EXPERIMENTFEST
2015
HANDBOOK
Chemistry
CONTENTS
Introduction
3
Welcome
4
Studying Chemistry
6
Metal Analysis by Atomic Absorption Spectroscopy:
Concentration of Sodium in Sports Drinks
8
Determination of Sulfate in Lawn Food
13
Analysis of Nitrate Ion Concentration in Water
18
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INTRODUCTION
ExperimentFest 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.
Callaghan Campus:
Physics - Friday 12 – Thursday 18 June, 2015
Chemistry - Monday 15 – Friday 19 June, 2015
Biology - Friday 12 - Friday 19 June, 2015
Ourimbah Campus:
Biology, Physics & Chemistry Monday 22 - Thursday 25 June, 2015
Tuncurry:
Physics - Saturday 20 June, 2015
All experiments are complemented by notes, follow-up discussions and
questions to enhance your learning experience.
For booking information contact:
Larry Milton on 0404 460 470; or
David Rushton on 4333 6965 or 0414 238 464
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WELCOME
Welcome to the Faculty of Science and Information Technology at the University 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 responsibly to
society. There are many education opportunities in science after high school. Here in the Faculty we
provide study and research programs in fastmoving 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 fellow 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. Eileen McLaughlin
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 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:
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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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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.
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•
•
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•
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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 please visit the University’s website:
www.newcastle.edu.au
For more information on the Faculty of Science and IT check out our website:
www.newcastle.edu.au/science-it
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METAL ANALYSIS BY ATOMIC
ABSORPTION SPECTROSCOPY:
CONCENTRATION OF SODIUM IN
SPORTS DRINKS
SAFETY GLASSES MUST BE WORN AT ALL
TIMES DURING THE LABORATORY SESSION
RISK ASSESSMENT
Hazard
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.
Extract from HSC Chemistry Syllabus:
Manufactured products, including food, drugs, and household chemicals, are analysed
to determine or ensure their chemical composition.
Gather, process and present information to interpret secondary data from
AAS measurements and evaluate the effectiveness of this in pollution control.
Atomic Absorption Spectroscopy (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 the sample.
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http://www.differencebetween.net/
object/difference-between-gatorade-and-powerade/
Introduction
AAS is a commonly used technique for determining the concentration of metal
ions. The 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.
A flame is needed to atomise a sample in order to measure it. The temperature
of the flame is selected based on what metal(s) are being analysed. As shown in
Table 1, the less reactive metals require more energy (heat) to be atomised. The
most common method for atomizing a sample is using an acetylene (C2H2) / air
mixture.
http://www.pelletlab.com/images/2303.jpg
Table 1: Flame gas temperatures and applications.
Gas Combination
Natural gas/air
Acetylene/air
Acetylene/N2O
Max. Temp. (K)
1800
2250
3000
Elements atomized
Group 1A (Li, Na, K)
Transition metals (Fe, Cu, Zn)
Group 2A (Ca, Mg, Ba) Group 3A (Al)
The detection limit between metals varies considerably. Under acetylene/air conditions, the detection limit for iron is 5 µg/L
(or 5 parts per billion, also expressed as 5 ppb), sodium – 0.3 ppb; zinc – 1.5 ppb and mercury – 150 ppb.
The basis of this analytical method is that for a given metal ion, the absorption is directly proportional to the concentration
(A a [ ]), that is, the more atomised metal present, the higher the absorbance. 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.
Remember a sports drink is not just sodium ions and water! It is important that we are aware of everything else in
the solution 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.
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Types of Interference
In this experiment we are primarily interested in determining the sodium content of sports drinks by AAS. However, potassium
is also present within sports drinks, and may affect the analysis of sodium (resulting in an interference).
In AAS, it is very important that the standards we prepare are as close as possible to the sample under
investigation. When using a calibration curve to calculate an unknown concentration, your results will only
ever be as good as your standards. Bad standards give bad results! In this investigation, the interference
due to other ions (potassium) is accounted for in by adding potassium as well as sodium to the standard
solutions.
The reason for adding potassium is simple. Alkali metals such as sodium have low ionisation potentials (lose
electrons easily) and readily ionise in the flame. This results in sodium ions being present as well as sodium
atoms, however, only sodium atoms absorb the unique radiation that is generated (via the lamp in the AAS)
and measured by the instrument. When this occurs, it is called ionization 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, keeping 100% of the sodium as atoms which
can be measured.
To show the effect from the presence of potassium, the sports drinks will be analysed using two sets of standard solutions.
One series will be made containing only sodium ions, while another series will be prepared with the same concentrations of
sodium ions and will also contain a large excess of potassium ions.
Principle
Your 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) Stock
10,000 mg/L Potassium (KCl) Stock
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Procedure
Preparation of Standards
1.
Prepare a series of four standard sodium solutions containing 0, 10.0, 20.0, 50.0 mg/L (ppm) of sodium. To do this,
measure accurately, 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 “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 5 mL aliquot (precisely known volume – use a pipette) of Gatorade into the provided 100 mL 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.
Analysis
7.
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.
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Results
Table 2: 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 3: Absorbance of Sports Drink.
Name of sports drink
Sodium concentration from bottle*
AAS Absorbance
Dilution Factor
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 Table 3 above would be expected to be closer to the “real” value?
Justify your choice.
2.
Draw a block diagram of an AAS unit and use this to briefly explain how the system operates. https://www.youtube.
com/watch?v=_KZjb9G3hB8
3.
What else could be analysed using AAS?
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DETERMINATION OF SULFATE IN LAWN FOOD
SAFETY GLASSES MUST BE WORN AT ALL
TIMES DURING THE LABORATORY SESSION
RISK ASSESSMENT
Hazard
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. Use in fume cupboard.
Pipette filling
Do not pipette by mouth. Use a pipette bulb to
fill pipette.
BaCI2
Harmful by inhalation. Toxic if swallowed.
Skin contact may produce health damage.
Avoid skin contact. Wash affected area well
with water.
Enriochrome Black T indicator
Cancer suspect agent. Cumulative effects may
result following exposure. Avoid skin contact.
Wash affected area well.
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 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
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Introduction
The traditional method used to determine sulfate concentration is using
gravimetric analysis. This is easily achieved by weighing the dried precipitate formed (barium sulfate) when barium ions are added to a sulfate
solution. Although this is a simple procedure, it is very 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.
Once filtered, the precipitate then needs to be oven dried to remove residual
water.
An alternative method, the one we will use, involves performing a back titration. A back titration is only used when the
endpoint of the forward reaction cannot be observed directly. This can occur when the reaction being measured is very slow,
produces side reactions, or when the analyte is an insoluble solid.
In this experiment, we cannot directly measure when the barium sulfate precipitate stops forming (due to the sulfate fertiliser
being completely consumed) because you cannot see through the already formed cloudy precipitate. Instead we use a colourometric indicator to measure the amount of excess barium that remains after the reaction is complete.
Barium ions react with sulfate ions in a 1:1 ratio as shown in the ionic equation representing the reaction below. If we know
how much barium is present before the reaction, and how much barium remains after the reaction, any difference between
these values must be how much has reacted with the sulfate. Because of the 1:1 ratio of the reaction, this difference will also
equal the amount of sulfate present.
The concentration of barium ions in solution can be easily measured by
completing a complexometric titration with ethylenediaminetetraacetic acid
(EDTA) using a metallochromic indicator.
The indicator Eriochrome black T (EBT) forms a red coloured complex
when bound to metal ions (e.g. Ba), and a blue complex when metal ions
are absent.
EDTA is a strong chelate - a species that reacts with metal ions to form a
http://en.wikipedia.org/wiki/Eriochrome_Black_T
stable metal complex. It reacts in a 1:1 ratio with barium ions by ‘stealing’
them away from the weaker eriochrome black complex. When enough EDTA
is added in the titration, all barium ions are no longer complexed with the
eriochrome black indicator which results in the solution turning blue,
signaling the end of the titration. This is illustrated in the reaction scheme
below:
Ba[EBT]
+
[EDTA] +
Ba[EDTA] + [EBT]
Red
Blue
We have chosen ammonium sulfate [(NH4)2 SO4] as the lawn food to analyse.
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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.
Titrations
This analysis requires each group to perform 2 titrations. The first of these is a ‘blank’ titration. It determines the amount
EDTA required to react with all the Ba2+ ions before the reaction with sulfate. The second titration is the ‘test’ titration. It
measures the amount of Ba2+ ions that remain in solution after the reaction with sulfate.
When completing the titrations, 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.
Interferences
This 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.
The presence of other metal ions (such as copper, zinc, etc) will lead an inaccurate estimation of sulfate. These ions will
also react with EDTA and if present, will lead to more EDTA being added than what is needed to react with the barium ions
in solution. The sample analysed today does not contain these interferences.
Supplied solutions:
Freshly prepared fertiliser solution[(NH4)2SO4] of known concentration (6.6 g/L)
Freshly prepared ~0.2 M barium chloride BaCl2 solution
Freshly prepared 0.1 M EDTA solution (Note: record the exact concentration – 4 decimal points).
Concentrated ammonia buffer solution NH3/NH4Cl (CAUTION: see safety note above).
Freshly prepared Eriochrome Black T indicator solution.
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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 buffer
Add 6 drops of Eriochrome black T 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 demonstrator if you are unsure.
6.
•
•
•
•
•
Transfer the clear filtrate to a 250mL conical flask. This will be used to carry out the test titration as follows:
Transfer all of the filtrate into a 250 mL conical flask
Add 10mL of concentrated ammonia/ammonium chloride buffer
Add 6 drops of Eriochrome black T 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 Vt.
7.
Repeat steps 6 for the second fertilizer sample.
8.
Use the information in the Results Table to calculate the percentage of sulfate ion in lawn fertilizer.
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Results:
Concentration of fertiliser (g/L)
6.6 g/L
Exact concentration of EDTA (4 d.p)
M
Volume of EDTA required for blank (Vb)
Titration 1 (2 d.p)
Titration 2 (2 d.p)
Average Titre: (2 d.p)
mL
mL
mL
Volume of EDTA required for test solution (Vt)
Titration 1 (2 d.p)
Titration 2 (2 d.p)
Average Titre: (2 d.p)
Difference in volume of EDTA (corresponding to Ba2+ used,
thus SO42- present) -Vb - Vt
moles of EDTA consumed
moles of sulfate
Initial mass of fertiliser in 50.0 mL sample
Mass of sulfate determined experimentally
mL
mL
mL
mL
moles
moles
g
g
Follow-up activities:
1.
Can you suggest any changes in the experimental procedure that may improve the accuracy?
2.
What was the reason for adding the ammonia/ammonium chloride buffer?
3.
When is it useful to use the back titration method of volumetric analysis?
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ANALYSIS OF NITRATE ION
CONCENTRATION IN WATER
SAFETY GLASSES MUST BE WORN AT ALL
TIMES DURING THE LABORATORY SESSION
RISK ASSESSMENT
Hazard
Substance, Apparatus, Procedure.
Precaution/ Action
Burette Filling
Use a plastic funnel to fill burette;
remove equipment from retort stand
and fill over sink.
Extract from HSC Chemistry Syllabus:
•
Perform first-hand investigations to use qualitative and
quantitative tests to analyse and compare the quality of water samples.
http://altura.speedera.net/ccimg.catalogcity.
com/210000/211600/211630/Products/5673991.jpg
•
Gather, process and present information on the range and
chemistry of the tests used to: monitor possible eutrophication of
waterways.
Introduction
Nitrate and phosphorous 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 (organic matter) being present.
Spectrophotometry
When 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 (400700nm), 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.
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Spectroscopy can also be carried out using other parts of the electromagnetic spectrum such as the ultra-violet (40010nm) and infrared regions (700-1050nm). 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 spectrum. 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. The Beer-Lambert Law quantitatively relates the extent of absorption of light to the concentration of
absorbing species in solution. 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:
It
= 10 - εcd
Io
where:
I0 It
ε
c
d
is the intensity of the incident light of definite wavelength,
is the intensity of the transmitted light of the same wavelength,
is the molar absorptivity (in units of L mol-1 cm-1) at that wavelength. It is a constant for a pure substance,
is the concentration of the absorbing substance (in mol dm-3, i.e. mol L-1),
is the pathlength (in cm) of the absorbing substance.
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Using mathematics (a log transformation) this equation may be simplified to a linear form:
A = εcd
where
A = absorbance
source
log10
I
o
I
t
polariser
diffraction grating
sample
detector
spectrum
signal
processing
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 mono-chromator.
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.
In this experiment you will use a spectrophotometer to determine the concentration of nitrate in an unknown water
sample. This will be carried out in the UV region (400-10nm) 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.
This experiment relies on the same analytical concept that was utilized in the AAS experiment. The basis of this analytical
method is that for a given analyte (nitrate) ion, the absorption is directly proportional to the concentration (A α [ ]), that is,
the more nitrate ions present, the higher the absorbance of radiation. In this experiment nitrate ions absorb wavelengths of
UV light.
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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, adding 1 mL HCl to each flask, and then making them up to the mark in a 100 mL volumetric
flask with ultra-pure water.
2.
Prepare you sample in a similar fashion by diluting 20 mL of the natural water sample in a 100 mL volumetric flask
with HCl (1 mL) 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.
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Results
Solution
Wavelength Maximum
(nm)
Absorbance at Maximum
0 mg/ L standard
5 mg/ L standard
10mg/ L standard
15mg/ L standard
Unknown
Unknown Concentration=
Follow-up activities:
1.
How would nitrates & phosphates get into the water supply? i.e where do they come from?
2.
What happens when too much nitrate & phosphate are both present in a water system?
3.
How would the levels of nitrate & phosphates be monitored by a governing body such as Hunter water?
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