analyzed using flame atomic absorption spectrophotometry and the

copper concentration compared to the predicted value.

Procedure

You are responsible for repeating the following dilution and measurement until satisfactory accuracy is

attained.

Solution Preparation

Prepare approximately 1 L of approximately 0.2% (v/v) nitric acid solution by adding a sufficient volume

of concentrated nitric acid to deionized water. This solution will be used as the solvent for all copper

solutions. Store the solution in a plastic bottle in your drawer (you will need it for Experiments 3 and 6).

Obtain from your instructor a copper standard solution. Record the concentration of the solution in

your notebook.

Dilution Calculations

Based on the copper concentration in the standard solution, calculate the total dilution factor needed to

reach a concentration of approximately 2.5 mg/L. The maximum dilution factor for a single dilution step

is 10-fold. You will use a 50-mL volumetric flask to prepare each dilution. Based on the total dilution

factor you calculated and the 10-fold limit for each step, determine how many serial dilutions you must

perform and which pipet(s) you will use.

Dilution of the Copper Standard

Using a volumetric pipet or digital pipetter, accurately deliver an appropriate volume (based on your

calculations above) of the copper standard solution into the clean 50-mL flask. Dilute to the mark with

your 0.2% nitric acid solution. Rinse a 125-mL plastic bottle three times with a small amount of the

diluted solution, transfer the remaining solution into the bottle, and label with the copper concentration

in mg/L. This solution is copper dilution 1.

Using a volumetric pipet or digital pipetter, accurately deliver an appropriate volume (based on your

calculations above) of the copper dilution 1 solution into the clean 50-mL flask. Dilute to the mark with

your 0.2% nitric acid solution. Transfer the second diluted solution to a 125-mL plastic bottle and label

with the copper concentration in mg/L. Continue the serial dilution if needed according to your

calculations.

17

Determination using FAAS

Bring the bottle containing your final dilution of copper solution (~2.5 mg/L) to the instrument for

analysis. Following the standard operating procedures for the instrument, measure and record a

concentration for your final diluted solution and for a reference solution provided by your instructor.

Calculations

Use the measured concentrations of copper in your diluted solution (DS) and the reference solution (RS)

to calculate a corrected concentration for your diluted solution according to the following equation:

DSmeasRSmeas

DScorrectedRSknown

Use the corrected concentration of copper in your diluted solution to calculate the concentration of

copper in the original standard solution. Compare your calculated value to the known copper

concentration in the standard. Acceptable accuracy is indicated by a percent recovery in the range 95

105%. If your results do not have acceptable accuracy, repeat the dilution. If your results still do not

meet expectations, consult the instructor.

You must show the instructor your notebook with acceptable results prior to moving on to

Experiment 3!

18

3 Determination of Copper in a Penny by Flame Atomic Absorption

Spectrophotometry (FAAS)

Introduction

According to the United States Mint, pennies minted after 1837 are composed of a mixture of metals.

From 1864 to 1982, pennies were a 95% copper alloy (with the exception of 1943). Starting in 1982 the

composition was changed to a zinc core with a thin layer of copper (2.500%) plated onto the surface.

Small variations in the minting process will produce pennies with slightly different compositions. Also,

once in circulation, physical wear and chemical oxidation may change the copper content. The purpose

of this experiment is to accurately determine the amount of copper in a penny using flame atomic

absorption spectrophotometry (FAAS). Results from different students and for different pennies will

then be statistically evaluated.

Procedure

You will use the 0.2% nitric acid solution from Experiment 2B. Obtain from your instructor a solution

containing one dissolved penny. Record any information given about the original penny and the solution

in your notebook.

Dilution Calculations

Perform a calculation to determine the necessary dilution factor of the stock penny solution needed to

bring the copper concentration into appropriate range (0.5 10 mg/L) for analysis using the FAAS

instrument. Remember, the maximum dilution factor for a single dilution step is 10-fold. You will use a

50-mL volumetric flask to prepare each dilution. Based on the total dilution factor you calculate, and the

10-fold limit for each step, determine how many serial dilutions you must perform and which pipet(s)

you will use.

Dilution of the Penny Solution

Using appropriate pipets and/or digital pipetters, accurately dilute the original penny solution according

to your calculations. Use 0.2% nitric acid as the solvent for all dilutions. Transfer the diluted solution to a

labeled 125-mL plastic bottle. Repeat the dilution from the beginning until you have three individually

prepared dilutions of the original penny solution.

Determination using FAAS

Bring the bottles containing your three diluted penny solutions to the instrument for analysis. Detailed

instructions for operating the instrument are available in a separate document.

19

Calculations

Use the measured concentration of copper in each diluted penny solution to calculate the concentration

of copper in the original penny solution. Remember to correct your concentration using the reference

solution following the equation from 2B:

DSmeasRSmeas

DScorrectedRSknown

Using the undiluted concentrations and information about the original penny, calculate the percent by

weight of copper in the penny for all three trials. Calculate the mean, standard deviation, percent

relative standard deviation, and 95% confidence interval for the copper composition based on your

measurements. Report your mean value to the instructor. The instructor will post the mean copper

compositions for each student on Blackboard. Using the class data, calculate the mean, standard

deviation, relative standard deviation, and 95% confidence interval for the copper composition based on

all measurements made by all students who analyzed the same penny.

Report

The report for this experiment is abbreviated and should include only a cover page, results and

discussion, and references. See the instructor handout for more information on preparing this report.

20

4 Buffer Preparation

Introduction

A buffer is a mixture of a weak acid and its conjugate base that resists changes in pH when small

amounts of other acids or bases are added. Buffers are most commonly described by the Henderson-

Hasselbalch equation:

p p a log([A ]

[ A])

The buffer capacity is a measure of the resistance of a solution to pH change and is expressed

numerically as the number of moles of strong acid or base required to change the pH of one liter of

buffer by one unit. Buffer capacity is influenced by the concentration and pH of the buffer solution as

well as the identity of the weak acid/base chosen to prepare the buffer.

In this experiment you will prepare a buffered solution of a given pH and concentration starting from a

solid salt.

Procedure

The instructor will assign each student a different buffer pH. The total buffer concentration must be

0.050 M. Available for buffer preparation are ~1 M solutions of HCl and NaOH and the salts listed

below:

Salts Available sodium bicarbonate sodium carbonate

monosodium citrate disodium citrate

monosodium phosphate disodium phosphate trisodium phosphate

The pKa values of the acids associated with the aforementioned salts are listed in the table below:

Acid pK1 pK2 pK3 Carbonic Acid 6.251 10.329 ---

Citric Acid 3.128 4.761 6.396* Phosphoric Acid 2.148 7.198 12.375

21

* Due to significant disparity among reported values in the

chemistry literature, we advise against relying on this pKa value.

Following the instructions provided with the pH meter, calibrate the meter using the available pH

standards. Your instructor will demonstrate proper care of the pH electrodes.

Choose an appropriate salt and prepare at least 100 mL of your assigned buffer (use a volumetric flask!)

in deionized water. Record the measured pH and actual formal concentration in your notebook.

Volumetrically transfer 25.00 mL aliquots of the buffer solution into two clean, dry 100-mL beakers. To

one beaker add 250 L of a 0.50 M HCl (provided), mix well, and determine the pH. To the second

beaker add 250 L of 0.50 M NaOH (provided), mix well, and determine the pH.

Calculation of Results

On a note card, neatly record your name, assigned buffer pH and concentration, actual buffer pH and

concentration, the salt used to prepare your buffer, and the measured pH values after addition of strong

acid and base. The information on your note card will be used to verify that your buffer was prepared

correctly.

Your note card must be approved by the instructor prior to moving on to Experiment 5!

22

5 Introduction to Instrument Calibration and Method Validation

Calibration is a crucial step in any quantitative analysis involving instrumentation. During calibration, the

measured instrument response is correlated with the known analyte concentrations in a series of

standard solutions (external standards). Ideally, a graph (calibration curve) of signal vs. concentration

should be linear with a coefficient of determination (R2) greater than 0.99.

One method for assuring that a calibration is accurate is to analyze an independent sample which

contains a known amount of analyte. If the concentration of that sample is known to the analyst, it is

called a check standard (or calibration check) and the calibration accuracy is determined by the percent

recovery:

CmeasuredCknown

00

If the concentration of the validation sample is known only by a second party, it is called a quality

control sample (or performance test sample).

In this experiment you will prepare accurate external calibration standards from which you will generate

a calibration curve and determine the concentration of a quality control sample. The external

calibration standards will be made in a buffered solution that you prepare.

Procedure

Buffer Preparation

The instructor will assign each student a new buffer pH and concentration (different than the one made

in experiment 4). Quantitatively prepare at least 500 mL of the assigned buffer solution using deionized

water. See Experiment 4 details for available reagents.

Preparation of Quality Control Sample

Transfer approximately 50 mL of your buffer solution into a labeled 125-mL bottle and give the bottle to

your instructor. A quality control sample will be prepared using your buffer as the solvent and returned

to you for analysis.

Preparation of External Standard Solutions

Prepare a series of solutions as detailed below using the stock indicator solution provided by the

instructor. The solvent for all dilutions will be your assigned buffer solution. Use a 50-mL volumetric

flask to prepare a series of solutions with indicator concentrations of 5, 10, 20 and 40 mg/L. Remember,

do not dilute more than 10-fold in any dilution step. Store the standards in labeled 125-mL plastic

bottles.

Determination using UV-Vis Spectrophotometry

Instructions for operating the UV-Vis instrument are available in a separate document. Before using the

instrument, prepare six disposable 1-cm cuvettes for analysis by filling ~2/3 full with your four

23

standards, the quality control sample, and your buffer solution as a blank/zero. Collect five replicate

measurements for each cuvette. For samples at p 4.0, collect all data at 615 nm. For samples at pH

< 4.0, collect all data at 445 nm.

Calculation of Results

Create a calibration curve by plotting the absorbance for each of the four standards versus

concentration. Insert a linear trendline and verify that the coefficient of determination (R2 value) is

greater than 0.99. If the R2 value is less than 0.99, repeat the experiment using new external standard

solutions.

Use the LINEST function in Excel to generate linear least squares parameters and use them to calculate

the concentration of the quality control sample and its associated uncertainty (sx; see equation 4-27 in

the textbook). On a note card, neatly record your name, the R2 value from your calibration curve, the

concentration sx, and the 95% confidence interval for your quality control sample. On the back of your

notecard, show the calculation for your uncertainty and 95% confidence interval. The information on

your note card will be used to verify that your external standards were prepared correctly.

Your note card must be submitted to the instructor prior to moving on to Experiment 6.

24

6 Determination of Ca2+ by Ion Chromatography (IC) and

Flame Atomic Absorption Spectrophotometry (FAAS)

Introduction

Calcium is a major component of seawater (along with chloride, magnesium, sodium, sulfate, and

potassium). Calcium is an essential nutrient for any marine organism that forms a calcium carbonate

(CaCO3) shell or skeleton (e.g. shellfish, coral). Two crystalline forms of CaCO3 (calcite and aragonite)

can precipitate from seawater; the Ksp for calcite is 3.36 x 10-9 and that for aragonite is 6.0 x 10-9.

Ksp = [Ca2+] [CO3

2-]

According to the solubility equation, the ocean is supersaturated with respect to calcium carbonate.

Supersaturated systems are normally driven to thermodynamic equilibrium by precipitation. In the case

of seawater, the presence of magnesium ions inhibits precipitation of calcium carbonate. Many marine

organisms use a combination of calcite and aragonite crystals along with various biomolecules to build

large, complex composite structures (e.g. mollusk shells). An organisms ability to control precipitation

of the two types of calcium carbonate is determined in part by the carbonate ion concentration (most

easily determined as alkalinity see Expt. 7) and calcium ion concentration in seawater.

There are several methods for determining calcium in aqueous samples. The most common are

complexometric titration with EDTA, use of a calcium ion-selective electrode (ISE), ion chromatography,

and flame atomic absorption spectrophotometry. In this experiment, you will compare the latter two

methods.

Procedure

Prepare 1 L of approximately 0.2% (v/v) nitric acid solution (use any remaining solution from Expt. 2B

before making additional nitric acid solution). Assume that full strength nitric acid is 70% (v/v) nitric

acid.

Obtain from your instructor two calcium check standard solutions, one for the IC determination and one

for the FAAS determination. Your instructor will provide the concentrations of the check standards.

Seawater Sample Preparation

Obtain a seawater sample from the aquarium in a clean, dry 50-mL beaker. The diluted seawater

samples for both methods must be prepared on the same day.

IC Seawater Sample: Dilute the seawater 25-fold with 8 Mcm water to produce 50 mL of a sample

solution for analysis. Make sure not to dilute more than 10-fold in any dilution step. Transfer the

solution to a labeled 125-mL plastic bottle for storage.

FAAS Seawater Sample: Dilute the seawater 250-fold with your 0.2% nitric acid solution to produce 50

mL of a sample solution for analysis. Make sure not to dilute more than 10-fold in any dilution step.

Store the diluted sample in a labeled 125-mL plastic bottle.

25

Determination using IC

IC External Standard Preparation: Prepare a series of calcium external standard solutions as detailed

below using the commercial 000 mg/L standard solution. The solvent for all dilutions will be 8 Mcm

water. Use a 50-mL volumetric flask to prepare a series of four solutions with concentrations in the

range 5-40 mg/L Ca2+. As always, do not dilute more than 10-fold in any dilution step. Store the

standards in labeled 125-mL plastic bottles.

IC Determination of Ca2+: Prepare a syringe and a 0.22 m filter. Transfer 3 mL of the diluted seawater

sample into the syringe and filter approximately 1.5 mL into a labeled autosampler vial. Using fresh

syringes and filters, transfer approximately 1.5 mL of each of your external standard solutions into

separate, labeled autosampler vials. Also transfer approximately .5 mL of 8 Mcm water blank, and

the IC check standard into separate, labeled autosampler vials. Give all of your prepared vials (seven

total) to your instructor for loading into the IC autosampler. It is very important that all vials be properly

labeled to indicate both their contents and your group name.

The instructor will load the vials containing the blank and external standard solutions into the

instrument autosampler in order of increasing concentration followed by the check standard and

seawater sample. Three replicate runs will be performed for the seawater sample and check standard.

Determination using FAAS

FAAS Stock Ca2+ Standard Preparation: Obtain a sample of primary standard grade calcium carbonate

(CaCO3) from your instructor in a clean dry weighing bottle. Weigh accurately by difference

approximately 0.1 g of calcium carbonate into a 100-mL volumetric flask. Add approximately 10 mL of

deionized water and then add drop-wise a 50% (v/v) HNO3 solution until effervescence ceases and no

solid remains. The HNO3 reacts with the calcium carbonate according to the following reaction.

CaCO3 (s) + 2H3O+ (aq) + 2NO3

- (aq) Ca2+ (aq) + 3H2O (l) + CO2 (g) + 2NO3- (aq)

Dilute to the mark with your 0.2% nitric acid solution. Transfer to a 125-mL plastic bottle, label and

calculate the concentration of Ca2+ in mg/L.

FAAS External Standard Preparation: Prepare a series of calcium external standard solutions as detailed

below starting from your stock standard solution. The 0.2% nitric acid solution will be used as the

solvent for all dilutions. Use a 50-mL volumetric flask to prepare a series of four solutions with

concentrations in the range 0.5-4 mg/L Ca2+. As always, do not dilute more than 10-fold in any dilution

step. Store the standards in labeled 125-mL plastic bottles.

FAAS Determination of Ca2+: The blank, external standard series, and check standard should be in

labeled 125-mL plastic bottles. The seawater sample must be in a 15-mL centrifuge tube. Detailed

instructions for operating the instrument are available in a separate document.

26

Calculations and Report

The report for this experiment is abbreviated and should include only a cover page, results and

discussion, and references. See the instructor handout for more information on preparing this report.

27

7 Potentiometric Titration of Sodium Carbonate

and an Unknown Soda Ash Sample

Introduction

Many acid-base titrations are difficult to accomplish using a visual indicator for one of several reasons.

Perhaps the analyst is color-blind to a particular indicator color change, there may not be a suitable

indicator available for a particular type of titration, the solutions themselves may be colored, opaque or

turbid or it may be desired to automate a series of replicate determinations. In such situations,

potentiometric titration, using a glass hydronium ion (H3O+) selective electrode, a suitable reference

electrode and a sensitive potentiometer (a pH meter) may be advantageous.

In the first part of this experiment, you will perform a potentiometric titration of the primary standard

grade sodium carbonate from Experiment 1. In the second part of this experiment, you will perform a

potentiometric titration on a soda ash sample containing an unknown amount of sodium carbonate. Be

sure to collect an unknown soda ash sample from your instructor, place it in a weigh bottle and dry it

in the oven (at 140 oC for two hours) prior to beginning Part B of this experiment. Carefully label this

soda ash sample so as to not confuse it with your unknown base from Experiment 1.

Theory

Any acid-base titration may be conducted potentiometrically using a pH electrode. A pH electrode is

composed of an indicator electrode, selective for H3O+, and a stable reference electrode. The indicator

electrode contains a thin glass membrane specially fabricated to preferentially exchange H3O+. When

the electrodes are immersed in a solution, the potential difference between the electrodes is measured

in millivolts (mV). Each glass electrode is slightly different, due to the difficulty of reproducing the glass

membrane. Therefore, it is necessary to calibrate the pH meter and electrode using at least two buffer

solutions of accurately known pH in order to convert the electrode potential to a pH value. If the pH

meter was calibrated properly, the pH meter will automatically convert the potential difference into a

pH value. To increase the accuracy of pH measurements, solutions should be gently mixed (either with a

small magnetic stir bar or via gentle swirling by hand).

The main objective of a titration is to recognize the end point(s), at which the reacting species are

present in stoichiometrically-equivalent amounts. To help identify the end point(s), you should plot

your titration data (as pH versus volume of titrant added) in real-time.

You will employ two mathematical procedures (first-derivative plots and nonlinear least-squares

regression analysis via Microsoft Excels SOLVER function) to analyze titration data. The results of these

mathematical procedures will permit you to calculate pKa1 and pKa2 for carbonic acid and to compare

your findings with the pKa values reported in your textbook. Output from the SOLVER function will also

permit you to calculate the composition of sodium carbonate in the unknown soda ash sample.

28

Procedure

Part A Titration of Primary Standard Grade Sodium Carbonate

The sample for this part of the experiment is the solid sodium carbonate and the titrant is ~0.12 M

standardized HCl, both from Experiment 1. Accurately weigh 0.1 g of sodium carbonate into a 100-mL

beaker and add volumetrically 25.00 mL of deionized water. Based on the amount of sodium carbonate

weighed out and the previously standardized concentration of your HCl (this actual value is not 0.12 M),

calculate both equivalence volumes for the titration. Set up Excel to plot your titration as you go.

Titrate your sample as follows:

1. Standardize your pH meter. After standardization, rinse the electrode using DI water and dab

the exterior (plastic portion) dry with a Kimwipe. Avoid touching the glass portion at the bottom

of the electrode.

2. Place the pH electrode into the solution to be measured together with a magnetic stir bar. The

pH reading will appear on the screen. Allow the value to stabilize before reading. Record the

initial pH of the solution.

3. For the following steps, do not waste time trying to add a specific volume. After each addition of

acid, allow the pH meter to stabilize and then record the pH and exact volume (If the pH has not

stabilized after one minute, consult your instructor).

4. Begin the titration by adding approximately 1.0 mL of titrant at a time until the volume added is

approximately 1 mL less than the first equivalence volume.

5. From approximately 1 mL before through 1 mL after the equivalence point, add approximately

0.3 mL of titrant at a time.

6. Repeat steps 4 and 5 for the second equivalence point.

7. Continue adding approximately 1.0 mL of titrant at a time to reach a final volume of 5 mL past

the second equivalence point. Do not take data points lower than a pH of 2.5.

Part B Titration of Unknown Soda Ash Sample

The sample for this part of the experiment is a new unknown soda ash sample (not your unknown from

Experiment 1) and the titrant is ~0.12 M standardized HCl. You must dry your unknown soda ash

sample in the over at 140 oC for two hours before proceeding. Accurately weigh between 0.45 g and

0.55 g of your soda ash sample into a 100-mL beaker and add volumetrically 25.00 mL of deionized

water. Set up Excel to plot your titration as you go.

Titrate your sample as follows:

29

1. Place the pH electrode into the solution to be measured together with a magnetic stir bar. The

pH reading will appear on the screen. Allow the value to stabilize before reading. Record the

initial pH of the solution.

2. For the following steps, do not waste time trying to add a specific volume. After each addition of

acid, allow the pH meter to stabilize and then record the pH and exact volume (If the pH has not

stabilized after one minute, consult your instructor).

3. Begin the titration by adding approximately 1.0 mL of titrant at a time until the pH is near 9.0.

4. From pH 9.0 to pH 7.0, add approximately 0.3 mL of titrant at a time.

5. From pH 7.0 to pH 5.5, add approximately 1.0 mL of titrant at a time.

6. From pH 5.5 to pH 2.5, add approximately 0.3 mL of titrant at a time.

Do not collect data at pH < 2.5.

Calculations and Post-Lab Assignment

See the instructor handouts for more information on completing the post-lab assignment.

30

8 Determination of Caffeine by High Performance Liquid

Chromatography (HPLC) and UV-Visible Spectrophotometry

Introduction

With the development of agriculture some 10,000 5,000 years BCE in Anatolia and somewhat later in

the Indus Valley, China, and Mesoamerica, plants became an important component of the human diet.

Indeed today three plants, wheat, corn and rice are the major components of the diet of a large section

of the worlds population. Plants cannot only be a source of nutrients, but many plants also produce

compounds, as a part of their metabolism, that can affect humans in other ways. Some plants produce

compounds that are poisonous (hemlock); others can cause depression or act as stimulants while others

can cause profound psychological effects (opium poppies, cannabis, mushrooms). Many of these

compounds and their plant source are banned from legal use. Some though, are tolerated by society.

Caffeine is one such compound that is tolerated. Caffeine occurs naturally in the leaves of the tea plant

and in the beans of the coffee plant. Caffeine is also added to many foods and drinks, particularly sodas

for its mental stimulating effect. People vary widely in their tolerance to caffeine. For some, it has little

effect, for others, it can cause serious problems. The wide variation in peoples tolerance to caffeine

requires that the caffeine content in food and beverages be clearly stated on the label.

Caffeine can be determined in foods by a variety of methods. The two most versatile are high-

performance liquid chromatography (HPLC) and UV-visible spectrophotometry (UV-Vis). In this

experiment, you will analyze for caffeine in beverages of your choice using these two methods of

analysis

Procedure

Samples

Everyone will analyze Coca-cola Classic (provided by the instructor). In addition, select several (at least

three) beverages that you would like to analyze for caffeine, and bring these additional beverages to

lab with you. You may also bring tea leaves or ground coffee beans to be brewed in the lab. Please do

not bring any milk-containing beverages or Mountain Dew to analyze. These drinks are colloidal

suspensions and cannot be filtered.

Determination using HPLC

External Standard Preparation: Obtain from your instructor a sample of pure caffeine in a clean, dry

weighing bottle. Also obtain a caffeine check standard, for which the concentration is known. Prepare

0.500 L of an approximately 350 mg/L solution of caffeine with 18 Mcm water. Caffeine is slow to

dissolve in water, so make sure it has all gone into solution before filling the flask all the way up to the

mark. Calculate the exact concentration of your solution in mg/L (ppm). Prepare five calibration

solutions volumetrically by serial dilution. Transfer the solutions to clean, labeled plastic bottles for

storage.

31

HPLC Standard Preparation: Transfer approximately 3 mL of a standard solution into a syringe, and filter

1.5 mL through a 0.22-m filter into a labeled HPLC vial. Repeat for each external standard and the

check standard. Also fill a vial with 18 Mcm water to use as your blank.

HPLC Sample Preparation: You will need to filter 30 mL of each beverage through a 0.45 m filter.

Transfer the filtered beverages into labeled 50-mL centrifuge tubes, and cap. Many beverages will need

to be diluted for proper measurement. You will prepare HPLC samples starting from the undiluted and

filtered beverage samples prepared in the 50-mL centrifuge tubes. The dilution factor for HPLC samples

will be based on the amount of caffeine in your beverage and may be different from the dilution factor

for UV-Vis analysis. If you do not expect caffeine to be present, you do not need to dilute the beverage.

For moderately-caffeinated drinks (e.g. coffee, tea, coke, etc.) prepare a 25-fold dilution of the filtered

beverage. For highly-caffeinated drinks (e.g. energy drinks) prepare a 50-fold dilution of the filtered

beverage. Transfer approximately 3 mL of a diluted beverage into a syringe, and filter 1.5 mL through a

0.22-m filter into a labeled HPLC vial. Repeat for each beverage sample.

Once you have prepared all 10 HPLC vials, place them in a beaker labeled with your name and lab

section and give the samples to your instructor

HPLC Determination: The instructor will load the vials containing the blank and external standard

solutions into the instrument autosampler in order of increasing concentration followed by the check

standard and beverage samples. Three replicate runs will be performed for the beverage samples and

check standard.

Determination using UV-Vis

UV-Vis Sample Preparation: Many beverages will need to be diluted for proper measurement with the

UV-Visible spectrometer (see determination below). You will prepare UV-Vis samples starting from the

diluted beverage samples prepared for the HPLC. Check each dilution by the UV-Vis determination

below, and then dilute further as necessary.

UV-Vis Standard Preparation: The unfiltered check standard and external standard solutions you used

for the HPLC determination of caffeine can be used for this method.

UV-Vis Determination: To verify that your beverages fall within the range of your calibration curve, you

will first test your caffeine calibration standards and your diluted beverage samples on a portable Ocean

Optics UV-Vis spectrophotometer. Record the UV-Visible spectrum of your most concentrated caffeine

standard solution over the wavelength range of 210 500 nm. Determine which absorbance peak in the

spectrum is most suitable for quantitative determination and record the wavelength of maximum

absorption (max). Record the UV-Visible spectrum of your diluted beverage samples, ensuring that the

maximum absorption of the peak you will use for quantitative determination is less than that of your

most concentrated caffeine standard. If your beverage absorbance falls outside the calibration range,

you will need to determine an appropriate volumetric dilution of the beverage to bring it within the

range and repeat the measurement.

32

Once you have confirmed that your diluted beverage solutions are within an appropriate concentration

range, you may proceed to acquire data on the Cary UV-Vis instrument. Instructions for operating the

Cary UV-Vis are available in a separate document. Use 8 Mcm water as your blank and 1-cm

disposable cuvettes from the prep room. Measure and record the absorption of all five of your

standards, your 8 Mcm water blank, the check standard, and your beverages at the max (determined

earlier). Collect 5 replicate measurements for each cuvette.

Calculations and Report

See the instructor handout for more information on preparing this report.