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National 5 Chemistry

Chemical Analysis

Teacher/Technician Guide

Contents

Teacher’s Guide

Page 3 Introduction and background

Page 4 Investigation A1 – “Calcium analysis of water”

Page 12 Investigation A2 – “Calcium analysis of milk”

Page 16 Investigation B – “Iron in tea and cereals”

1 Investigation C – “Chloride in seawater”

Technicians’ guide

Page 25 Investigation A – calcium investigations

Page 28 Investigation B – “Iron in tea and cereals”

Page 30 Investigation C – “Chloride in seawater”

Page 32 Risk assessments

Page 2


Practical Assignment Chemical Analysis: Teacher/Technician

Introduction and Background

The ability to know the precise composition of a substance is always going to be important: whether it be finding the percentage of metal in an ore to

see if it is suitable for mining or analyzing the level of pollutants in drinking water to ensure public health.

These days most of the analysis is highly automated using complex (and very expensive) equipment such as mass spectrometers or gas chromatographs which are out of the reach of most schools.

Traditional analytical techniques still have their place though, especially in the field as a way of getting initial values before taking samples back for further analysis. Additionally, the analyses are a useful showcase for some good chemical techniques.

These activities can be used to:

provide evidence for the National 5 Chemistry assignment

provide an opportunity for learners to become familiar with the use of titrations as an analytical technique. The data produced can be used to provide a context within which learners can practice calculations involving the mole and balanced equations.

A simple introduction is given for each of the experiments in the candidate guide. This sets the scene for each experiment and contains very limited background information.

Prior knowledge of redox and compleximetric titrations is not required at National 5 level. To make the chemistry accessible, each titration reaction is represented by a simplified balanced equation. This allows candidates to access interesting experimental chemistry at a level appropriate to National 5.

Before starting these activities it would be very useful if candidates had experience of doing titrations.



Where does it fit?

The Chemistry

Why is Chemical analysis topical?

Investigation A1 - “Calcium in water”

Background

Drinking water contains small amounts of salts and minerals dissolved from rocks that the water has passed through. Across Britain there is considerable variation in the concentration of different ions present in tap water.

Calcium ions, Ca2+, in drinking water can supplement the calcium in our diet and may be beneficial to our health. Some popular bottled waters are advertised as being high in dissolved minerals.

In high concentrations, Ca2+ ions can be a cause of “water hardness”. Hard water is not a health hazard but can form an unpleasant scum with soap as well as causing washing machines, irons and heating boilers to break down. The determination of water hardness is a useful test that provides a measure of quality of water for households and industrial uses. Originally, water hardness was defined as the measure of the capacity of the water to precipitate soap. Soap scum is formed when the calcium ion binds with the soap. This causes an insoluble compound that precipitates to form the scum you see. Soap actually softens hard water by removing the Ca2+ ions from the water.

The concentration of calcium ions can be measured by titrating a sample of water using a chemical known as EDTA.

Ca2+ + Na2C10H14N2O8 Ca C10H14N2O8 + 2Na+

calcium ion EDTA calcium compound sodium ions

The calcium ion concentration can be determined by titration with a chelating agent, ethylenediaminetetraacetic acid (EDTA), usually in the form of disodium salt. The titration reaction is:



The Ca2+(aq) ion is determined at a high pH, by adding NaOH solution to precipitate any Mg2+(aq) ions present in the water as Mg(OH)2(s).

Murexide indicator is used which changes from pink to purple when the endpoint is reached.

As the titre of EDTA is directly proportional to the concentration of calcium ions, candidates can compare the calcium ion levels in different samples without the need to carry out concentration calculations. Titre values can be used to rank the samples in order of increasing calcium ion concentration. This can be compared with the order found using literature/internet data.

Possible InvestigationsThere is a variety of different factors candidates can investigate. For instance,

the calcium content can be compared:

in different brands of mineral water

in water samples from around the UK*

in samples passed through different domestic water filters

*Details on how to prepare simulated water samples from different UK locations are provided in the technician guide.



Media Items1. A video providing an explanation of how calcium ions enter water supplies.

https://www.youtube.com/watch?v=ebygQes5Wig

2. Scottish water have a very useful breakdown of calcium content of water across the country.

http://www.scottishwater.co.uk/-/media/Domestic/Files/You-and-Your-Home/Water-Quality/ScottishWaterHardnessData2015.pdf?la=en

3. Map of England showing calcium carbonate levels

http://www.dwi.gov.uk/consumers/advice-leaflets/hardness_map.pdf

4. Information on EDTA titration of calcium ions

http://www.titrations.info/EDTA-titration-calcium

5. World Health Organisation document about calcium in drinking water

http://www.who.int/water_sanitation_health/dwq/chemicals/hardness.pdf

6. The average calcium content (along with other minerals) for the different bottled waters that they sell are provided on supermarket websites such as Tesco, Sainsburys etc.

7. How water filters work

http://www.explainthatstuff.com/howwaterfilterswork.html



http://www.explainthatstuff.com/howwaterfilterswork.html

http://www.who.int/water_sanitation_health/dwq/chemicals/hardness.pdf


http://www.dwi.gov.uk/consumers/advice-leaflets/hardness_map.pdf



https://www.youtube.com/watch?v=ebygQes5Wig

The ExperimentYou will need0·01 mol l-1 EDTA solution (if your water sample is very pure, you may need to use a 0.001 mol l-1 solution)

Murexide indicator

1 mol l-1 sodium hydroxide solution (NaOH)

Funnel

Clamp and stand 3 cm3 dropper or 5/10 cm3 measuring cylinder

50 cm3 burette 25 cm3 pipette and safety filler

100 cm3 conical flask

See technician’s guide for details of the reagents.* In the Pupil booklet it suggests 0.01 mol l-1 EDTA solution but for water samples very low in calcium, you may need to use 0·001 mol l-1 EDTA solution.** A solution will give greater consistency of colour intensity but it is easier to use the powder ground with sodium/potassium chloride added on a spatula tip – this gets around any issues of stability – though that is not much of a problem with murexide.



Method1. Using the funnel, fill a 50 cm3 burette with 0·01 mol l-1 EDTA solution, making

sure the tip is full and free of air bubbles.

2. Using a pipette, add 25·0 cm3 of your water sample into a 100 cm3 conical flask.

3. Add 2 cm3 of 1 mol l-1 sodium hydroxide to the flask using a dropper or a small measuring cylinder.

4. Add a spatula tip of murexide indicator powder

5. Remove the funnel from the top of the burette and note the reading on the burette.

6. Titrate the water sample using the 0·01 mol l-1 EDTA solution until the colour changes from pink to purple and then read the burette to the nearest 0·1 cm3.

7. Repeat the titration until your titres agree to within 0·2 cm3.



ExtensionTotal Hardness Determination.

As was mentioned in the introduction, permanent hardness of water is due to the presence of calcium and/or magnesium ions in the water; almost always both but in varying proportions.

The total hardness, therefore is a combination of the two concentrations. It may be of interest to compare two water samples of equal ‘hardness’ to see if they are actually the same, or indeed to see if two samples with the same calcium concentration have the same level of ‘hardness’.

To do this, we need to determine the amount of magnesium in the water. It is not possible (or at least not straightforward) to do this directly but it is fairly easy to determine the total hardness. Subtracting the calcium hardness from this will give the concentration of magnesium.

For the calcium determination, the pH of the solution is raised to pH 12 or above. This causes the magnesium salts to precipitate out as insoluble magnesium hydroxide.

The total hardness titration is carried out at a lower pH, about pH 10 produced by an ammonia buffer, and using a different indicator, Eriochrome black T (aka solochrome black)

Preparation

Ammonia Buffer

1. Dissolve 17·5g of ammonium chloride (NH4Cl) in 142 cm3 of concentrated ammonia (0·880).

2. Dilute to 250 cm3 with distilled water.

Eriochrome Black T preparation

The easiest method is for the Eriochrome Black T to be ground with potassium or sodium chloride as described for the murexide in the calcium titration. A spatula-tip of the powder can then be added.If a liquid indicator is desired, for instance to ensure a consistent colour intensity, it can be prepared as follows:

1. Put on gloves and protective eyewear and weigh out approximately 0·5 g of solid Eriochrome Black T, (EBT) on a balance and transfer it to a small beaker or flask. Add about 50 cm3 of 95 percent ethanol (IDA) and swirl the mixture until the EBT has fully dissolved.



2. Weigh out 4·5 g of hydroxylamine hydrochloride on a balance and transfer it to the beaker or flask containing the EBT. Swirl until the hydroxylamine hydrochloride has fully dissolved.

3. Transfer the solution containing the EBT and hydroxylamine hydrochloride to a 100 cm3 graduated cylinder. Add enough 95 percent ethanol (IDA) to bring the total volume to exactly 100 cm3

4. Transfer the EBT solution from the 100 cm3 graduated cylinder to a dropper bottle and label the bottle "0·5% Eriochrome Black T in ethanol."

Tips & Warnings

EBT indicator solutions typically exhibit very short shelf lives. Always prepare a fresh EBT solution when performing complexometric titrations.

Hydroxylamine hydrochloride is highly toxic and corrosive to skin and mucous membranes. Avoid direct skin contact. Wear rubber gloves and protective eyewear at all times when handling this compound.

Ethanol is flammable. Avoid working near open flames or other possible sources of ignition.

You will need0·01 mol l-1 EDTA solution (if your water sample is very pure, you may need to use a 0.001 mol l-1 solution)

Murexide indicator


Funnel



100 cm3 conical flask

Method

1. Fill a 50 cm3 burette with 0·01 mol l-1 EDTA solution, making sure the tip is full and free of air bubbles.

2. Add 25·0 cm3 of an unknown hard water solution into a 100 cm3 beaker.



3. Add 5 cm3 of ammonia buffer to the beaker.

4. Add 0·5 cm3 of Eriochrome Black T indicator.

5. Titrate with the 0·01 M EDTA until the colour changes from wine red to pure blue. Read burette to +/- 0·1 cm3.

6. Repeat the titration until the final volumes agree to +/- 0·2 cm3.



Investigation A2 - “Calcium in milk”IntroductionMilk, and other dairy produce are extremely important sources of calcium in the diet. It is very important for:

helping build strong bones and teeth

regulating muscle contractions, including heartbeat

making sure blood clots normally

A lack of calcium could lead to a condition called rickets in children and osteomalacia or osteoporosis in later life.

The same technique as for water analysis, EDTA titration, can be used to determine the concentration of calcium in milk, though using a higher concentration of EDTA to reflect the higher concentration of calcium.

As the titre of EDTA is directly proportional to the concentration of calcium ions, candidates can compare the calcium ion levels in different samples without the need to carry out concentration calculations. Titre values can be used to rank the samples in order of increasing calcium ion concentration. This can be compared with the order found using literature/internet data.

Possible investigationsThere is a variety of different factors candidates can investigate..

For instance,

the calcium content can be compared:

in milks from different sources (cow, goat, soya etc)

in treated milk (skimmed, homogenized, semi-skimmed, UHT etc)

in baby milks

in powdered milks

Media Items



1. A table of values for calcium in various milks and other foods

https://www.iofbonehealth.org/osteoporosis-musculoskeletal-disorders/osteoporosis/prevention/calcium/calcium-content-common-foods

2. A leaflet from the British Dietitians Association that gives information about milk and its dietary importance.

https://www.bda.uk.com/foodfacts/Calcium.pdf

3. The average calcium content (along with other minerals) for the different bottled waters that they sell are provided on supermarket websites, such as Tesco, Sainsburys etc.

4. Detailed data about the nutritional content can be found in tables from Public Health England, here.

https://www.gov.uk/government/uploads/system/uploads/attachment_data/file/416932/McCance___Widdowson_s_Composition_of_Foods_Integrated_Dataset.xlsx

(This gives you a large dataset in spreadsheet form – Open the tab (on the bottom) labelled inorganics. All sorts of mineral values are given, including calcium)

5. A BBC Good Food article on milk and nutrition

https://www.bbcgoodfood.com/howto/guide/which-milk-right-you

6. Information on EDTA titration of calcium ions


The ExperimentYou will need0·1 mol l-1 EDTA solution Murexide indicator


Funnel



100 cm3 conical flask 100 cm3 measuring cylinder

Distilled water White tile




https://www.bbcgoodfood.com/howto/guide/which-milk-right-you






PreparationA convenient way to use murexide indicator is by trituration.

A small amount of indicator, 0·1 g, is ground in a pestle and mortar with 20 g of potassium (or sodium) chloride until it is fully mixed. A spatula tip of the powder can then be added to the solution to titrate.

Method

1. Using a funnel, fill the burette with 0·1 mol l-1 EDTA solution, making sure the tip is full and free of air bubbles.

2. Using a pipette, add 10·0 cm3 of milk to the 100 cm3 conical flask.

3. Using the measuring cylinder, add 40 cm3 of distilled water to the flask.

4. Add 5 cm3 of 1 mol l-1 sodium hydroxide using a 3 cm3 Pasteur pipette or a small measuring cylinder.

5. Add a spatula tip of murexide indicator powder.


7. Titrate with the 0·1 mol l-1 EDTA until the colour changes from ‘salmon’ pink to ‘orchid’ purple*. Read the burette to the nearest 0·1 cm3.

8. Repeat the titration until the titres agree to within 0·2 cm3.

* The colour change is not as clear as it is for water samples but is still clear enough to see.



Investigation B – “Analysis of Iron in foods”

IntroductionIn this experiment the sample is dissolved in nitric acid which oxidises the iron to the ferric-state, Fe3+. Addition of excess iodide under mildly acidic conditions results in quantitative iron reduction to the ferrous-state, Fe2+, and simultaneous oxidation of the iodide to iodine.

2Fe3+ + 2Iˉ 2Fe2+ + I2

Iodine produced in the iron reduction is titrated with standard thiosulfate to a starch end-point.

I2 + 2S2O32- 2Iˉ + S4O62-

If students are simply comparing the levels of calcium in different samples, as long as the same sodium thiosulfate solution is used in each experiment its concentration does not need to be accurately known so it can be simply taken as made up.

If, however, you wish to use this experiment to determine actual concentrations of calcium ions, for example with Higher or AH students, as sodium thiosulphate is not a primary standard it will have to be standardised before use. This can be done by using your thiosulphate solution to titrate the iodine produced when an unmeasured excess of potassium iodide is added to a known volume of an acidified standard potassium iodate solution (iodate is a primary standard). The amount of iodine is known and thus the concentration of thiosulphate can be determined.

There are a few foods that will work using this method but we have only tested tea and breakfast cereal – there is no reason, however, why other readily ‘ashable’ foods could not be chosen too.

From a practical point of view, it might be preferable if the ‘ashed’ samples are prepared by technicians though there is no specific reason why pupils should not carry this out if it is wished.

As the titre of thiosulfate is directly proportional to the mass of iron present in each sample, candidates can use their titre values to rank the foods without having to calculate the mass of iron present. This can be compared with the order found using literature/internet data.



Possible investigationsThere is a variety of different factors you can investigate. For instance: Iron levels could be determined

in different types of tea

in teas from different countries

in breakfast cereals made from different crops (wheat, oat, corn or rice)

in organic, branded or own-label products

Media Items1. Information from the NHS about iron in the diet.

http://www.nhs.uk/Conditions/vitamins-minerals/Pages/Iron.aspx

2. Detailed data about the nutritional content can be found in tables from Public Health England, here.


(This gives you a large dataset in spreadsheet form – Open the tab (on the bottom) labelled inorganics. All sorts of mineral values are given, including iron)

3. Research paper with data on iron (and other mineral content) of various teas.

http://www.agriculturejournals.cz/publicFiles/50276.pdf

4. Information on iron intake and content in various foods

http://www.uhs.nhs.uk/Media/Controlleddocuments/Patientinformation/Digestionandurinaryhealth/Adviceforimprovingyourironintake-patientinformation.pdf

5. SSERC documents about iron and manganese in tea, including some sample data

http://www.sserc.org.uk/advanced-higher-revised/3069-iron-and-manganese-in-tea

6. Information on iodometric titrations

http://www.titrations.info/iodometric-titration



http://www.titrations.info/iodometric-titration






http://www.agriculturejournals.cz/publicFiles/50276.pdf




http://www.nhs.uk/Conditions/vitamins-minerals/Pages/Iron.aspx

7. RSC Classic Chemistry Experiments – Iron in breakfast cereal

http://www.rsc.org/learn-chemistry/resource/download/res00002108/cmp00000462/pdf

8. A Times of India article on the prevalence of iron filings in tea.

http://timesofindia.indiatimes.com/city/pune/Zero-iron-filings-in-tea-powder-is-not-possible/articleshow/18600259.cms

The experimentYou will need

Preparing the solution

Sample of food or tea 2 mol l-1 nitric acid solution

Access to a balance (2dp) crucible

Bunsen burner, tripod and pipe-clay triangle

100 cm3 beaker

25 or 100 cm3 measuring cylinder 50 cm3 volumetric flask

Funnel and filter paper

The titration

20 cm3 pipette and safety filler 100 cm3 flask

funnel 0·01 mol l-1 sodium thiosulfate solution

1% starch solution burette and stand

Dropper (for adding starch) white tile

MethodPreparing the solution

1. Accurately weigh about 2·0 g of tea/breakfast cereal into a crucible and roast it in a fume cupboard for several minutes until all the tea has turned to ash and no more smoke is coming off.

A significant amount of smoke is likely to be produced – It may be that the technician will prepare the extracts (or at least do the burning). If the pupils are doing it then there will need to be good ventilation or use of a fume cupboard.







2. Allow the ash to cool and wash it into a beaker using 2 mol l-1 nitric acid. [CORROSIVE]

3. Add a further 20 cm3 of 2 mol l-1 nitric acid [CORROSIVE] is added and boil the mixture for 5 minutes.

4. Let the mixture cool again and then filter it (to make sure any unburned carbon that could possibly remain in the mixture and affect the result is removed).

5. Place the filtrate is then placed in a 50 cm3 standard flask and made up to the mark using distilled water.

The titration

1. Using a funnel, fill the burette with 0·01 mol l-1 sodium thiosulfate solution, making sure the tip is full and free of air bubbles.

2. Using a pipette and safety filler, add 20·0 cm3 of the food extract to a conical flask.

3. Add 1·0 g of potassium iodide. The solution should now go brown.


5. Titrate the solution in the conical flask using the 0·01 mol l-1 sodium thiosulfate in the burette.

6. When the yellow colour has almost gone, add 1 cm3 of starch solution to produce a dark blue/black solution.

7. Continue titrating until the solution goes clear and colourless (and remains clear and colourless for at least 1 minute). Read the burette to the nearest 0·1 cm3.

8. Repeat the titration until the titres agree to within 0·2 cm3.



Investigation C – “Chloride in sea water”

IntroductionThis method determines the chloride ion concentration of a solution by titration with silver nitrate. As the silver nitrate solution is slowly added, a precipitate of silver chloride forms.

Ag+(aq) + Clˉ(aq) → AgCl(s)

The end point of the titration occurs when all the chloride ions are precipitated. Then additional silver ions react with the chromate ions of the indicator, potassium chromate, to form a red-brown precipitate of silver chromate.

2 Ag+(aq) + CrO42–(aq) → Ag2CrO4(s)

This method can be used to determine the chloride ion concentration of water samples from many sources.

As the titre of silver nitrate is directly proportional to the concentration of chloride ions, candidates can compare the chloride ion levels in different samples without the need to carry out concentration calculations. Titre values can be used to rank the samples in order of increasing chloride ion concentration. This can be compared with the order found using literature/internet data.

Possible investigationsThere is a variety of different factors you can investigate. For instance: The level of chloride ions could be determined:

In samples of water from different seas

In water sampled at different points in an estuary

Media Items1. A simple explanation of the oceans’ salinity

http://oceanservice.noaa.gov/facts/whysalty.html

2. Average composition of seawater and salinity of various seas.

http://dardel.info/IX/other_info/sea_water.html

3. A list of the salinity of various bodies of water.

https://en.wikipedia.org/wiki/List_of_bodies_of_water_by_salinity



https://en.wikipedia.org/wiki/List_of_bodies_of_water_by_salinity

http://dardel.info/IX/other_info/sea_water.html

http://oceanservice.noaa.gov/facts/whysalty.html

4. How salinity varies as you travel up an estuary.

https://books.google.co.uk/books?id=pXwSDAAAQBAJ&pg=PA8&lpg=PA8&dq=varying+salinity+in+forth+estuary&source=bl&ots=HSH0dWU7ok&sig=yZZZOYfcHV5VSEL2vEVfaaOXIYs&hl=en&sa=X&ved=0ahUKEwj91YLp8sbTAhVpBsAKHYZ0BU8Q6AEIXzAI#v=onepage&q=varying%20salinity%20in%20forth%20estuary&f=false

5. A guide to the Mohr method for determination of chlorides

http://www.titrations.info/precipitation-titration-argentometry-chlorides-Mohr

6. A World Health Organisation about chlorides in drinking water.

http://www.who.int/water_sanitation_health/dwq/chloride.pdf

The ExperimentEquipment Needed

Preparing dilute samples of seawater

20 cm3 pipette and safety filler 100 cm3 volumetric flask

Titration

diluted sea water sample 250 cm3 conical flasks

10 cm3 and 100 cm3 measuring cylinders 0·1 mol l-1 silver nitrate

1 mol l-1 potassium chromate indicator burette and stand

white tile funnel

Solutions NeededSilver nitrate solution: (0·1 mol l−1)

If possible, dry 5·0 g of AgNO3 for 2 hours at 100°C and allow to cool.

Accurately weigh about 4·25 g of solid AgNO3 and dissolve it in 250 cm3 of distilled water in a conical flask.

Store the solution in a brown bottle.

Potassium chromate indicator solution: (approximately 0·25 mol l-1)

Dissolve 1·0 g of K2CrO4 in 20 cm3 distilled water.

Sample Preparation



http://www.titrations.info/precipitation-titration-argentometry-chlorides-Mohr




If the seawater contains traces of solid matter such as sand or seaweed, it must be filtered before use.

Dilute seawater by pipetting a 20 cm3 sample into a 100 cm3 volumetric flask and making it up to the mark with distilled water.

Titration1. Pipette a 10·0 cm3 aliquot of diluted seawater into a conical flask and add

about 50 cm3 distilled water and 1 cm3 of chromate indicator

2. Titrate the sample with 0·1 mol l-1 silver nitrate solution. Although the silver chloride that forms is a white precipitate, the chromate indicator initially gives the cloudy solution a faint lemon-yellow colour. Before the addition of any silver nitrate the chromate indicator gives the clear solution a lemon-yellow colour.

3. The endpoint of the titration is identified as the first appearance of a red-brown colour of silver chromate

4. Repeat the titration with further aliquots of diluted seawater until concordant results (titres agreeing within 0·2 cm3) are obtained.



Additional Notes1. This titration should be carried out under conditions of pH 6·5 – 9·0. At higher pH

silver ions may be removed by precipitation with hydroxide ions, and at low pH chromate ions may be removed by an acid-base reaction to form hydrogen chromate ions or dichromate ions, affecting the accuracy of the end point.

If you are analysing samples of water as described then this will not be a problem.

2. It is a good idea to first carry out a “rough” titration in order to become familiar with the colour change at the end point.

3. The Mohr titration is sensitive to the presence of both chloride and bromide ions in solution and will not be too accurate when there is a significant concentration of bromide present as well as the chloride. However, in most cases, such as seawater, the bromide concentration will be negligible.



Technician GuideInvestigation A - CalciumEach group will need

EDTA solution* Murexide indicator**

1 mol l-1 NaOH 1x Burette

Clamp and stand 1 x 100 cm3 beaker for topping up burette with EDTA

100 or 250 cm3 flasks for titrations - 1 (to be washed out after each titration) or more

Small funnel for topping up burette.

Spatula for adding indicator 3 cm3 pasteur pipette (or 5/10 cm3 measuring cylinder) for adding NaOH

Samples of different milks Samples of different waters***

Preparation

* EDTA solutionIf possible, dry the disodium salt of EDTA for several hours or overnight at 80°C, allow to cool.

For calcium in water, this should be 0·01 mol l-1 BUT – if the water is very low in calcium then a lower concentration such as 0.001 mol l-1 will be needed

Weigh 1·86 g of the dried EDTA salt and dissolve it in 500 cm3 of distilled water in a volumetric flask.

(for waters that are very low in calcium, it may be necessary to dilute the EDTA further (1:10) to get a reasonable titre.

For calcium in milk, it should be 0·1 mol l-1

Weigh 4·65 g of the dried EDTA salt and dissolve it in 500 cm3 of distilled water in a volumetric flask.

**Murexide preparationThe easiest way to do this is a method called trituration. In a pestle and mortar add 0·1g of indicator powder to 20g or potassium or sodium chloride and grind thoroughly.



To use – add a spatula-tip of the salt/indicator powder to the solution.

*** Water preparationIn Scotland, most tap waters are low in calcium.

The easiest way to get round this is to purchase various mineral waters – they tell you the mineral content, including the calcium content, on the label. You can decant the water and suggest for instance that they are waters from different springs.

Alternatively, you can make artificial hard water

Add 0·7g of calcium sulphate-2-water to 1 litre of water in a bottle. Leave overnight to dissolve.

This gives you a solution that has 360 ppm of calcium in it – equivalent to very hard water areas like York and Lincoln.

To get water samples representative of other parts of the UK, dilute as follows:

Hard water eg Leicester 250ppm 69 cm3 made up to 100 cm3

Moderately hard eg Cheltenham 150 ppm 42 cm3 made up to 100 cm3

Slightly hard eg Blackpool 100 ppm 27 cm3 made up to 100 cm3

Or for Scotland

moderately sofy eg Moffat 24.5 6.8 cm3 made up to 100 cm3

moderately hard eg Shetland 52.1 14.6 cm3 made up to 100 cm3

Hard (eg Tiree) 110 30.5 cm3 made up to 100 cm3

(Note that Scottish Water uses ‘Hard’ and ‘Soft’ at slightly different levels.

Calcium sulphate produces what is known as permanent hardness.

If the experiment is looking at the effect of boiling water on calcium concentration, you will probably want to make up some temporary hard water.

Take 445 cm3 of freshly made limewater Bubble carbon dioxide through the solution so that the calcium carbonate

precipitates. Continue bubbling it until the solution goes clear again. Dilute the solution to 1 litre.

Assuming all the calcium has ended up as calcium hydrogen carbonate, this will give you a concentration of 360 ppm.

If you want, you can then make up dilute solutions as above.



Calcium hydrogencarbonate is not stable, it will slowly return to CO2 and calcium carbonate.



Investigation B – “Analysis of Iron in foods”

In this experiment the sample is dissolved in nitric acid which oxidises the iron to the ferric-state, Fe3+.

Sodium thiosulphate is not a primary standard so it will have to be standardised before use.

There are probably lots of foods that will work using this methods but we have only tested tea and breakfast cereal.

Each group will need

Access to a balance (2dp) crucible

Bunsen burner, tripod and pipe-clay triangle*

100 cm3 beaker

Funnel and filter paper 100 cm3 flask

50 cm3 volumetric flask Burette and stand

pipette

2 mol l-1 nitric acid** 0·01 mol l-1 sodium thiosulphate solution

1% starch solution

* A significant amount of smoke is likely to be produced – It may be that the technician will prepare the extracts (or at least do the burning). If the pupils are doing it then there will need to be good ventilation or use of a fume cupboard.

Preparing the solution

1. Accurately weigh about 2·0 g of tea/breakfast cereal into a crucible and roast it in a fume cupboard for several minutes until all the tea has turned to ash and no more smoke is coming off.

2. Allow the ash to cool and wash it into a 100 cm3 beaker using 2 mol l-1 nitric acid. [CORROSIVE]

3. Add a further 20 cm3 of 2 mol l-1 nitric acid [CORROSIVE] is added and boil the mixture for 5 minutes.

4. Let the mixture cool again and then filter it (to make sure any unburned carbon, that could possibly remain in the mixture and affect the result, is removed).



5. Place the filtrate in a 50 cm3 standard flask and make up to the mark using distilled water.

** 2 mol l-1 nitric acid is corrosive. Goggles to BS EN166 3 will be needed.



Investigation C – “Chloride in sea water”

This method determines the chloride ion concentration of a solution by titration with silver nitrate. As the silver nitrate solution is slowly added, a precipitate of silver chloride forms.

The end point of the titration occurs when all the chloride ions are precipitated. Then additional silver ions react with the chromate ions to form a red-brown precipitate of silver chromate.

Each group will need

burette and stand 10 and 20 cm3 pipettes/measuring cylinders.

100 cm3 volumetric flask 250 cm3 conical flask(s). If they are in short supply, pupils can wash theirs out between titrations.

10 cm3 and 100 cm3 measuring cylinders

0·1 mol l-1 silver nitrate

1 cm3 pasteur pipette 20 cm3 pipette and filler*

0·25 mol l-1 potassium chromate indicator

* If this is not easily accessible, the fact that the density of seawater is so close to that of distilled water, 1·025 compared to 1·000, means the aliquot can be measured by mass. 20 cm3 of seawater has a mass of 20·5g

Preparation

Silver nitrate solution: (0.1 M)

If possible, dry 5·0 g of AgNO3 for 2 hours at 100°C and allow to cool.

Accurately weigh 4·25 g of solid AgNO3 and dissolve it in 250 cm3 of distilled water in a conical flask.

Store the solution in a brown bottle.

Potassium chromate indicator solution: (approximately 0·25 mol l-1 )

Dissolve 1·0 g of K2CrO4 in 20 cm3 distilled water.



Water

If the seawater contains traces of solid matter such as sand or seaweed, it must be filtered before use.

Seawater can be prepared artificially by

EITHER

Purchasing marine salts from an aquatic centre

OR

Making up your own

Just make up solutions of sodium chloride

Dead sea – a 29% solution

Red sea – a 4·1% solution

North sea – a 3·4% solution

Black sea – a 2% solution

Baltic sea – a 0·8% solution

Estuaries, if you are unable to get samples from an actual estuary, you can make up representative samples for the different zones:

Mouth 3·4%

Lower estuary 2·7%

Middle estuary 2·1%

Inner estuary 1·2%

Upper estuary 0·25%

Dilute the seawater by pipetting a 20 cm3 sample into a 100 cm3 volumetric flask and making it up to the mark with distilled water.

Alternative microscale titrationPrepare the solutions as above

As well as those you will need equipment for a microscale titration – see the SSERC website for details



Activity assessed Testing water for calcium/magnesiumDate of assessment 26th July 2013Date of review (Step 5)SchoolDepartment

Step 1 Step 2 Step 3 Step 4List Significant hazards here:

Who might be harmed and how?

What are you already doing? What further action is needed?

Action by whom?

Action by when?

Done

EDTA is a skin, eye and respiratory irritant

Technician preparing solutions.

Wear gloves and eye protection. Avoid raising dust.

Sodium hydroxide is corrosive1M sodium hydroxide solution is corrosive

Technician preparing solutionsTechnician, teacher or pupils by splashes

Wear gloves and goggles (BS EN166 3).Wear goggles (BS EN166 3).

Ammonia .880 is corrosive and the fumes are toxic (Cat 3)

The ammonia buffer is corrosive and gives off toxic fumes (Cat 3)

Technician preparing buffer solution.

Technician, teacher or pupils by splashes or inhaling fumes

Wear gloves and goggles (BS EN166 3). Handle in a fume cupboard

Wear goggles (BS EN166 3). Work in a well-ventilated areas and keep lid off bottle for as short a time as possible.

SSERC Risk Assessment (revised version November 2009)(based on HSE ‘5 steps to risk assessment’)

2 Pitreavie Court, South Pitreavie Business Park, Dunfermline KY11 8UBtel : 01383 626070 fax : 01383 842793e-mail : [email protected] web : www.sserc.org.uk

http://www.sserc.org.uk/

mailto:[email protected]

Step 1 Step 2 Step 3 Step 4Murexide indicator (ammonium purpurate) has no significant hazardEriochrome black T is an eye irritant

Ethanol is flammable

Hydroxylamine hydrochloride is harmful by ingestions/skin contact, a skin/eye irritant, a skin sensitiser a category 2 carcinogen and can damage organs on repeated exposure.

Eriochrome Black T indicator solution is a skin sensitiser and a category 2 carcinogen.

Technician preparing solution.



Technician, teacher or pupils by splashes

Wear eye protection. Avoid raising dust.

Keep away from sources of ignition. Wear gloves and eye protection.

Wear gloves and goggles (BS EN166 3).

Wear gloves and goggles (BS EN166 3).

The reaction mixture is of no significant hazard.

Description of activity:

Water samples are titrated against EDTA solution. Using murexide and eriochrome black T indicators. The solution is made alkaline by pH 10 ammonia buffer for the total hardness or sodium hydroxide for the magnesium.

Additional comments:

Activity assessed Analysis of Iron in tea/cerealDate of assessment 28th April 2017Date of review (Step 5)SchoolDepartment




Action by whom?

Action by when?

Done

Burning Tea/cereal produces irritating smoke

Anyone nearby by inhalation of the smoke.

If more than a very small amount, carry out in a fume cupboard.

Sulphuric acid is extremely corrosive

Technician making up dilute solution

Wear gloves and face shield (or chemical resistant goggles EN 166 3 if the quantity is not large). Always add acid to water.

1M sulphuric acid is corrosive

Pupil/teacher by splashes during experiment

Wear gloves and chemical resistant goggles EN 166 3

Nitric acid is highly corrosive and oxidizing


Wear gloves and face shield (or chemical resistant goggles EN 166 3 if the quantity is not large). Keep away from





Step 1 Step 2 Step 3 Step 4flammables and reducing agents.

2M Nitric acid is corrosive

Pupil/teacher by splashes during experiment

Wear gloves and chemical resistant goggles EN 166 3

potassium manganate VII is a powerful oxidiser (and harmful if swallowed)


Keep away from flammables and reducing agents. Avoid raising dust.

0.01M potassium manganate VII has no significant hazard.Potassium iodide is an eye irritant

Pupil (or technician) weighing out solid

Wear eye protection. Avoid raising dust.

Iodine – the concentration of iodine in the solution is low enough to be of no significant hazardSodium thiosulphate is of no significant hazard.

Description of activity:Tea/cereals (or other foods) are burned and the ash boiled with 2M nitric acid to convert all the Iron to Iron III. The solution, diluted with water has potassium iodide added which reacts with Iron III to produce iodine. This is titrated with sodium thiosulphate using a starch indicator near the end point.


Activity assessed Mohr titration of chloride (Silver nitrate)Date of assessment 28th April 2017Date of review (Step 5)SchoolDepartment




Action by whom?

Action by when?

Done

Silver nitrate is an oxidising agent and is corrosive to skin and eyes.

Technician by splashes while preparing solutions

Avoid raising dust. Keep away from flammables and reducing agents. Wear gloves and goggles EN 166 3.

Potassium chromate is a mutagen and carcinogen. It is also a skin/eye and respiratory irritant and a skin sensitiser.The 1M solution has the same properties.

Technician while making up solution and pupils/teacher by splashes when using.

Avoid raising dust. Wear gloves and goggles EN 166 3.





Seawater is of low hazard but if genuine seawater is used it is best to boil the sample before use to destroy any potentially harmful micro-organisms.The reaction mixture is still classed as mutagenic and carcinogenic due to the chromate.

Description of activity:

Samples of seawater (real or artificial) are titrated against silver nitrate using potassium chromate as an indicator.


The chromate is very hazardous to the environment. To dispose, filter the reaction mixture and keep the residue (a mixture of silver chloride and silver chromate) for disposal by registered contractor. If the filtrate is yellow, meaning there is unreacted chromate, acidify to approximately pH 2 and add sodium hydrogensulphite to reduce to Cr(III). Precipitate the Cr3+ as hydroxide, filter and keep for disposal by a licensed contractor.

extension - scottish schools education research … · web viewhard water is not a health hazard...

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