2012 Experifest_chemistry Booklet

8/12/2019 2012 Experifest_chemistry Booklet

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HANDBOOK

Chemistry

EXPERIMENTFEST 2012

www.newcastle.edu.au


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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


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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 Newcastles Faculty of Science and Information Technology and

takes place at both the Callaghan and Ourimbah (Central Coast) campuses of

the University of Newcastle.

Callaghan Campus:

Physics - Mon 18 Thurs 21 June, 2012

Chemistry - Mon 18 Fri 22 June, 2012

Biology - Fri 15 - Fri 22 June, 2012

Ourimbah Campus:

Biology, Physics & Chemistry Mon 25 - Thurs 30 June, 2012

Tuncurry:

Physics- Saturday 23 June, 2012

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 Biology, cementing classroom theory

and providing a good basis for examination preparation.

Experiments include:

Enzymes

Kidney Function

Measuring oxygen satration in the blood

All experiments are complemented by notes, follow-up discussions and

questions to enhance your learning experience.

For booking information contact:

Larry Miltonon 49 469 159 or 0404 460 470; orDavid Rushtonon 4333 6965 or 0414 238 464


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WELCOME

Welcome 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 Technology

University of Newcastle


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CHEMISTRY


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STUDYING CHEMISTRY

Why 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 publicand 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


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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 Universitys

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/


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ANALYSIS OF NITRATE ION

CONCENTRATION IN WATERSAFETY 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.

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

Concentrated HCl Avoid skin contact, wash with copious quantities of water. Wear

safety glasses. Use in fumehood.

Extract from HSC Chemistry Syllabus: Identify that water quality can be determined considering:

- concentrations of common ions

- total dissolved solids

- hardness

- turbidity

- acidity

- dissolved oxygen and biochemical oxygen demand

Perform rst-hand investigations to use qualitative and quantitative tests to analyse and compare the

quality of water samples

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1600/211630/Products/5673991.jpg


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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 fromdecomposing 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 thevisible 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.

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.


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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:

where:

I0 is the intensity of the incident light of definite

wavelength,

It is the intensity of the transmitted light of the same

wavelength,

is the molar absorptivity (in units of L mol-1cm-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:

where:

Measurement of the absorbance of light by solutionsis 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.


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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- 300 nm toselect 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 its absorbance and reading the concentration from the Standard

Curve.

Procedure

1. 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 Conc HCl) and ultra-pure water.

2. Prepare your sample in a similar fashion by diluting 20 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 190-300 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 thenatural water sample.

NO3

-


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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 =


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DETERMINATION OFPHOSPHORIC ACID IN COLA

SAFETY GLASSES MUST BE WORN AT ALLTIMES DURING THE LABORATORY SESSION


Procedure.

Precaution / Action



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

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 rst-hand investigation to determine the concentration of a domestic acidic

substance using computer-based technologies.


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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 xed potential

(Ag/AgCl). The potential of this electrode can be directly

related to pH via an equation of the form:

E = Eglass- 0.05915 pH

where Eglassis 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 theobserved 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 inection 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 rst inection point for H3PO4occurs at complete

conversion to dihydrogen phosphate ion (H2PO4-).

Reaction complete at rst equivalent-point is thus:

H3PO4+ OH- H2PO4

-+ H2O (1)

www.seeinc.com/moreinfo/ AccumetXL25.html


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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 inection 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-

PO43-

+ 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 inection point can be very difcult to observe and we will not be trying

to detect it in this experiment. From the volume of titrant (OH-) consumed

between the rst and second inection 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 2.5

Fraction Titrated


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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. Stirr ing 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 0.10M NaOH

record the exact concentration) sodium hydroxide solution, recording the pH after each

addition. The first two additions of titrant volume (Vb) can large (about 1 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 Vbbased on the change inpH / V. Make sure that you calculate

pH/V for each new Vbaddition 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 0.1mL - 0.2mL.

Drop-wise addition of NaOH in the vicinity of the equivalence point is essential tocreate a satisfactory plot. Make sure that you calculate pH/mL as you go.This

permits precise determination of the end points.

4. Once you have passed through the first end point pH / V 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 1stand

2ndendpoints (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 pH / mL

Example

0.00

5.10

2.55

2.75

2.75 - 2.55 5.10 - 0.00

= 0.04


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mL NaOH added (Vb) Solution pH pH / mL


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

base (NaOH) added (mL).

9. Generate a graph of this data, plotting pH (vertical axis) against Vbvolume 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 (V2on 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 dif ference? 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 pH/mL

verses volume of NaOH (Vb) is prepared. The expression pH/mL 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


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DETERMINATION OFSULFATE IN LAWN FOOD



Procedure.

Precaution / Action

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

Ammonia/Ammonium

Chloride solution

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

Use in fumehood


BaCl2solution. Harmful by inhalation. Toxic if swallowed. Skin contact may produce

health damage. Avoid skin contact. Wash affected area well.

Eriochrome Black T

indicator

Cancer suspect agent. Cumulative effects may result following

exposure. Avoid skin contact. Wash affected area well.

Hot Plate: 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 chemistry

involved

Analyse information to evaluate the reliability of the results of the aboveinvestigation and to propose solutions to problems encountered in the

procedure.

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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 particularlycomplex 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.

TitrationsIn carrying out the titrations, use two samples (i.e. carry out the titrations in duplicate) and record your values to 2 deci-

mal 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 re-

lying on this value to determine the remaining barium, this figure will be artif icially inflated. Thus when this inflated 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 exactconcentration 4 decimal points)

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

Eriochrome Black T indicator solution.

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 thecarboxyl units of EDTA have been deprotonated, thereby enabling the metal complex to form.


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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/ammonium chloride solution and 3-4 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 BaSO4precipitate

from the first sample using the Millipore filtration apparatus on you bench. Make sure that you

place a 0.45m filter paper in the apparatus before commencing to f ilter.

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 10mL concentrated ammonia/

ammonium chloride solution and 3-4 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 Vtin 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.

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.


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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

Titration 2 ...mL

Average Titre: ...mL

Volume of EDTA required for test solution (Vt)

Titration 1 ...mL

Titration 2 ...mL

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?


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METAL ANALYSIS BY ATOMIC ABSORPTION SPECTROSCOPY:

CONCENTRATION OF SODIUM INSPORTS DRINKS



Procedure.

Precaution / Action




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 specic

absorbance data provided to read off the corresponding concentration for the

sample.

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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 includingpollutants - 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 limitsfor metals vary considerably and

depends on the method used to atomise the sample. The

most common method for atomising samples is by ame

using an acetylene / air mixture. This produces a ame

temperature of about 2,500C, 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 thedetection 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 [ ]) 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 aninterference.

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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 ame, 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

ame, 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 ofthe 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


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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 fourseparate 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 5mL aliquot (precisely known volume use a pipette) of

Gatorade into a 100.0 mL volumetric ask. 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 asks and thoroughly mix the solutions by inverting the asks

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 theabsorption 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.


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Results

Table 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 sportsdrink 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 briey 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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2012 Experifest_chemistry Booklet