The Outreach Program is supported by CSIRO Education

 

Analytical Instrument

Workshops

Student Notes

Name……………………………………………… School……………………………………………….. Instrument………………………………………… Demonstrator……………………………………..

CHEMISTRY OUTREACH

2

Table  of  Contents  

SAFETY  IN  THE  LABORATORY   3  

INTRODUCTION   4  

ANALYTICAL  INSTRUMENTS   7  

PART  A  -­‐  CHROMATOGRAPHY   7  INTRODUCTION   7  PRACTICAL  1:    HIGH  PERFORMANCE  LIQUID  CHROMATOGRAPHY   8  DETERMINATION  OF  CAFFEINE  IN  SOFT  DRINK   8  PRACTICAL  2:  GAS  LIQUID  CHROMATOGRAPHY   10  DETERMINATION  OF  ALCOHOL  IN  WINE  BY  GAS  CHROMATOGRAPHY   10  CHROMATOGRAPHY  DISCUSSION  QUESTIONS   12  

PART  B  -­‐  SPECTROSCOPY   14  INTRODUCTION   14  PRACTICAL  3:  UV-­‐VISIBLE  SPECTROSCOPY   16  SPECTROPHOTOMETRIC  ANALYSIS  OF  IRON   16  PRACTICAL  4:    ATOMIC  SPECTROSCOPY   21  ANALYSIS  BY  ATOMIC  EMISSION  &  ABSORPTION  OF  LIGHT   21  SPECTROSCOPY  DISCUSSION  QUESTIONS   25    

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SAFETY IN THE LABORATORY

The University of Melbourne has adopted the internally recognised systems Safety MAP and Environmental Management System (ISO14001) to ensure a safe and environment-friendly workplace for all staff, students and visitors. As a visitor to the University, you are responsible for adopting safe practices and you are required to comply with all relevant University and School of Chemistry rules and procedures. The Laboratory rules and Safe Work Procedures set out in this manual must be adhered to at all times and the direction of staff must be followed. If you have any concerns about the safety or environmental impact of any activity in this practical class, raise them with a staff member. Report all injuries, accidents or incidents immediately to your demonstrator.

LABORATORY RULES AND SAFETY PRACTICES 1. You will be issued with safety glasses with side shields conforming to Australian

Standard 1337, and these must be worn at all times. Prescription spectacles with polycarbonate lenses are acceptable provided they are fitted with side shields (available from optometrists). If you wear contact lenses, safety goggles (which provide a total seal around the eyes) must be worn instead of safety glasses.

2. You will also be issued with a laboratory coat, and this must be worn in the laboratory along with shoes that enclose the feet. Thongs, sandals and open style shoes are prohibited.

3. Long hair must be fastened securely. 4. Eating, drinking, smoking or chewing gum is not permitted in the laboratory. 5. Pipetting by mouth is prohibited. Safety pipettes are provided and must always be

used. 6. Listening to portable media players (iPod etc.) is not permitted in the laboratory.

In any chemical laboratory there is always a potential danger from accidental splashing or spillage of chemicals, cuts from broken glass, burns from touching hot apparatus or splashing hot liquids and fire. Part of your training in practical chemistry is to learn procedures to minimise these dangers and allow safe working conditions. All laboratory glassware mist be handled with due care. Hot objects should be allowed to cool before handling. If it is essential to handle a hot object, use a cloth or tongs to hold the object. In the case of a major chemical spillage or fire, evacuation of the laboratory may be needed.

EVACUATION PROCEDURE When the alarm sounds: Stop what you are doing and turn off electricity and any gas taps. Listen to evacuation instructions. Move quickly from the laboratory using the nearest exit, taking only personal belongings. In the passageway, the fire wardens will direct you to the building exit and to the assembly point well clear of the building. You may only re-enter the building when the chief fire warden gives the all clear.

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INTRODUCTION

Chemical analysis is very important. Almost all commercial products are subject to some form of analysis to ensure quality, and to identify and determine the amounts of constituents. Analysts test the primary constituents (are we getting what we have paid for?) and attempt to detect impurities that might be harmful or interfere with the uses of the product. Chemical analysis is also important in optimising processes to ensure efficient and cost effective operation. For example, in a packet of chips, the dyes, plastic and foil in the wrapper are analysed to ensure that they are safe and free form contaminants. The fertiliser, water and soil used to grow the potatoes are tested, as are the ingredients in the steel used to slice the potatoes, the oil used for cooking, and the salt and flavourings in the final product. The ingredients of most of the processed products that we use are similarly analysed. Australia is a world leader in the development and manufacture of instruments for chemical analysis. Two examples are

• the invention of atomic absorption spectroscopy by Sir Alan Walsh and

• the invention of the flame ionisation detector for chromatography by Dr. Ian McWilliam.

TYPES OF ANALYSIS Qualitative analysis identifies a compound in a sample. Quantitative analysis detects the amount of a compound in a sample.

REQUIREMENTS FOR AN ANALYSIS 1. A reason for analysis with clear aims and goals. 2. A representative sample taken in accordance with the aims and goals of the analysis. 3. An analytical procedure by which the analyte (the substance to be analysed for) is

determined. The procedure must be capable of the required precision. 4. The accurate recording of data and accurate calculation of a result. 5. A sensible estimate of the reliability or precision of the result. 6. The presentation of the result in a form that can be readily understood and upon

which action can be taken if required.

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INSTRUMENTS Modern analytical instruments take specific physical measurements of one or more properties of the sample being analysed. We exploit differences between properties of different compounds when we devise analytical methods. The values of these measured properties are transformed to an electrical signal that has an intensity directly related to the amount or concentration of the analyte (the compound being determined in the sample). The electrical signal is then displayed on a meter, chart recorder or computer. To convert this electrical signal to a concentration, we must calibrate our instrument.

CALIBRATING THE INSTRUMENT Instruments are calibrated using standards that contain different, but precisely known amounts of analyte - ranging from zero to an amount known to be greater than the amount present in the sample. The graph of the instrument’s reading plotted against amount of analyte in the sample is called the calibration curve. The “zero” standard is often called the blank. Analysts determine the amount of analyte in the sample by using calibration curves to compare the analytical response of their sample to the response of the standards. This is illustrated in the following diagram.

Depending upon the instrument and the analysis, the calibration curve may not be a straight line passing through zero but may actually be curved or give a non-zero response when the analyte is absent. This is not important. As we are never quite sure of the shape of the calibration curve it is necessary that the analytical response of the sample is within the calibration range of the standards, and that this response is reproducible. In the above example the unknown gave a reading of 3480. If we go up to our calibration curve, trace across from 3480 to our line and then down to the horizontal axis, we see the value of the concentration of the unknown was therefore 7.2

ANALYSIS

0

500

1000

1500

2000

2500

3000

3500

4000

4500

5000

0 1 2 3 4 5 6 7 8 9 10

Concentration/(units)

Inst

rum

ent

resp

onse

unknown

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QUALITY OF MEASUREMENT Three problems that can affect the quality of measurement are listed here. Understanding and overcoming some of these problems is important to endure good analysis.

Quality problem Reason Dealing with the problem Blunders and mistakes • mixing up samples

• not following instructions

• examine results with care

• draw graphs by hand rather than computer – people spot mistakes much better than computers.

Random errors • no analytical process can be reproduced exactly

• mixing, measuring and instrumental noise can vary

• repeat the measurement a number of times and calculate the mean value(s)

• produce a calibration curve (see below)

Systemic errors • standard solutions used in

two labs may be prepared from different materials

• one lab may be more competent than another

• all measuring instruments must be calibrated according to international standards

• all analyses must be able to be compared with results obtained using certified standard materials obtained from the US National Bureau of Standards

• all laboratories are certified by an independent agency. In Australia this agency is the National Australian Testing Association (NATA).

Producing a calibration curve is almost equivalent to repeating the measurements. We assume that the points will lie on a smooth curve. The average distance of the points from the calibration curve is used as a reasonable measure of the reliability or scatter of the data. This is easily estimated by inspecting the graph. In the analysis on the previous page the average distance of the points from the line is approximately 200, so the uncertainty in the measurement of the instrument response is 200. We can estimate that true value if the response for the sample is 3500 ± 200. The corresponding uncertainty in the concentration of the unknown is found by using the graph to determine the concentrations corresponding to the upper and lower uncertainty values for our measurement. The upper value (3500+200 = 3700) corresponds to a concentration of 7.6, and the lower value (3500-200 = 3300) gives a concentration of 6.8. Therefore, the concentration of unknown is 7.2 ± 0.4

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ANALYTICAL INSTRUMENTS

PART A - CHROMATOGRAPHY

INTRODUCTION Chromatographic analysis separates a mixture of substances into its individual components so that the identity or amount of each substance present can be determined. There are a number of different types of chromatography, including

• gas liquid chromatography (GLC)

• high performance liquid chromatography (HPLC)

• paper chromatography

• ion exchange chromatography. All chromatographs have at least two components – a mobile phase and a stationary phase. In paper chromatography, a sample is placed on a strip of paper (the stationary phase) that has one end sitting in a small amount of solvent (the mobile phase). The components of the sample separate because they are adsorbed and desorbed on to the paper differently as they are carried along it by the mobile phase. Components that strongly adsorb to the paper will take a longer time to travel than components that are weakly adsorbed onto the paper. Other chromatographic instruments are based on the same principles, but they are more complicated. Instruments differ in: (i) the way the samples are added (ii) the way that different substances in the sample are separated – in particular the type of stationary phase and mobile phase. (iii) the way in which the amount of each substance is measured.

DETECTION A device detects substances at the end of the column by measuring changes in the properties of the mobile phase. In gas liquid chromatography a common type of detector is the Flame Ionization Detector - a very simple device invented in Australia in the early 1960's. In this detector the substance is burned in a hydrogen flame. The flames of burning organic substances conduct electricity and the conductivity of the flame is directly proportional to the concentration of the substance in the flame. The changes are shown as peaks on a spectrum – the total amount of a substance in the sample is related to the area under the peaks. This detector measures the amount of substance present but not its identity. It will only work with combustible substances. Other detectors can be used for identifying the substances. These include mass spectroscopy and infra-red spectroscopy.

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PRACTICAL 1: HIGH PERFORMANCE LIQUID CHROMATOGRAPHY

DETERMINATION OF CAFFEINE IN SOFT DRINK

INTRODUCTION HPLC is one of the standard tools for quality control of complex materials. Their operating conditions – with components in solution and at room temperature – make these instruments suitable for the analysis of complicated biological molecules, drugs and chemical species that can decompose under the gas-phase, higher temperature conditions used for GLC. The instruments are expensive: $20,000 for a basic unit to $70,000 for a totally automated system. In HPLC the stationary phase in the column is more tightly packed than in other types of chromatography. The columns used in this experiment are packed with particles that are less than 0.005 mm in diameter. The total surface area of these tiny particles with which the components can interact is enormous - in a 15 x 1 cm column it is about half the size of a soccer field. Since there is very little space between the particles, it is necessary to apply a high pressure to the mobile phase in order to force it through the column at a reasonable rate. The pump maintains a precise flow rate so that the positions of the peaks in time can be used to identify the species in a sample. This is done by comparing the chromatographs of the test sample with those of prepared standards of the particular species to be determined. The common peak is an indication of the standard. A small sample (0.02 mL) is injected into the injector port where the mobile phase moves it through the column. Each component of the sample has different physical properties and structure, and moves at a different rate through the column. Thus, the components will be separated according to the size, shape and polarity of the molecules. Molecules that bind weakly to the particles in the column will move through the column quickest. As each set of molecules reaches the end of the column, a detector (most often an ultra-violet spectrometer) recognizes their presence and records a peak. Substances are identified by the length of time they are retained in the column – identical substances will have identical retention times under identical conditions. The area of the peak is proportional to the amount of that particular species in the injected sample. It is important that the volume of the sample of liquid injected is always the same. When this is so the area of the peak is proportional to the concentration of the component in the sample.

SAFETY • Always check with your demonstrator before turning on any electrical equipment.

• Use safety pipettes, never pipette by mouth. Methanol is poisonous.

• Be careful when using syringes – they have very sharp tips

• Always wear safety glasses

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DETERMINATION OF CAFFEINE CONTENT Procedure You are supplied with a standard aqueous solution of caffeine with a concentration of 1000 mg/L. Using the autopipette and 10 mL standard flasks prepare the following solutions

Flask Volume of added caffeine solution

Add distilled water to make a

Caffeine concentration

1 0.1 mL 10 mL solution 10 mg/L

2 0.3 mL 10 mL solution 30 mg/L

3 0.5 mL 10 mL solution 50 mg/L

4 0.7 mL 10 mL solution 70 mg/L

Your demonstrator will set up the HPLC so it is ready for use, and show you how to use it. 1. Clean the syringe by rinsing several times with the solution to be injected. 2. Follow your demonstrator’s instructions to inject the sample. Make sure that there is

no air in the syringe. 3. Make sure the injector lever is in the LOAD position, and fill the sample loop. 4. Move injector lever to INJECT position. Data will start collecting automatically. 5. Repeat steps 1 - 4 for each sample (standards and sodas). 6. Ask your demonstrator if the soft drink sample was diluted before it was injected. If

so, by how much? ________________ Run each of the solutions.

RESULTS Solution Retention time - tr Ccaffeine/(mg/L) Peak area

1

2

3

4

Sample

Analysis of data 1. Use the standard caffeine samples to identify the caffeine peak. What is the retention

time of the caffeine? ____________ 2. Use the retention time to determine if caffeine is present in the soft drink sample. 3. To quantitatively determine the amount of caffeine in the sample, measure the area of

the peaks due to caffeine in the standards, and construct a calibration curve by plotting caffeine concentration on the horizontal axis and peak area on the vertical axis.

4. Record the caffeine peak area in the soft drink sample chromatogram, and use the calibration curve to determine the concentration of caffeine in the soft drink sample.

DON’T FORGET TO ANSWER THE CHROMATOGRAPHY QUESTIONS ON PAGE 12 & 13

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PRACTICAL 2: GAS LIQUID CHROMATOGRAPHY

DETERMINATION OF ALCOHOL IN WINE BY GAS CHROMATOGRAPHY

INTRODUCTION Like any other chromatographic technique, the movement of a sample through a gas chromatography (GC) column depends on the way it moves between the mobile and stationary phase. The mobile phase in GC is a gas (usually nitrogen or helium) that is passed continually through the apparatus. The stationary phase is a viscous liquid. To begin the analysis, the sample is injected into the gas stream, and the sample injection port is often heated to vaporize liquid samples. The flowing gas stream then carries the sample into the column, where a portion of the sample will dissolve in the column liquid.

SAMPLE (gas) SAMPLE (dissolved in column liquid). So a section of the column now has column liquid containing dissolved sample in contact with gas containing no sample. As the solution moves down the length of the column, the sample is transferred from the liquid to the gas and back again, slowing the sample’s passage. The sequence of steps is illustrated in the diagram.

When a sample containing a number of different substances is injected, each substance travels through the column at a different speed and takes a different amount of time (the retention time, tr) to reach the end of the column. The retention time depends on the solubility of the sample in the column liquid. Gas chromatography is used in this experiment be used to determine the identity of alcohols in alcoholic beverages. In this analysis, an internal standard (propanol) is used to compensate for differences that may arise because of the nature of the sample and standards, and small variations in the volume of sample injected into the gas chromatograph. By adding the same amount of internal standard to the standards and sample, we can use it as a constant reference, and take the ratio of the instrument’s response due to analyte (ethanol) to the response of the internal standard (propanol), to correct for any handling errors. Safety • Always check with your demonstrator before turning on any electrical equipment.

• Be careful with the syringes. The tips are very sharp. Do not point them at yourself or anyone else.

• Be careful when handling the solutions of ethanol, propanol and the mixed alcohols.

• Always wear safety glasses.

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DETERMINATION OF ALCOHOL IN WINE Procedure Take 5 clean 25 mL volumetric flasks and use the autopipettes to make up the following solutions.

Solution Ethanol Propanol Water propanol of volumeethanol of volume Ethanol

Conc.

1 1 mL 5 mL to line 4%

2 2 mL 5 mL to line 8%

3 3 mL 5 mL to line 12%

4 4 mL 5 mL to line 16%

Wine Sample 5 mL of propanol and then wine to the line.

?

Run the sample and standards on the gas chromatograph and record the peak area for ethanol and propanol in the results table. Note that 1000 µL = 1 mL. Procedure • Fill a 1 µL syringe with 0.1 µL of ethanol

• Inject the sample into the gas chromatograph as per demonstrator’s instructions

• Wait for the peaks to appear

• Repeat using samples of propanol.

• What is the retention time for ethanol? _____________ Propanol? ____________

• Inject your standards and sample into the instrument and complete the results table below.

RESULTS Solution ethanol peak area propanol peak area

!

ethanol peak areapropanol peak area

1

2

3

4

Wine Sample

Analysis of data – calibration curve

• Plot

!

ethanol peak areapropanol peak area

on the vertical axis, and propanol of volumeethanol of volume on the horizontal

axis of a calibration curve

• Read off the value of propanol of volumeethanol of volume for your wine sample

• Calculate the % ethanol in the wine sample, using the equation:

% ethanol = propanol of volumeethanol of volume in the wine x

!

volume of propanoltotal volume

of wine x 100

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CHROMATOGRAPHY DISCUSSION QUESTIONS 1. Briefly describe the method you used in your analysis. A flow chart with the various

steps involved might be helpful. Did the method change from the description in these notes?

2. Briefly describe the principles of operation and the major components of the

instrument you used. A labelled diagram might be helpful. 3. What is the purpose of the mobile phase? Of the stationary phase?

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4. What is the purpose of the standards (caffeine for HPLC, ethanol and propanol for GC)?

5. Was the analytical procedure qualitative or quantitative, or both? Why? 6. What uncertainties (or errors) were involved in the procedure? 7. Were there any unexpected results? 8. Why does the syringe have to be carefully rinsed before each use? 9. How could you be certain a peak in the sample was your intended analyte (caffeine or

ethanol) and not another substance with a similar retention time?

CONCLUSION Concentration of caffeine in the cola drink = _____________mg/L or Concentration of ethanol in wine = _____________________mg/L Concentration stated by the manufacturer = ______________mg/L

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PART B - SPECTROSCOPY

INTRODUCTION Spectroscopic instruments monitor the way that light or some other form of electromagnetic radiation interacts with the atoms, molecules and subatomic particles in a sample. They can be used to identify chemical species, measure chemical concentrations and check chemical purity.

THE ELECTROMAGNETIC SPECTRUM The electromagnetic spectrum begins with the low energy, long wavelength radiowave region. Visible light is approximately in the middle of the spectrum. At the top end, we have gamma rays with very short wavelengths. This range of radiation is usually divided into a number of spectral regions according to the wavelength of the radiation, as listed below.

This image is reproduced courtesy of Andor Technology © Andor Technology plc. http://www.andor.com/learning/

The various regions of the electromagnetic spectrum interact with atoms and molecules in different ways. For example, when molecules absorb microwaves, they rotate faster, and absorption of infrared waves makes molecules vibrate more. These interactions can be used to qualitatively identify the different molecules. Chemists also use electromagnetic radiation for quantitative analysis, because the fraction of radiation that is absorbed by a sample is related to the concentration of the absorbing species in that sample.

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SPECTROSCOPIC INSTRUMENTS

As shown above, spectrometers contain a source of light or other electromagnetic radiation, a cell that contains the sample (the sample absorbs some of the light), and a monochromator, which selects radiation of the wavelength to be measured. A detector measures the intensity of the light that strikes it and produces an electrical signal. This signal is transformed so that it can be read from the display. Before analysing the sample it is necessary to take a “blank” reading by measuring the amount of radiation detected with no sample in place. The amount of light absorbed by the sample is determined by subtracting the sample reading from the blank reading. The proportion of light that is absorbed at a given wavelength depends on the concentration of the coloured species in the solution, and on the length of the cell through which the light passes. If we take our absorbance measurements at a fixed wavelength and use the same sample cell and instrument settings each time (as we will in this experiment), the absorbance is directly proportional to concentration.

DETERMINING CONCENTRATIONS BY SPECTROSCOPY. 1. Prepare standard solutions of known concentration and scan through the spectrum

for one of these solutions to determine to find the wavelength at which the sample absorbs the maximum amount of light. Analyses are carried out at this wavelength in order to maximize the accuracy of the determination.

2. Measure the absorbance of the standard solutions at the wavelength of maximum absorbance determined in step 1.

3. Construct a calibration curve (page 5), usually by plotting concentration on the horizontal axis and absorbance on the vertical axis.

4. Measure the absorbance of the unknown sample and determine the concentration from the calibration curve.

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PRACTICAL 3: UV-VISIBLE SPECTROSCOPY

SPECTROPHOTOMETRIC ANALYSIS OF IRON

INTRODUCTION Outer-shell electrons absorb ultra-violet and visible radiation as they move between energy levels. In molecules, the wavelength of radiation that is absorbed is related to the molecular structure, so the absorbance spectrum may be used to qualitatively identify compounds. UV/vis spectroscopy is much more commonly used for the quantitative analysis of compounds. When visible light of a suitable wavelength passes through a coloured solution, some of the light is absorbed. Your eye does not see the colour absorbed by a sample – it sees the colour transmitted, which is the colour that is complimentary to the absorbed colour. The table below details the colours observed when samples absorb different coloured light.

Wavelength /(nm) (1 nm = 10-9 m)

Colour absorbed Colour seen

400-435 violet yellow-green

435 -480 blue yellow

480-490 green-blue orange

490- 500 blue-green red

500- 560 green purple

560- 580 yellow-green violet

580-595 yellow blue

595-605 orange green-blue

605-750 red blue-green

Absorbance

Wavelength

Absorbance

Concentration

Fig. 1 Absorbance spectrum Fig. 2 Calibration curve If the solution that gave the absorbance spectrum in Figure 1 is absorbing blue light (with a wavelength about 430 nm), it will appear red. Red light (with wavelength about 700 nm) is not absorbed.

430 nm ↓

700 nm ↓

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Detection of Iron Iron is an essential human nutrient. To detect iron by UV/visible spectroscopy, analytical chemists often rely on the reaction of Fe3+ with thiocyanate ions (NCS-). This reaction gives an intensely red-coloured product that may be used as a qualitative test for the presence of Fe3+. The product of the reaction probably has the formula [Fe(OH)(NCS)(H2O)4]+.

[Fe(H2O)6]3+(aq) + H2O + NCS-

(aq) [Fe(OH)(NCS)(H2O)4]+(aq) + H3O+(aq)

The hydrolysis reaction of [Fe(H2O)6]3+ with water (below) can compete with the above reaction and interfere with the results. We add strong acid to the sample solution to drive the hydrolysis reaction backwards, ensuring all iron in the solution is present as the desired chemical species.

[Fe(H2O)6]3+(aq) + H2O [Fe(H2O)5OH]2+

(aq) + H3O+(aq)

In this experiment, we prepare a series of solutions with known concentrations of [Fe(OH)(NCS)(H2O)4]+ from a standard Fe3+

(aq) solution, and compare them to a solution prepared from an iron supplement. Most of the iron in the dietary supplement is in the Fe2+ form. Before the iron can be accurately complexed with the thiocyanate, all of the Fe2+ must first be oxidised to Fe3+ by reacting it with hydrogen peroxide (H2O2), which is a strong oxidising agent.

Safety • Always check with your demonstrator before turning on any electrical equipment.

• Never pipette by mouth. Use the dispensers supplied for 4 M HNO3 and 10% KNCS, and safety pipettes for other solutions.

• Always wear safety glasses.

PROCEDURE In today's practical you will: 1. Prepare a solution from your iron supplement sachet 2. Add hydrogen peroxide (H2O2) to oxidise any Fe2+ to Fe3+ 3. Add potassium thiocyanate (KCNS) to give an intensely red-coloured solution 4. Measure the absorbance of light by this solution. The amount of light that is absorbed

by the sample is proportional to the iron concentration 5. Measure the absorbance of light by solutions that have a precisely known

concentration of iron, and plot a calibration curve of absorbance versus concentration of iron in these solutions

6. Use this calibration curve to determine the iron concentration in the sachet. As the solution that you prepare is relatively unstable, you will need to make up the standards so that they are ready at the same time as the sachet sample. When they are all prepared, you should measure the absorbance of the standards as soon as possible at a wavelength of 480 nm using the spectrophotometer.

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Part 1. Prepare a calibration curve • The Fe3+ standard solution has a concentration of approximately 2 x 10-4 M. The exact

concentration is written on the bottle. What is it? ____________________

• Write “iron standard solution” on a beaker. Rinse the beaker with a small amount of the standard solution and then place 40 mL of the solution into the beaker.

• Label five 25 mL volumetric flasks with numbers 1 to 5

• Use an autopipette to transfer the amounts of iron standard solution listed in the table to these flasks.

• Add the 4M HNO3 and 10% KNCS solutions from their dispensers.

• Mix the solutions well by stoppering and shaking, and check the colour variation in the flasks by eye.

Flask Fe3+ solution 4 M HNO3 solution

10% KNCS solution

Water Fe3+ concentration

1 0 mL 2 mL 2 mL to the line

2 1 mL 2 mL 2 mL to the line

3 2 mL 2 mL 2 mL to the line

4 3 mL 2 mL 2 mL to the line

5 4 mL 2 mL 2 mL to the line

Using a clean Pasteur pipette each time, fill 5 cuvettes with your solutions and arrange them in order of concentration. Your demonstrator will show you how to handle the cuvettes.

Part 2. What frequency of light is absorbed? Use the spectrophotometer to record the spectrum of the solution in flask 4 over a range of approx. 400 – 700 nm. Include this spectrum with the report. Questions (Discuss the following questions as a group, but answer them individually.) 1. At which wavelength is the maximum amount of light absorbed? (Consult your

demonstrator if this value is not between 400 and 500 nm). 2. Give two reasons why 530 nm would be less suitable for use in the analysis.

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Part 3. Determining the iron content of the supplement To perform the analysis, we empty the sachet then convert all the iron to a form that strongly absorbs light, by oxidising any Fe2+ to Fe3+ and adding KNCS, which produces an intensely coloured solution. (i) Preparing the sample solution • Write down the concentration of iron in the sample stated by the manufacturer

_________________________. Our analysis will determine a much more precise value.

• Empty the contents of the sachet into a small beaker

• Use an autopipette to add 1.0 mL of this liquid to a 250 mL volumetric flask.

• Add 1ml of 10% hydrogen peroxide (H2O2) solution to your flask.

• Add 10 mL of 4M HNO3 and 10 mL of 10% KNCS solution

• Use a plastic pipette to make up to the mark with distilled water. This is the stock solution of your standard.

(ii) Analysing the stock solution Transfer an amount of this sample solution to a cuvette. Measure the absorbance of this test solution at the wavelength of maximum absorbance that you determined in Part 2. Record this absorbance in the table below. After you have plotted you calibration curve (Part 4) determine the concentration of Fe3+ in this solution, and thus calculate the concentration in your stock solution and then the sachet. Part 4. How much light is absorbed? Set the spectrometer to the wavelength where the maximum amount of light was absorbed in section 2. Measure the absorbance of each of your solutions and complete this table

Solution Fe3+ concentration Absorbance 1

2

3

4

5

Sample unknown

Use this information to plot a calibration curve with absorbance on the vertical axis and iron concentration on the horizontal axis. Plot the values of absorbance against concentration of iron in the unknown. Fit the graph to a straight line (check with your demonstrator if your line isn’t very straight).

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RESULTS What is the iron concentration in the sachet solution (in the 250 mL flask)? Assuming the sachet contains exactly 25.00 mL of iron supplement, what is the mass of iron in the sachet? We assume all the iron exists in the sachet as Fe3(PO4)2. Calculate the mass of Fe3(PO4)2 in the sachet. (Mr values Fe = 55.8, P=31.0, O=16.0) DON’T FORGET TO ANSWER THE SPECTROSCOPY DISCUSSION QUESTIONS ON PAGE 25 & 26.

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PRACTICAL 4: ATOMIC SPECTROSCOPY

ANALYSIS BY ATOMIC EMISSION & ABSORPTION OF LIGHT

INTRODUCTION Sir Alan Walsh of the CSIRO invented the technique of atomic absorption spectroscopy (AAS) in 1955. It is now widely used for the efficient analysis of many substances (animal, vegetable and mineral) for low concentrations of elements. Atomic absorption spectroscopy is usually only applied to analysis for metals because the wavelengths for most non-metals are outside an easily accessible region of the spectrum. You will investigate the techniques of atomic emission and atomic absorption spectroscopy in this experiment.

SUMMARY This experiment is in two parts. Part 1: Qualitative examination of atomic emission lines and bands Part 2: Quantitative analysis of sodium by atomic absorption spectrophotometry

PART 1: EMISSION BY LITHIUM, CALCIUM, COPPER, STRONTIUM, SODIUM AND POTASSIUM Introduction AES measures the light that is emitted by atoms in a flame. When a sample containing vaporized metal atoms is heated in a flame the atoms absorb energy, which excites the outer s electrons in the atom to become p electrons. The excited atoms then release their excess energy as light. The intensity of the light emitted is proportional to the concentration of atoms in the flame. In the same way as in AAS, the solvent carrying the cations evaporates and the ions are converted to atoms in the flame, for example: Na+

(g) + e- (in flame) → Na(g) The atoms of each element have characteristic colours and wavelengths at which emission of light occurs. The procedure works best for the alkali elements lithium (Li - deep red), sodium (Na - yellow), and potassium (K - lavender), which give sharp single colours (or 'resonance lines'). Unlike AAS, no separate light source is necessary for AES since the atoms in the flame emit the light being measured here. (i) Examining the Instrument The demonstrator will explain the components of the instrument. You should note:

• the gas inlet and burner (used to create a hot flame in which the sample is broken into atoms)

• the air inlet and aspirator vessel (used to carry sample solution into the flame)

• light source – hollow cathode lamp

• electronic readout system

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(ii) Emission colours A series of solutions will be supplied containing:

Analyte Concentration Na 500 mg/L as NaCl (1.27 g/l)

Li 500 mg/L as LiCl (3.05 g/l)

Ca 2000 mg/L as CaCO3 (5.00 g/l + 20 ml conc. HCl)

Sr 1000 mg/L as Sr(NO3)2 (2.42 g/l)

K 5000 mg/L as KCl (9.53 g/l)

Cu 5000 mg/L as CuSO4•5H2O (19.64 g/l)

Procedure 1. Using the same technique demonstrated previously, aspirate the solutions provided. 2. Record your observations in the table below

Element Group Period (in the Periodic Table)

Colour of Flame

Li+

Na+

K+

Ca2+

Sr2+

Cu2+

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PART 2: QUANTITATIVE ANALYSIS OF SODIUM BY ATOMIC ABSORPTION SPECTROPHOTOMETRY (AAS) Introduction AAS measures light of a specific frequency that is absorbed by atoms in a hot flame. These atoms are introduced into the flames as a solution of ions. The high temperature of the flame evaporates the solvent and converts the ions into atoms. The light is produced in an AAS lamp by bombarding atoms with electrons so that they are excited and vaporised. The atoms emit light as they relax. AAS instruments have the same element in the light and sample. In this analysis, we use a sodium lamp that emits light at 589.0 and 589.6 nm (yellow), the same frequency that is absorbed by atoms in the flame. The fraction of light that is absorbed is proportional to the concentration of the metal ion in the sample. So, as the atoms absorb light, the response of the detector decreases. The higher the concentration of the sodium, the more light absorbed by the flame. Procedure (i) Standard preparation 1. Label 6 250 mL volumetric flasks with the sodium concentrations listed in the table. 2. Use an autopipette to add the standard solution to these flasks, and make them up to

the mark with distilled water. The standard sodium solution contains 200 mg/L of Na+.

Solution Volume of standard sodium

solution

Sodium concentration

(mg/L) 1 0 0

2 1.0 1.0

3 2.0 2.0

4 3.0 3.0

5 4.0 4.0

6 5.0 5.0

(ii) Sample Preparation Soy Sauce 1. A stock solution of Soy Sauce will be provided to you that has 1 mL of soy sauce

diluted to 200 mL with water. Label this solution sample 1. 2. Using the autopipette take 1.0 mL of this solution and make up to the mark with

distilled water in a 100 mL volumetric flask. Label this solution sample 2.

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Gatorade 1. Using the autopipettes take 1 mL of the Gatorade sample and make up to the mark

with deionised water in a 100 mL standard flask. Label this solution sample Gatorade 1/100.

(iii) Measurements using AAS instrument 1. Aspirate the standards and samples and note the absorbance values in the table

below. 2. Use graph paper to plot a calibration curve with the absorbance (A) of your standards

on the vertical axis and concentration of sodium on the horizontal axis. 3. Draw the best straight line/smooth curve for the data points, including the blank (0

ppm). 4. Use the graph to read off the concentration of sodium in each of the "unknown"

samples.

Solution Sodium concentration (mg/L)

Absorbance

Standard 1 0 Standard 2 1.0 Standard 3 2.0 Standard 4 3.0 Standard 5 4.0 Standard 6 5.0 Soy sauce Gatorade

QUESTIONS 1. Why is it, in general, best to measure the samples in increasing order of

concentration? 2. Why is it a good idea to use the same aspirator for the set of measurements?

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SPECTROSCOPY DISCUSSION QUESTIONS 1. Briefly describe the method used in your analysis. A diagram illustrating the various

steps involved might be helpful. 2. Briefly describe the principles of operation and the major components of the

spectrometer you used. A labelled diagram might be helpful. 3. Is there anything else in the instrument (aside from the sample) that could absorb

light from your light source? Has this affected your results? Why/why not?

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4. Was the analytical procedure qualitative or quantitative, or both? Explain fully. 5. What is the analyte content (iron or sodium) as stated by the manufacturer? 6. What uncertainties (or errors) were involved in the procedure? 7. Were there any unexpected results? 8. What are some other applications or uses of the analytical technique that you

investigated?

CONCLUSION Mass of iron in the sachet = _______________________g or Concentration of sodium in soy sauce = _____________mg/L Concentration stated by the manufacturer ____________