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HANDBOOKChemistry

2010

EXPERIMENTFEST

www.newcastle.edu.au

Experiment Fest 2010

CONTENTS

Introduction 3

Welcome 4 Studying Chemistry 6

Analysis of Nitrate Ion Concentration in Water 8

Determination of Phosphoric Acid in Cola 13

Determination of Sulfate in Lawn Food 20

Metal Analysis by Atomic Absorption Spectroscopy: Concentration Of Sodium In Sports Drinks 25


INTRODUCTIONExperimentFest is an experiment program designed to provide enriching educational experiences for senior high school students who are studying Physics, Chemistry and Biology. ExperimentFest is supported by the University of Newcastle’s Faculty of Science and Information Technology and takes place at both the Callaghan and Ourimbah (Central Coast) campuses of the University of Newcastle. It is run at the Callaghan Campus from 21-25 June 2010 for Physics and Chemistry; and 18-25 June 2010 for Biology; and at the Ourimbah Campus from 28 June– 1 July 2010. Students may attend either morning sessions (9am-12pm) or afternoon sessions (12:45pm-3:45pm). Physics Experimentfest is also running at Tuncurry (June 26 2010).

The activities allow students to engage in a range of hands-on experiments that are difficult to organise within a school setting, all under the supervision of University Staff and Postgraduate students. Each experiment is chosen to complement the NSW HSC syllabus for Chemistry, cementing classroom theory and providing a good basis for examination preparation.

Experiments include: Determination of phosphoric acid in cola Metal analysis by atomic absorption spectroscopy: concentration of sodium in sports drinks

Determination of sulphate in lawn food An analysis of nitrate ion concentration in water.

All experiments are complemented by notes, follow-up discussions and questions to enhance your learning experience.

For booking information contact:Larry Milton on 49 469 159 or 0404 460 470; or David Rushton on 4333 6965 or 0414 238 464.


WELCOMEWelcome to the Faculty of Science and Information Technology at the Univer-sity of Newcastle. ExperimentFest is a wonderful chance to give you practical experience which complements your classroom learning while giving you a first hand look at University life and facilities. Science is an exciting field of study, allowing you to move with the times and contribute actively and respon-sibly to society. There are many education opportunities in science after high school. Here in the Faculty we provide study and research programs in fast-moving modern fields that make our world work.

The Faculty staff and students who will be taking you through the experiments today are involved in contemporary science research. Please ask questions and utilise your time with them.

Take this day to enjoy being out of the class room, exploring science with fel-low students and participating in valuable experiments and discussions which will help you in your HSC and beyond.

I wish you well in your studies. I hope you apply yourselves to the learning process with enthusiasm and you enjoy your time at the University. We hope to see you studying with us in the future!

Best wishes,

Prof. Bill Hogarth Pro Vice-Chancellor – Faculty of Science and Information TechnologyUniversity of Newcastle

CHEMISTRY


STUDYING CHEMISTRYWhy study chemistry?Chemistry is truly the “central science”. It is the science of the molecular scale, and it is at the molecular level where major advances are being made in many diverse areas such as medicine, drugs, nanotechnology, new materials, and the environment. A sound knowledge of chemistry is required to fully understand many other areas of science, and this is why the study of chemistry is either compulsory or recommended by many other disciplines in the University. Chemistry opens the door for many careers because training in chemistry is essential for many positions in industry, is highly desirable for science teaching, and is useful for careers in the public service and management. Both the public and the private sectors increasingly draw their higher management echelons from chemistry graduates.But, most importantly, it is just so fascinating! If you want to understand the workings of the world around you - then chemistry is for you!Opportunities for further studies in Chemistry:The Bachelor of Science degree program at the University of Newcastle provides a foundation of knowledge, skills and attributes that allows graduates to be employable not just today but into the future and to contribute actively and responsibly to society. Majoring in Chemistry, you have the opportunity to sample and/or specialise in any one of the following: Analytical Chemistry Chemistry Chemometrics (double major) Environmental Chemistry Forensic Chemistry Geological Chemistry Industrial Chemistry Materials Chemistry Medicinal Chemistry Surface and Colloid Chemistry

Research in Chemistry at the University of Newcastle: There are various groups here at the University which are committed to research in chemistry. Groups include:

Analytical and Environmental Chemistry Battery Materials and Applied Electrochemistry Coordination and Bioinorganic Chemistry Marine Natural Products and Chemical Ecology Polyoxymetalates and Catalysis Surface and Colloid Group


“For me chemistry represented an indefinite cloud of future potentialities…” - Primo Levi

Careers in Chemistry: The Faculty of Science and IT care about our students and are interested in giving as much direction as possible to those making career choices and beyond. The possible career paths listed below include a range of opportunities for graduates at degree, honours, and post graduate study levels.

Analytical Chemist Clinical Research Coordinator Developmental Chemist Environmental Chemist Energy Technologist Forensic Chemist Geochemist Industrial/Production Chemist Laboratory Manager Laboratory/Research Assistant Meteorologist Organic/Synthetic Chemist Pharmaceutical/Medicinal Chemist Reproductive Medicine / IVF Chemist Research Scientist Science Information/Education Officer Science/Chemistry Teacher Sciences Technician Scientific Patent Attorney / Technical Advisor Scientific Policy Officer Scientific Writer

For more information on these career paths, please visit the University’s careers website: www.newcastle.edu.au/service/careers/majors/

For more information on the Faculty of Science and IT check out our website:http://www.newcastle.edu.au/faculty/science-it/


Experiment Fest 2010Page 8

ANALYSIS OF NITRATE ION CONCENTRATION IN WATER

SAFETY GLASSES MUST BE WORN AT ALLTIMES DURING THE LABORATORY SESSION

RISK ASSESSMENTHazard Substance, Apparatus,

Procedure.Precaution / Action

Burette filling: Use a plastic funnel to fill burette; remove equipment from retort stand and fill over sink.

Ammonia solution Toxic by inhalation. Causes burns. Risk of serious damage to eyes.

Pipette filling: Do not pipette by mouth. Use a pipette bulb to fill pipette.

BaCl2 solution. Harmful by inhalation. Toxic if swallowed. Skin contact may produce health damage. Avoid skin contact. Wash affected area well.

Eriochrome Black Tindicator

Cancer suspect agent. Cumulative effectsmay result following exposure. Avoid skincontact. Wash affected area well.

Extract from HSC Chemistry Syllabus:• Identify data, plan, select equipment and

perform firsthand investigations to measure the nitrate content of an unknown sample using spectrophotometry

• Analyse information to evaluate the reliability of the results of the above investigation and to propose solutions to problems encountered in the procedure.

http://altura.speedera.net/ccimg.catalogcity.com/210000/211600/211630/Products/5673991.jpg


IntroductionNitrate and phosphorus are together known as nutrients when found in water bodies. These two species, if present in sufficient quantities, allow excessive plant growth and eventual stagnation of the water. This process is known as eutrophication. Nitrogen is tied up in biological systems. Nitrate and nitrite ions are important indicators of pollution by organic materials as nitrogen from decomposing organic substances often ends up as nitrate or nitrite ions. The determination of nitrate is often difficult because of the low levels found, and the distinct possibility of interfering materials being present.

SpectrophotometryWhen determining nitrate it is important to choose an analytical method that suits both the interferences that are present and the level of analyte in the solution. Spectrophotometry is an excellent analytical method. In visible spectrophotometry, light impinges on a coloured substance causing certain wavelengths of light to absorbed. The remaining wavelengths of light in the visible spectrum are reflected and transmitted to the eye of the viewer. For example, a purple dye absorbs wavelengths in the green-yellow parts of the visible spectrum and reflects red and blue wavelengths of light.

Figure 1: Absorbance spectrum of a cyanine dye

max = 558 nm

1,1'-diethyl-2,2' cyanine iodide

N NH3CH2C CH2CH3

Spectroscopy can also be carried .be carried out using other parts of the electromagnetic spectrum such as the ultra-violet and infrared regions. The principle of evaluation remains the same however, provided you have a device for detecting the intensity of the radiation in the chosen region of the spec-trum. In this experiment you will utilize UV radiation to quantify the amount of nitrate ion in an unknown sample.

The Beer-Lambert Law.The intensity of a coloured solution depends upon the concentration of the coloured component and the path length through which the light travels. Laws describing the absorption of light were originally formulated by Lambert (1760) and Beer (1852), and today are combined to give a single law, which quanti-tatively relates the extent of absorption of light to the concentration of absorb-ing species in solution, known as the Beer-Lambert Law. Importantly this law refers to a single wavelength and not to the absorption band as a whole.


The Beer-Lambert Law states that, at a given wavelength, the relationship between the intensity of the incident and transmitted monochromatic light in terms of the concentration of the absorbing species in solution and the pathlength is given by:

I tI o

= 10 - cd

I o I t

d

concentration (c)

where:I0 is the intensity of the incident light of definite

wavelength,

It is the intensity of the transmitted light of the same wavelength,

ε is the molar absorptivity (in units of L mol-1 cm-1) at that wavelength. It is a constant for a pure substance,

c is the concentration of the absorbing substance (in mol dm-3, i.e. mol L-1),

d is the pathlength (in cm) of the absorbing substance.

Using mathematics (a log transformation) this equation may be simplified to a linear form:

A = cd

I oI t

log10A = absorbance

where:

Measurement of the absorbance of light by solutions is accomplished using a spectrophotometer. In a spectrophotometer a light source emits a range of wavelengths which are passed into a monochromator. The monochromator in turn selects a single wavelength of light which is passed through the sample to be measured. The incident light beam is absorbed to some extent by the sample and the intensity of the transmitted light (of the same wavelength as the incident light) is measured by a photoelectric device.

The electrical circuit is so designed that the output of the instrument is calibrated to read directly in absorbance or percentage transmission.

polarisersource

sample

detector

spectrum

signalprocessing

diffraction grating


In this experiment you will use a spectrophotometer to determine the concentration of an unknown nitrate containing solution in the UV region of the electromagnetic spectrum. Good results may be obtained if the water contains little or no organic matter (which also absorb in the UV). The method will involve taking an absorbance spectrum of nitrate ion in the range 190- 400 nm to select the strongest absorbance wavelength for your measurements. You will then construct a calibration plot, otherwise known as a Standard Curve (see below) from the absorbances of the nitrate samples of known concentration.

The concentration of the unknown NO3- sample will then be determined by

measuring it’s absorbance and reading the concentration from the Standard Curve.

Abs

orba

nce

[NO3-] M

known absobances of standard solutions.

unknown absobance.

Procedure1. Prepare nitrate calibration standards for 0, 5, 10 and 15 mg/L by adding 0,

5, 10 and 15 mL, respectively, of the 100 mg/L stock nitrate solution, and then making them up to the mark in a 100 mL volumetric flask with HCl (i.e. 1mL of 0.1M HCl) and ultra-pure water.

2. Prepare your sample in a similar fashion by diluting 5 mL of the natural water sample in a 100 mL volumetric flask with HCl and ultra-pure water.

3. Using your 0 mg/L standard record a background absorbance spectrum in the range 190-400 nm.

4. Measure the absorbance spectrum of the remaining standards and sample over the same wavelength range.

5. Prepare a calibration graph by plotting the absorbance maximum in each spectrum and hence determine the concentration of nitrate ion in the natural water sample.


ResultsSolution Wavelength Maximum (nm) Absorbance at Maximum

0 mg/L standard

5 mg/L standard

10 mg/L standard

15 mg/L standard

Unknown

Unknown Concentration =


DETERMINATION OF PHOSPHORIC ACID IN COLA

SAFETY GLASSES MUST BE WORN AT ALLTIMES DURING THE LABORATORY SESSION




pH electrodes Use with care. Remove electrode from solution when stirring.


0.1 M NaOH solution. Avoid skin contact, wash with copious quantities of water. Wear Safety Glasses

0.1 M H3PO4 solution. Harmful by inhalation. Causes burns. Risk of serious damage to eyes. Wear safety glasses.

Extract from HSC Chemistry Syllabus:The Acidic Environment Contextual Outline (extract)Acids such as ascorbic acid (vitamin C) and citric acid occur in many foods. Many drinks contain carbonic acid and some contain phosphoric acid. Other acids, such as benzoic acid and acetic acid, are added to drinks and food to act as preservatives. Hydrochloric acid is secreted into the human stomach to assist in the digestion of food, especially of proteins to amino acids.

Extract from Chemistry Stage 6 Syllabus:Perform a first-hand investigation to determine the concentration of a domestic acidic substance using computer-based technologies.


IntroductionDetermination of the concentration of acids or bases is routinely carried out by titration using an acid / base indicator such as phenolphthalein. Selection of the correct indicator is important to ensure that the endpoint (at which point the indicator changes colour) corresponds with the equivalence point (the exact point where the acid or base is neutralized) in the titration. In situations where coloured solutions are involved, determination of an end-point by a colour change is not practical. Under these circumstances, a pH electrode (shown right) is often used.

Given that cola is coloured, we will use a pH electrode to determine the concentration of one of the components of cola, phosphoric acid, by titration with dilute sodium hydroxide. Titrations where a pH electrode is used are often termed potentiometric titrations.

PrincipleThe electrical potentials of the electrode in the pH meter (shown right) is dependent upon the hydrogen ion concentration of the solution in which it is immersed. The thin walled glass bulb at the end of the electrode which contains a platinum contact immersed in a buffer solution coupled to a reference electrode of fixed potential (Ag/AgCl). The potential of this electrode can be directly related to pH via an equation of the form:

E = E°glass - 0.05915 pH

where E°glass is a non-constant potential characteristic of the glass electrode. The equation shows that the electrical potential of the probe is therefore a function of pH, consequently the electrode may be used to monitor changes in solution pH. An incremental plot of the observed pH readings against volume of titrant (NaOH in this experiment) added, yields a curve with one or more regions in which there is a rapid change in pH. These inflection points correspond to the completion of proton-transfer equilibria for H3PO4; that is, they indicate titration equivalence points. Further treatment of the data with Excel permits a more accurate determination of the equivalence points.

The first inflection point for H3PO4 occurs at complete conversion to dihydrogen phosphate ion (H2PO4

-). Reaction complete at first equivalent-point is thus:

H3PO4 + OH- H2PO4- + H2O (1)

www.seeinc.com/moreinfo/ AccumetXL25.html


This point is significant, as there is a rapid increase in pH with only a small addition of base. When the data is plotted, this first inflection point will be obvious. As this is the first equivalence point (see plot below), it should be approached carefully so that a precise measure of the volume of base required can be obtained.

Continued addition of base results in the conversion of the dihydrogen phosphate species to the mono-hydrogen form and a second inflection point appears on the titration curve. i.e.

H2PO4- + OH- HPO4

2- + H2O (2)

Again, this second equivalence point will be obvious, as there is a rapid change in pH with the addition of only a small volume of base. Further addition of NaOH leads to the formation of phosphate ion:

HPO42- + OH- PO4

3- + H2O (3)

An animated graph showing the change in concentration of the species present in a phosphoric acid titration can be viewed at the following website:http://chemistry.beloit.edu/Rain/moviepages/phosphoric.htm. The last inflection point can be very difficult to observe and we will not be trying to detect it in this experiment. From the volume of titrant (OH-) consumed between the first and second inflection points, the concentration of the phosphoric acid component can be calculated.

Typical Titration Curve for Phosphoric Acid.14

12

10

8

6

4

2

0

pH

0 0.5 1 1.5 2 2.5Fraction Titrated


Procedure: Phosphoric acid content of Coca-ColaThe beverage “Coca-Cola” contains significant amounts of phosphoric acid and phosphate salts and these can be evaluated by titration. However, since beverages also contain carbonic acid, a necessary preliminary step is elimination of this species. To shorten the lab, the Coke has been boiled already.

pH ELECTRODES ARE FRAGILE AND VERY EXPENSIVE! They are easily damaged, so please take care when handling them !

1. Take a 100.0mL aliquot of the boiled beverage in a 250mL beaker. 2. Place the beaker on the magnetic stirrer, add the stirrer “bug” (more correctly a magnetic

follower) and carefully clamp the burette into position above the beaker (so that you can comfortably add base) and the pH electrode into position. Turn the stirrer on and make sure that it does not come into contact with the pH electrode. Stirring should be a steady, constant rate – not too rapid. This will help achieve effective mixing with each addition of the base.

3. With the pH electrode in place, titrate the solution with standard (approximately 0.05M – record the exact concentration) sodium hydroxide solution, recording the pH after each addition. The first two additions of titrant volume (Vb) can large (about 5 mL), as the pH change at this point in the titration is only small (see diagram on page 15). The volume added should however be decreased beyond this point the pH will begin to rise rapidly. You will need to make adjustments to Vb based on the change in ΔpH / ΔV. Make sure that you calculate ΔpH/ΔV for each new Vb addition as this will show when you are nearing the titration end-point (by becoming large). When the pH is changing rapidly record the volume of addition should be reduced to no more than 0.2mL.

Drop-wise addition of NaOH in the vicinity of the equivalence point is essential to create a satisfactory plot. Make sure that you calculate ΔpH/ΔmL as you go. This permits precise determination of the end points.

4. Once you have passed through the first end point ΔpH / ΔV will fall off in magnitude. Have the supervisor check your results at this point, then proceed to ascertain the second endpoint using the same approach in step 3. The volume of NaOH may again be increased in the initial stages (~5 mL) until you approach the second end point. The volume of NaOH added between the 1st and 2nd endpoints (V2) will be about the same as that observed for the first equivalence point (V1).

5. When the second inflection point has been reached, turn the magnetic stirrer off, remove the pH electrode from the solution, wipe carefully with a clean tissue and place in the original flask

of distilled water. Leaving pH electrodes in solutions of high pH can damage them. 6. The titrated solution can be discarded down the sink.

DO NOT THROW THE BUG OUT WITH THE SOLUTION! Replace it on the magnetic stirrer.

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Results: Concentration of NaOH solution ………… M

mL NaOH added (Vb) Solution pH ΔpH / ΔmL

Example 0.00 5.10

2.552.75

2.75 - 2.55 5.10 - 0.00 = 0.04


mL NaOH added (Vb) Solution pH ΔpH / ΔmL


Data Handling / Calculations8. Enter your results into an Excel spreadsheet under the column headings pH and Vb mL, where Vb is the volume of

base (NaOH) added (mL).9. Generate a graph of this data, plotting pH (vertical axis) against Vb volume of base added (mL). Label this plot “pH

verses volume of NaOH added”.

Results10. From the graph, identify the volume of NaOH required for the titration of H3PO4

- the first equivalence point (V1

on graph p15). Use this volume to calculate the concentration of the phosphoric acid present in the cola. Ex-press your answer in moles per litre of cola (i.e. the molarity). Repeat this calculation using the volume of NaOH

required for the neutralisation of H2PO3- (Eq. 2) - the NaOH volume between the first and second equivalence

points (V2 on graph p15).

Follow-up activities:1. Was there any difference between the volumes of NaOH required to achieve the first and second equivalence

points in the titration? What could account for this difference? Which of the NaOH volumes is likely to be the more accurate? Explain your choice.

2. Why was the cola sample boiled prior to analysis?3. A more accurate way to estimate the equivalence point is to perform a differential plot. Here a plot of ΔpH/ΔmL

verses volume of NaOH (Vb) is prepared. The expression “ΔpH/ΔmL“ shows the rate of change of slope along the titration curve yielding a very accurate end point. (see below). Instructions on how to prepare a differential plot us-ing Microsoft Excel may be found at.: A differential plot of a titration curve may be viewed at

http://facstaff.bloomu.edu/eschultz/chem%20116%20spring%2006/lab/Generating%20Titration%20Curve.pdf

7.00

6.00

5.00

4.00

3.00

2.00

1.00

0.00

12

10

8

6

4

2

0

pH

0 10 20 30 40 50 60mL Base added

H 7


DETERMINATION OF SULFATE IN LAWN FOOD

SAFETY GLASSES MUST BE WORN AT ALL TIMES DURING THE LABORATORY SESSION




Ammonia solution Toxic by inhalation. Causes burns. Risk of serious damage to eyes.


BaCl2 solution. Harmful by inhalation. Toxic if swallowed. Skin contact may produce health damage. Avoid skin contact. Wash affected area well.

Eriochrome Black Tindicator

Cancer suspect agent. Cumulative effects may result following exposure. Avoid skin contact. Wash affected area well.

Hot water bath: Burn risk. Avoid skin contact.

Extract from HSC Chemistry Syllabus:Identify data, plan, select equipment and perform firsthand investigations to measure the sulfate content of lawn fertilizer and explain the chemistryinvolved

Analyse information to evaluate the reliability of the results of the above investigation and to propose solutions to problems encountered in the procedure.http://altura.speedera.net/ccimg.catalogcity.com/210000/211600/211630/Products/5673991.jpg


IntroductionThe traditional method of determining sulfate is by precipitation as the barium sulfate, filtering, “ashing” the filter paper (i.e. heating until the filter paper is turned to ash with effectively zero weight) then weighing the residual matter. This is termed a gravimetric method, as it relies on weighing. While not a particularly complex procedure, it is time-consuming. The solution has to be boiled for nearly 1 hour to ensure that the barium sulfate precipitate forms particles large enough to be successfully filtered. This process is termed digestion.

An alternative method, the one we will use, involves the addition of a known number of moles (excess) of barium ions (added as a soluble barium salt), which then precipitates as barium sulfate. After removing the precipitate, the residual barium in solution may be determined by titration with ethylenediamine tetraacetic acid (EDTA). This method is termed a “back titration”. EDTA is a strong chelate - a species that reacts with metal ions to form a stable metal complex. It reacts in a 1:1 ratio with barium forming a soluble complex (see equation). EDTA titrations are carried out under basic conditions, often buffered to maintain a highly basic pH environment. This is to ensure that all of the OH protons in the carboxyl units of EDTA have been deprotonated, thereby enabling the metal complex to form.

To assist the completion of the experiment, samples of the lawn fertilizer pretreated with barium ion and predigested will be available. You will prepare fresh samples for the group following you.


TitrationsIn carrying out the titrations, use two samples (i.e. carry out the titrations in duplicate) and record your values to 2 decimal places (that is, estimate between the graduations on the burette). Ideally, the two results should be within 0.2 mL to be acceptable. Record your results in the Table below.

To simplify calculations, a “blank” titration is carried out first, where the volume of EDTA required to react with the Ba2+ ions before the reaction with sulfate is determined. Record all titration values to second decimal point.

InterferencesThis method will not give a meaningful result in the presence of phosphate, as the solid barium phosphate will also form, leading to an incorrect estimation of the amount of barium ions which reacted with the sulfate.

In addition, the presence of other metal ions (such as copper, zinc, etc) will lead to low sulfate estimation, as these ions also react with the EDTA. Their presence would lead to more EDTA being used than for the barium alone. As we are relying on this value to determine the remaining barium, this figure will be artificially inflated. Thus when this inflat-ed value is subtracted from the initial moles of barium added, we would underestimate the actual sulfate concentration.

Supplied solutions:Fertiliser solution of known concentration (6.6 g/L)

0.2 M Barium chloride BaCl2

0.1 M EDTA solution (Note: record the exact concentration – 4 decimal points)

Concentrated ammonia solution NH3 (CAUTION: see safety note above)

Eriochrome Black T indicator solution.


Procedure:

Steps 1 and 2 should be carried out in duplicate (2 samples). Get steps 1 and 2 underway before commencing the titration.

1. Pipette 1 x 50.0mL of the supplied ammonium sulfate solution into two separate beakers.2. Pipette 25.0mL of 0.2M barium chloride into each beaker, then cover the solution with a

watchglass and place on the hotplate in the fume cupboard. These solutions will be used by the group undertaking the experiment after you. Take two of the predigested samples from the hotplate for your experiment.

3. Return the fertilizer solutions to your bench, place them in an ice bath to cool. While this is taking place, carry out a blank titration on the 0.2M barium chloride solution as follows:

• Pipette 25.00mL of 0.2M barium chloride into a 250 mL conical flask • Add 10mL of concentrated ammonia, 0.5g of ammonium chloride and a 6 drops of

indicator • Titrate with 0.1M EDTA from red/purple to a blue endpoint. • Record the volume of EDTA required in the Results Table below as Vb.4. Repeat step 3 to obtain a duplicate result for the blank titration. Record the volume in the table

and average the two blank titration results.5. Remove the now cooled fertilizer solutions from the ice bath and filter the BaSO4 precipitate

from the first sample using the Millipore filtration apparatus on you bench. Make sure that you place a 0.45μm filter paper in the apparatus before commencing to filter.

This apparatus is very expensive. Take great care! Make sure that you use the filter paper and not the separating paper! Check with the demonstrator if you are unsure.

6. Transfer the clear filtrate to a 250mL conical flask. Add 0.5g of ammonium chloride, 10mL of concentrated ammonia and 8-10 drops of the indicator solution.

7. Repeat steps 5 -7 for the second fertilizer sample.8. Titrate sample 1 with 0.1M EDTA until the colour changes from red/purple to blue. Record this

volume to 2 decimal places as Vt in the Results Table below.9. Repeat step 8 with the second sample of the fertilizer solution.10. Use the information in the Results Table to calculate the percentage of sulfate ion in lawn

fertilizer.


Results:

Concentration of fertiliser (g/L) 6.6 g/L

Concentration of EDTA (exact concentration) ………………… M

Volume of EDTA required for blank (Vb)

Titration 1 …………..….mL


Average Titre: …………..….mL

Volume of EDTA required for test solution (Vt)



Average Titre: …………..….mL

Difference in volume of EDTA (corresponding to Ba2+

used, thus SO42- present) - Vb - Vt

…………..….mL

... moles of EDTA consumed. ………………moles

... moles of sulfate ………………moles

Initial mass of fertiliser in 50.0 mL sample ……………....g

Mass of sulfate determined experimentally ……………....g

% sulfate in the lawn food

Follow-up activities: 1. Calculate the percent of sulfate if the sample had been pure ammonium sulfate. Account for any differences.

2. Can you suggest any changes in the experimental procedure that may improve the accuracy?

3. What was the reason for adding the ammonium chloride?

4. Why is the concentration of the barium chloride not required in the calculations?

5. How could you minimise the interferences identified above?


METAL ANALYSIS BY ATOMIC ABSORPTION SPECTROSCOPY:

CONCENTRATION OF SODIUM IN SPORTS DRINKS

SAFETY GLASSES MUST BE WORN AT ALL TIMES DURING THE LABORATORY SESSION





Extract from HSC Chemistry Syllabus:Some foodstuffs are monitored for the presence of particular metals, e.g. lead, to ensure that we are not consuming poisons that can accumulate in our bodies.

Gather, process and present information to interpret secondary data from AAS measurements and evaluate the effectiveness of this in pollution control.

AAS allows the detection of very small concentrations from samples of air, water or food. This activity depends on your ability to manipulate data and dilution factors. The absorbance values obtained using solutions of known concentration enable you to draw a calibration graph. Use the specific absorbance data provided to read off the corresponding concentration for thesample.

http://www.suntory.com/news/2004/img/8648.jpg


IntroductionAtomic Absorption Spectroscopy (AAS) is a commonly used technique for determining the concentration of metal ions. This analysis of metals is important in a range of applications from trace analysis of metals – including pollutants - in water and soil to quality control of materials such as iron and special steels.

Quality control in products is an important commercial tool. In this instance, the presence of salts such as sodium in sports drinks is critical for their usefulness in replacing the salt loss by athletes through perspiration. Too little salt and the product will not be effective. However, the addition of too much salt may diminish the taste of the product, as well as slow the rate of absorption of the liquid in the stomach.

The detection limits for metals vary considerably and depends on the method used to atomise the sample. The most common method for atomising samples is by flame using an acetylene / air mixture. This produces a flame temperature of about 2,500°C, which is a relatively low temperature for this technique. Under these conditions, the detection limit for iron is 5 ng/L (or 5 parts per billion, also expressed as 5 ppb), copper – 1ppb; zinc – 1.5 ppb and mercury – 150 ppb. While these are the detection limits, accurate quantitative analysis requires concentrations 10 to 100 times greater than the detection limit.

The basis of the method is that for a given metal ion, the absorption is directly proportional to the concentration (A α [ ]) of the species present. We use this principle to prepare a calibration curve and use this to determine the concentration of the species in the unknown.

As with any analytical method, the determination of metals by AAS can be affected by the presence of other species. It is important that we are aware of these and take steps to counter their effect. Any species that changes the observed signal, while the concentration of the metal remains unchanged, is termed an interference.

http://www.pelletlab.com/images/2303.jpg


Types of InterferenceIn this experiment we are primarily interested determining the sodium content of sports drinks by AAS. In addition, we are interested in the possible interference from the presence of potassium in obtaining the correct reading for the sodium. The sports drink sample you will analyse has both sodium and potassium present. In AAS, it is very important that the standards we prepare are as close as possible to the sample under investigation.

Because alkali metals such as sodium have low ionisation potentials, they readily ionise in the flame, changing the species responsible for the absorption and leading to errors in estimating the concentration. This is termed ionisation interference. The simplest way to overcome this effect is to add another alkali metal that will ionise more easily than the sodium, such as potassium. This is termed an ionisation suppressor. By having a high concentration of electrons in the flame, ionisation of the sodium is suppressed.

PrincipleYour analysis will determine the concentration of sodium present in a range of sports drinks. A series of standard solutions of sodium will be prepared. Our selection of the range of these standards is based on our understanding of the concentration of the unknown as well as the knowledge that high concentrations of substances cause deviations from the linear relationship between absorption and concentration.

Preparation of the sample for analysis involves dilution of the initial sports drink sample. This is required to adjust the concentration of sodium into the correct range for the operation of the instrument and to ensure that the absorption reading is within the range of the standards.

Supplied solutions:1,000 mg/L Sodium (NaCl) Stock10,000 mg/L Potassium (KCl) Stock


ProcedurePreparation of Standards1. Prepare a series of four standard sodium solutions containing 0, 10.0,

20.0, 50.0mg/L (ppm) of sodium. To do this, measure out 0, 1, 2 & 5mL of the sodium stock solution (from the supplied burette) respectively into four separate 100mL volumetric flasks and make each flask up to the mark with distilled water. Label them as as “Na stds-1, 2,” etc.

2. Prepare a series of sodium standards containing a constant amount of potassium ions (as KCl).

3. Prepare the same series (i.e. 0.0, 10.0, 20.0 and 50.0 mg/L) as in Step 1 above by measuring out the same volume of sodium stock solution into four new volumetric flasks. At this point, add 20mL of the 10,000 mg/L potassium (KCl) stock solution (measured with a measuring cylinder) to each volumetric flask and make up to the mark with distilled water. Label this series as “Na-K stds-1, 2,”etc.

Preparation of Sports Drink Sample.4. Pipette a 10.0 mL aliquot (precisely known volume – use a pipette)

of Staminade into a 100.0 mL volumetric flask. If Gatorade is used, dilute 5mL into the volumetric flask. Add 20mL of the 10,000 mg/L potassium stock solution (measured with a measuring cylinder), then add distilled water to the 100.0 mL mark. (What dilution factor have you just introduced?).

5. Stopper all flasks and thoroughly mix the solutions by inverting the flasks at least 6 times.6. Analyse these solutions by AAS for sodium as described below.

Analysis7. Measure the absorbance of the “Na stds” series (all 4 solutions). Solutions

are measured from lowest to highest concentration. Record these results in the Table below. From the data, prepare a plot of Absorption Vs Concentration Na stds.

8. Measure the absorbance of the “Na-K stds” series (all 4 solutions). Solutions are measured from lowest to highest concentration. Record these results in the Table below. From the data, prepare a plot of Absorption Vs Concentration Na-K stds.

9. Measure the absorbance of the diluted sports drink. Record the absorption value in the Table below.

10. From the absorption reading, determine the concentration of Na from each calibration graph.

You must take the dilution factor into account when reporting the actual concentration of sodium in the sports drink.


ResultsTable 1: Absorbance Values of Na standards

Concentration of Na stds(mg/L)

AAS Absorbance

Na-stds Na-K stds

0.0

10.0

20.0

50.0

Table 2: Absorbance of Sports Drink. Name of sports drink

Sodium concentration from bottle*

AAS Absorbance.

Concentration of sodium in sports drink determined from ‘Na stds’ graph.

Concentration of sodium in sports drink determined from ‘Na-K stds’ graph.

*Express all values in mg/L

Follow-up activity:1. Which of the two values you have reported in the Table above would be

expected to be closer to the “real” value? Justify your choice.2. If you were asked to determine the concentration of potassium in the

sports drink, what measures would you have to take to minimise ionisation interference?

3. What is the actual concentration of potassium used in the ‘Na-K stds’ series (express your answer in mg/L or ppm)?

4. Draw a block diagram of an AAS unit and use this to briefly explain how the system operates.

5. What procedure would you use to determine the amount of copper and arsenic which could be leached from a Koppers log?

6. How could a high concentration of sodium in the sports drink slow down water absorption by the stomach?

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