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University of Oregon Thailand Distance Learning Program Green Chemistry 2010

Copyright Kenneth M. Doxsee, University of Oregon Page 25 of 84

Connecting Solubility, Equilibrium, and Periodicity

Chemical Concepts

Equilibrium; solubility; solubility product; periodic properties; acid/base titration.

Green Concepts

Preventing waste; safer chemicals. (Consider Green Principles 1, 3, 11, and 12.)

Introduction

Chemical equilibrium is one of the core concepts in chemistry, and a firm understanding of

the concept is essential if one is to understand reaction chemistry, acid/base chemistry,

biochemistry, or any of myriad other facets of chemistry. Beginning students, however,

frequently have difficulty with the concept of equilibrium and regularly misunderstand both

its meaning and its implications. Crystallizationthe separation of a solid compound from a

solution of the compoundis a relatively simple phenomenon that may be used to highlight

some of the essential features of equilibrium in an intuitive and understandable way. Most

students have observed, for example, that sugar dissolves slowly in cold water (or coffee or

tea), but more rapidly and in larger amounts in hot water. Many have seen that cooling of a

hot solution of sugar can lead to the formation of crystalline sugar, illustrating that there is an

equilibrium between sugar molecules in solution and those in the solid state.

In this experiment, we will examine the solubility in water of three group II hydroxides,

Mg(OH)2, Ca(OH)2, and Sr(OH)2, using titrations to determine the amount present in a

saturated solution. From the resulting titration data, we will determine the solubility productsfor each of the three hydroxides, gaining familiarity with the concept of Kspand thereby with

consideration of simple solubility equilibria. Based on the trends in solubility products

determined for the three hydroxides, we will then form a hypothesis regarding the solubility

of the two group II hydroxides not examined experimentally, Be(OH)2and Ba(OH)2.

While the determination of solubility products is a common laboratory procedure, reported

experiments often rely on the precipitation of sparingly soluble compounds such as PbI2or

BaSO4, thereby risking student exposure to hazardous heavy metals and generating wastes

that are hazardous and costly to dispose of. In contrast, this experiment uses group II

hydroxides that, while requiring the usual amount of care appropriate for work with basic

substances, are neutralized and present no disposal hazards. The more toxic group II elements,beryllium and barium, are avoided, but students are able nonetheless to explore their

chemistry indirectly, by recognizing the trend in solubility displayed by the three compounds

and extrapolating to predict the solubilities of the remaining two.

Laboratory

Note: For this workshop, each group will titrate only a single group 2 hydroxide, then share

their results with the other groups.

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Materials needed

Saturated solution of Mg(OH)2 50 mL burette with clamp and stand

Saturated solution of Ca(OH)2 250 mL Erlenmeyer flask

Saturated solution of Sr(OH)2 Magnetic stir bar

0.002 M HCl Magnetic stir plate

Phenolphthalein solution (0.1% in ethanol)

Titrate each Group II metal hydroxide saturated solution with 0.002 M HCl using

phenolphthalein as the indicator, using the following procedure.

1. Fill buret with 0.002 M HCl solution. Record the starting volume of the titrant.

2. Obtain a sample of one of the analyteseither 50.0 mL of Mg(OH)2, 1.00 mL of

Ca(OH)2, or 0.500 mL of Sr(OH)2.

3. If you are using Ca(OH)2or Sr(OH)2, add approximately 50 mL of distilled water.

4.

Add 3 drops of phenolphthalein solution.

5. While stirring magnetically, add titrant until the indicator undergoes a color change.

Record the final volume of the titrant.

Questions

1. Calculate the Kspvalue for your sample, then share your data with other groups so that

you all have experimental results for the three hydroxides.

2. What periodic trend in Kspvalues do your experimental findings show? Did you find

the same periodic trend in Kspvalues as the trend given in the lab report?

3. Based on your results, which of the following values of Kspdo you think is most likely

for beryllium hydroxide, Be(OH)2? Explain your choice.

3.0 10-8

7.0 10-22

2.0 10-3

4. Based on your results, which of the following values of Kspdo you think is most likely

for barium hydroxide, Ba(OH)2? Explain your choice.

3.0 10-8

7.0 10-22

2.0 10-3

5.

a) How do beryllium hydroxide and barium hydroxide each compare to the Group

II metal hydroxides you did titrate in the lab in terms of ionic size and effective

nuclear charge (Zeff) on the ions?b) Based on the comparison of size and Zeff, which of the three values of Ksplisted

in question 3 would Coulombs Law predict for beryllium hydroxides Ksp?

c) Based on the comparison of size and Zeff, which of the three values of Ksplisted

in question 4 would Coulombs Law predict for barium hydroxides Ksp?

6. Of the five Group II metal hydroxides, Be(OH)2, Mg(OH)2, Ca(OH)2, Sr(OH)2, and

Ba(OH)2, beryllium hydroxide and barium hydroxide are definitely the most toxic.

The goal of this experiment was to determine the periodic trend in the solubility of

Group II hydroxide compounds. Explain how this experiment is more aligned with the

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concepts of green chemistry than an experiment calling for titrations of all five Group

II metal hydroxides.

Research Questions

1.

Why is beryllium so toxic, when all the surrounding elements (e.g., Li, Na, Mg) are sobenign?

2.

What are some uses of beryllium or its compounds?

References

This is an adaptation of Connecting Solubility, Equilibrium, and Periodicity in a Green,

Inquiry Experiment for the General Chemistry Laboratory,Kristen L. Cacciatore, Jose

Amado, Jason J. Evans, and Hannah Sevian*,J. Chem. Ed.2008, 85(2), 251-253. The

following pages provide a reproduction of this original article. A nice packet of supporting

information is available on-line at theJournal of Chemical Educationweb sitesee the

reference to this information at the end of the following article.

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Determination of Acetylsalicylic Acid in an Aspirin Tablet

Chemical Concepts

Acid-base chemistry; titration.

Green Concepts

Safer reagents; prevent waste. (Consider Green Principles 1, 2, 3, 4, 8, and 11.)

Introduction

This is a simple and inexpensive microscale experiment with aspirin (acetylsalicylic acid,

ASA) that can be successfully used in an introductory experimental chemistry course. It

introduces the concepts of acid/base chemistry and titration. Using a commercial aspirin

tablet, the experiment makes connections between the laboratory and real life, and can be

used to raise issues related to drug purity and analysis, as well as more complex related

matters, such as the manufacture of drug tablets, which contain various excipients, fillers, etc.

in order to provide them with desired shelf life, appearance, and solubility properties.

Acid/base titration is routinely taught in the high school and general chemistry laboratory.

Conceptually, such a titration should not generate hazardous waste, regardless of the scale at

which it is performed, since the titration leads to neutral, easily disposable solutions. In

practice, however, one would be hesitant to introduce the final solutions to the environment,

for fear that titrations were carried out incorrectly and had not properly neutralized the

solutions. Thus, a student titration experiment can lead to a large amount of waste that must

be treated as hazardous. In this experiment, we introduce a simple and effective means ofdramatically reducing the volumes required for titration experiments while maintaining an

appropriate level of accuracy and precision.

Laboratory

Materials needed

Porcelain mortar and pestle Standard 0.05 mol L-1

NaOH

1 commercial aspirin tablet Phenolphthalein indicator

50-mL beaker Two 1-mL syringes (insulin type)

Precision balance One three-way plastic stopcock

Ethanol, 20 mL One automatic delivery pipette tipDistilled or deionized water Silicone hot glue

Sample preparation and dissolution

1. Weigh a commercial aspirin tablet.

2. Grind the tablet to a powder in a porcelain mortar.

3. Transfer the powder to a 50-mL beaker, weighing before and after to determine the

amount transferred.

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4. Add 20 mL of ethanol and 20 mL of distilled water, gently swirling to effect

dissolution.

Sample filtration and titration

(When an aspirin tablet is dissolved in an ethanol/water mixture, the filler/binder from thecommercial aspirin tablet remains in suspension. It does not interfere with the final result, but

the endpoint of the titration is better detected if the solution is filtered.)

1.

Allow the mixture to reach room temperature, then filter it quantitatively into a 100-

mL volumetric flask containing 20 mL of distilled water.

2. Add additional distilled water to the 100-mL mark.

3.

After thorough mixing, aliquots of 1 mL are titrated with standard 0.05 mol L-1NaOH

contained in a 1 mL syringe (the exact concentration, previously determined by a

technician, is given to the students). The Figure illustrates the microtitration apparatus.

Insert the tip of the syringe into a plastic 3-way stopcock valve or an automatic pipette

tip.No needles are needed, and they should be avoided to prevent accidental pricks.Phenolphthalein is used as the indicator.

Automatic delivery pipet tip

1 mL syringe

(for titration)

1 mL syringe

(for refilling)

Three-way stopcock

TH E MEXICAN MICROSCALE

TITRATION SYSTEM

Data reporting and results

During the experiment, a data sheet must be completed by the students (Table 1). A sheet with

typical data obtained by a group of students using aliquots of 20 mL (and titrating with

normal size glassware) appears in Table 2 for comparison.

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Table 1. Data Sheet withActualExperimental Data and Results (using aliquots of 1 mL)

DETERMINATION OF ASA IN COMMERCIAL ASPIRINAspirin tablet

mass

gASA molar mass 180 g mol

.1

Analyzed massg

NaOH concentration mol L

.1

Titration with NaOH Results

sample volumeASA amount per sample

(in moles

and grams)1 mL

2 mLASA amount in analyzed

massg

3 mL %ASA in tablet %

average volume

usedmL

% error, assuming 99.3%

purity%

Table 2. Data Sheet with TypicalExperimental Data and Results (using aliquots of 20 mL)

DETERMINATION OF ASA IN COMMERCIAL ASPIRINAspirin tablet

mass

0.5819 gASA molar mass 180 g mol

.1

Analyzed mass 0.5716 g NaOH concentration 0.0462 mol L.1

Titration with NaOH Results

sample volumeASA amount per sample 5.17 10

-4

molor 0.0931 g1 11.20 mL

2 11.15 mLASA amount in analyzed

mass0.466 g

3 11.25 mL %ASA in tablet 81.5 %

average volumeused 11.20 mL

% error, assuming 99.3%purity 4.4 %

The calculations for Table 2 are as follows:

For equal equivalents, V1M1= V2M2

11.20 mL (0.0462 mol/L) / 20 mL = M2= 0.0258 M

The amount of ASA per 20-mL sample is then:

0.0258 (mol/L)(180 g/mol)(0.020 L) = 0.093 g or 5.17 x 10

-4

mol of ASA.

In 100 mL, this is equivalent to:

0.093 g (100 mL/20 mL) = 0.466 g of ASA dissolved from the tablet. Since 0.5716 g was the

total amount used, this means that the ASA content is

(0.466 g/0.5716 g) x 100 = 81.5%

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Since the US official standard regulates that each tablet must contain 500 50 mg ASA, we

will assume the ASA content to be 0.5000 g. Then, knowing that only 0.5716 g out of 0.5819

g was transferred, the total ASA dissolved is:

(0.5716 g /0.5819 g) x 0.5000 g = 0.9823 x 0.5000 g = 0.4911 g

The titration gave 0.466 g of ASA, which then gives an error of:

(0.4911 g - 0.466 g) / (0.4911 g) x 100 = 5.1% error

Since the ASA purity is not 100% but 99.3% (see below), the average error decreases to:

{[(0.4911 g x 0.993) - 0.466 g] / (0.4911 g x 0.993)} x 100 = 4.4% error

Results obtained by students (using aliquots of 20 mL) are typically 5% lower than the mean

value calculated from the manufacturers data. Tests with pure acetylsalicylic acid, performed

by a technician using the same method, showed an average value of 98.5% for the pure acidafter 24 repetitions. This result validates the method because in this particular case, the purity

of the acetylsalicylic acid employed was tested to be 99.3%. Lower ASA percentage values

determined by the students are probably due to losses arising from poor quantitative transfer

to the volumetric flask. (Typical aspirin tablets contain 75% to 90% ASA by weight.) There is

also the possibility that the ASA content in aspirin tablets is lower than 500 mg, as has been

reported in the literature.

Questions

1. Provide the data requested in Table 1.

2.

How do your results compare with the sample data provided in Table 2?

3. If your results differ, is the difference significant? If so, discuss possible reasons for

the difference.

Research Questions

1. Drug tablets usually contain many things in addition to the active drug compound.

What is an excipient? What is a filler? What is a binding agent?2. What other types of materials are used in preparing drug tablets? Why are they added?

3. Compare drug administration in tablet form vs. injected drugs. What are the

advantages of tablet dosage? What are some of the difficulties that must be overcomein order to administer a drug as a tablet?

References

This procedure was adapted by Jorge Ibaez (Universidad Iberoamericana, Mexico) from the

original paper, Katia B. Gusmo, Emilse M. A. Martini, and Suzana T. Amaral.

Determination of Acetylsalicylic Acid (ASA) in Aspirin: A General Chemistry Experiment.

Chem. Educator 2005, 10, 444-446.

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