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PART 3Monitoring and Management

2015

1

Prac Number

Prac name Syllabus Ref

16 Tests for cations perform first-hand investigations to carry out a range of tests, including flame tests, to identify the following ions:bariumcalciumleadcopperiron (Fe2+ and Fe3+)

17 Tests for anions perform first-hand investigations to carry out a range of tests, including flame tests, to identify the following ions:phosphatesulfatecarbonatechloride

18 Sulfate in lawn food: an example of gravimetric analysis

identify data, plan, select equipment and perform first-hand investigations to measure the sulfate content of lawn fertiliser and explain the chemistry involved

18a How much sulfate is in this fertiliser

identify data, plan, select equipment and perform first-hand investigations to measure the sulfate content of lawn fertiliser and explain the chemistry involved

19 Qualitative tests for assessing water quality

perform first-hand investigations to use qualitative and quantitative tests to analyse and compare the quality of water samples

20 Tests forheavy metal cations


21 Quantitative tests forassessing water quality


accuracy and validity

DEFINITIONS

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Experiment 16 TESTS FOR CATIONS

INTRODUCTION

Qualitative analysis is the process of identifying the substances present in a sample of matter. There are many possible tests, and in this experiment several tests including flame tests and precipitation reactions will be used.Some metal ions emit characteristic colours of light when heated in the blue flame of a Bunsen burner. Flame tests are therefore useful in identifying some cations.In precipitation reactions two ionic solutions are mixed with a solid precipitate being formed.The nature of chemical tests for identifying ions must be such that no two ions give the same results. By conducting appropriate tests the identity of the ion or ions in an unknown sample may be determined.

AIM

To carry out flame tests and a series of chemical reactions in order to devise tests for identifying the following cations in solution when these are the only ions that could be present. Ba2+ Ca2+ Pb2+ Cu2+ Fe2+ Fe3+

EQUIPMENT

Flame test PET bottlesSpray bottlesBunsen burnerSpotting test plates OR plastic sheets and coloured backing paper

CHEMICALS

Spray bottles containing solutions of Ca(NO3)2, Ba(NO3)2, Cu(NO3)2, Fe(NO3)2, Na NO3, KNO3 LiCl, SrCl2

Dropper bottles each containing one of the following solutions

Cation solutions Test solutions0.1mol/L Pb(NO3) 20.1mol/L Ba(NO3) 20.1mol/L Fe (NO3)30.1mol/L Cu(NO3)

1.0mol/L Ca(NO3)20.1mol/L FeSO4

1.0mol/L Na2SO4

1.0mol/L NaCl0.9mol/L NaOH1.0mol/L NH4OHAcidified 0.01mol/L KMnO4

0.1mol/L KIdistilled water

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RISK ASSESSMENT

What is the risk Why is it a risk How you could minimize the risk

METHOD Flame testsNOTE: Not all metal ions produce distinctive colours.

1 Use the spray bottle/PET set up with a Bunsen to test the flame colour of all the solutions in spray bottles.

2 Record colours of flames.3 Collect waste from bottom of PET bottle and rinse out the bottle.

RESULTS

COMPOUND IONS PRESENT IN COMPOUND FLAME COLOUR

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METHOD Precipitation reactions

USE EITHER SPOT TEST EQUIPMENT OR PLASTIC SHEET AND COLOURED BACKING PAPER. USE ONLY 2 DROPS OF SOLUTION FOR SPOT TESTS

1 Add 2 drops of each of the cation solutions onto plastic sheet or spotting plate.

2 Add 2 drops of the Cl– solution to each of the spots above. If no precipitate forms try

adding a few more drops.

3 Record the results in the table below. Write NP if no precipitate forms, and write ppt if

precipitate forms and give its colour.

4. Repeat steps 1–3 for SO42–.

5 Repeat steps 1–3 for NH4OH in micro test tubes; if precipitate forms add excess NH4OH.

RESULTS Precipitation reactions

TEST SOLUTION

CATIONSCI-

SO42– NH4OH EXCESS

NH4OH3

Pb2+

Cu2+

Ba2+

Ca2+

Fe2+

Fe3+

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ADDITIONAL TESTS

1 Add 2 drops of each of Fe2+ and Fe3+ solutions to separate spots on spotting plate. Add

2 drop of acidified MnO4– solution. Record the results.

2 Add 2 drops of Pb2+ solution to separate spots on spotting plate and add 2 drops of I–

solution. Record the results.

RESULTS

OBSERVATIONS

Fe2+ + MnO4–

Fe3+ + MnO4–

Pb2+ + I-

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QUESTIONS

1 Is the flame colour due to the metal ion or the anion present? Why?

2 Write an equation for each positive test

3 Give an additional identifying test for Pb2+.

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EXPERIMENT 17 TESTS FOR ANIONS

INTRODUCTION

In the previous experiment you conducted a series of tests to identify specified cations. In this experiment you will conduct a series of tests to identify specified anions.

AIM

To carry out a series of chemical reactions in order to devise chemical tests for identifying

the following anions in solution when these are the only ions that could be present. PO43-

SO42- CO32- Cl –

EQUIPMENTSemi Micro test tubes + rack OR OHP plastic OR spot test plates.

CHEMICALSDropper bottles each containing one of the following solutions.

Test solutions Anion solutions• 0.1mol/L Pb(NO3) 2 • 1mol/L Na2SO4• 0.1mol/L Ba(NO3) 2 • 1mol/L Na2CO3• 2mol/L HNO3 • 0.1mol/L Na3PO4• 0.05mol/L AgNO3 • 1mol/L NaCl• 0.9 mol/L NaOH

RISK ASSESSMENTWhat is the risk Why is it a risk How you could minimize the risk

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PROCEDURE

1 Add 10 drops of each of the anion solutions to each of 4 semi micro test tubes.

2 Add 10 drops of HNO3 to each of the semi micro test tubes and observe carefully.

3 Record the results in the table below. Write NP if no precipitate forms, and write ppt if precipitate forms and give its colour.

4 Repeat steps 1 –3 adding Ag+, then Pb2+ and then Ba2+.

RESULTS

TEST SOLUTIONS

ANION H+ Ag+ Pb2+ Ba2+

CO32–

CI–

SO42–

PO43–

QUESTIONS1. Write equations for each positive result.

2. Is lead ion useful in identifying anions? Explain why or why not.

3. When identifying anions in a solution where two anions may be present, it is necessary to destroy any CO32- before further testing for other anions can be carried out. How might this be done? (Consider the tests in this experiment.) Why is this necessary?

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TRY IT OUT

Your teacher will provide suitable solutions to the following investigations.

1 Use the results from the tests to identify the cation and anion (one each per

substance) in unknown substances.

2 Solids are amongst the unknowns. Your first step with them is to dissolve them in

acids. Use nitric acid.

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White precipitate

Cl-, PO43-, SO4

2-

SO42-

No precipitate

No bubbles produced

CO32-

Bubbles produced

Test 1: Add 2M HNO3

Test 2: Add AgNO3

PO43-

Unknown Anions

Cl-, PO43-, CO3

2-, SO42-

Precipitate forms

Test 3: Add Ba(NO3)2

Cl-

Testing Anions

White precipitateYellow precipitate

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Unknown Cations Ba2+, Ca2+, Pb2+, Cu2+, Fe2+, Fe3+

decolourisesDoes not decolourise

Test 4: Add a few drops of acidified potassium permanganate decolourises

Fe3+

Pb2+

Yellow precipitate of PbI2

Pb2+

Test 2:add KI.

Forms white precipitate

Fe2+, Fe3+

Blue precipitate forms which dissolves to from a royal blue solution when excess ammonia is added

Brown precipitate forms.

No precipitate forms

Test 2:Add SO42- ions, (Na2SO4 or H2SO4)

Ba2+, Ca2+, Cu2+, Fe2+, Fe3+

Pale green flame Dull red flame

Ba2+, Ca2+

Ca2+Ba2+

Test 3:Flame Test

forms white precipitate

Fe2+

Test 3:Add hydroxide ions. (NaOH or

Cu2+

No precipitate forms

Test 1:

Add chloride ions to sample.

Add 1 drop of the test reagent to 3 drops of the unknown solution either on the spotting plate or on the plastic sheet provided. For each test use a new sample unless otherwise stated.

If the sample is solid, add approximately 20 mg (1 rice grain) of the unknown to a test tube and dissolve in about 1 mL of distilled water or 2M nitric acid if it does not dissolve in water. Warm in a hot water bath if necessary.

Experiment 18 SULFATE IN LAWN FOOD: AN EXAMPLE OF GRAVIMETRIC ANALYSIS

INTRODUCTION

Sulfur in the form of sulfate is an important plant nutrient and is therefore found in many fertiliser preparations. As sulfates easily form precipitates with suitable metal cations, a gravimetric method based on this feature is a common way to analyse for sulfate. The usual metal ion used to form such a precipitate is barium, owing to the very low solubility of barium sulfate. The equation for this reaction is:

Ba2+(aq) + S042-(aq) BaSO4(s).

AIM: To determine the amount of sulfur (as sulfate) in a sample of lawn food.

EQUIPMENT

Electronic balance2 x 500mL beakers,stirring rod filter funnel and filter ring filter paper (high quality) sintered glass crucible (optional) , Buchner funnel and flask, mortar and pestle, hotplate,spatulathermometer

CHEMICALS

samples of lawn food, 1M hydrochloric acid, 0.5M barium chloride solution.Methylated spirits

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RISK ASSESSMENT

:What is the Risk Why is it a Risk? How will you minimise the risk?

METHOD

1.Grind up some lawn food in a mortar and pestle, and then accurately weigh out about 1 g.

2.Add this to about 100 mL of distilled water and dissolve as much as possible.

3.Filter the solution into a 400 mL (or larger) beaker. Wash the residue in the filter paper with several portions of distilled water.

4.Add about 25 mL of 1 M hydrochloric acid to the filtrate to dissolve any carbonates.

5.Make up to a total volume of 200mL with distilled water.

6.Heat to near boiling and then very slowly add about 20 mL of the barium chloride solution one drop at a time, stirring throughout.

7.Maintain the temperature just below boiling for 30 minutes, stirring occasionally.

8.Allow the precipitate to settle overnight.

9.DO NOT STIR. Observe carefully as you add a few drops of the barium chloride solution. If a precipitate forms, add a further 2 mL of the barium chloride solution.

10.Repeat this step until no further precipitate is observed to form. Try not to disturb the precipitate.

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11.Weigh a clean and dry Buchner funnel, together with 1 or 2 pieces of filter paper OR fit the Buchner with a sintered glass crucible.

12.Carefully filter the supernatant liquid and then the precipitate. Use the flat end of a spatula to scrape difficult-to-remove traces of precipitate into the funnel. Rinse the beaker a few times with water and add these washings to the filter paper.

13.Wash the precipitate with several portions of warm water followed by methylated spirits.

14.Dry the funnel ( or the sintered glass crucible) and contents in an oven at 110oC overnight.

15.Weigh when dry and calculate the mass of barium sulfate produced.

16.Use your results to calculate the percentage of sulfur in the lawn food tested.

RESULTS

Mass of lawn food=

Brand of lawn food=

Mass of barium sulfate produced=

ANALYSIS OF RESULTS

1. Calculate the mass of barium sulfate precipitated.

2. Calculate the moles of barium sulfate precipitated.

3. Write an ionic equation for the precipitation reaction.

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4. Calculate the number of moles of sulfate ions in the fertiliser sample.

5. Calculate the mass of sulfate present in the fertiliser sample.

6. Calculate the percentage by weight of sulfate in the fertiliser.

7. Compare your result against the manufacturers specification (if this is available). You may need to assume all the sulfate is derived from ammonium sulfate to determine the percentage by weight of ammonium sulfate in the fertiliser and then compare.

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QUESTIONS

1.In what parts of the method are errors likely to occur? Discuss the effect of these on your final result.

2.Why were the contents of the filter paper washed with water in step 3?

3. Why was the precipitate washed with water in step 13?

4. What is the "supernatant' solution?

5. Why is it tested with more barium chloride solution?

6. How reliable are the results of this investigation? Justify your answer.

7. What problems were encountered in the investigation? How might these be overcome to increase both reliability and validity of the test?

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Experiment 18 a HOW MUCH SULFATE IS IN THIS FERTILISER?

INTRODUCTION

Sulfur is essential for the growth of healthy plants. It is required to make some amino acids. Some soils can become deficient in sulfur and may need a fertiliser which adds sulfur (usually in the form of sulfate). The burning of fossil fuels produces oxides of sulfur which produce sulfuric acid when it rains. This leads to the phenomenon of "acid rain" which in some European countries has led to damage of forests and stone monuments. Efforts to reduce the pollution have led to such a reduction in acid rain, that some areas now have a sulfur deficiency!

In this investigation you will analyse a fertiliser called Epsom salts to see how much sulfate it contains.

You will do this by titrating a solution of Epsom salts with a solution of barium chloride of known concentration (the standard solution). The barium ions in the barium chloride react with the sulfate ions in the Epsom salts to form a precipitate of barium sulfate.

The indicator you will use for this titration is called alizarin red S (sodium alizarin sulfonate). The indicator remains yellow as long as any sulfate remains in solution. As soon as the sulfate has all reacted a pink complex of barium alizarin red S forms on the surface of the barium sulfate precipitate. Thus you can tell the end point of the titration has been reached because the mixture changes from pale yellow to pale pink.

In your practical notebook, write down the title of this investigation and, in your own words, write a sentence explaining its purpose.

EQUIPMENT

Burette Burette clamp 10mL Pipette Pipette filler Conical flasks x 2

Small Filter funnel Small beaker 250mL volumetric flask 100mL measuring cylinder Electronic balance

CHEMICALS

about 9g of some Epsom salts fertilizer distilled or deionised water,

methylated spirits

0.1mol/L hydrochloric acid in a dropper bottle

alizarin red indicator in a dropper bottle

about 100mL of the standard barium chloride solution (this will be about 0.05mol/L - the accurate concentration will be written on the bottle which contains the solution)

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RISK ASSESSMENT

What is the Risk Why is it a Risk? How will you minimise the risk?

METHOD

Make up solution of Epsom salts fertiliser

1. Accurately weigh out about 9g of the Epsom salts in a small, clean, dry beaker. Record this mass to two decimal places (This mass does not need to be exactly 9g, but whatever it is, you need to know it accurately.)

2. Completely transfer all the Epsom salt sample into the 250mL volumetric flask. Do this by washing the solid out of the beaker with distilled or deionised water. Use the funnel to help get the washings into the volumetric flask. Wash out the beaker several times into the flask to make sure all the Epsom salt sample ends up in the volumetric flask.

3. Make up the level of water in the volumetric flask to 250mL (up to the level marked on the stem of the flask). Use a dropper for the final tiny bits so you don't overshoot the mark. Turn the flask upside down (with your hand over the stopper!) and shake up the contents of the flask. Do this about ten times to make sure the solution is thoroughly mixed. Clearly label the flask so that you know what is in it (and who it belongs to).

Carry out the titration

1. Using a pipette and bulb pipette filler, transfer 10.0mL of the Epsom salts solution into a 250mL conical flask.

2. Add about 30mL of distilled water and 45mL of methylated spirits. Use a 100mL measuring cylinder for these measurements. Be careful with the methylated spirits. Avoid breathing it in and getting it on your skin.

3. Add two drops of alizarin red S indicator and then add the 0.1mol/L hydrochloric acid, drop by drop, until the indicator turns yellow.

4. Fill the burette with the barium chloride solution.

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5. Carry out a quick, rough titration to find out about how much barium chloride solution needs to be added. The end point has been reached when the mixture changes from a pale, pastel yellow to a pale pastel pink. Record your rough titration volume, but don't use it in your calculations later.

6. Pour the contents of the conical flask down the sink with plenty of water.

7. Prepare the Epsom salts solution, distilled water and methylated spirits mixture again. Add two drops of alizarin red S indicator and drops of 0.1mol/L hydrochloric acid until the indicator is yellow.

8. Fill up the burette again. Titrate the Epsom salts mixture with the barium chloride solution. Add the barium chloride solution until you are about two or three milliliters away from your rough titration volume. Add another three drops of the indicator solution. Continue your titration now very carefully - adding the barium chloride solution drop by drop until the mixture in the conical flask turns pale pink. Record your accurate titration volume.

9. Do the titration another two times to obtain three accurate volumes. Each time, after you have finished, pour the contents of the conical flask down the sink.

RESULTS

Amount of fertiliser used=

Draw up a table to record your titration values. Calculate the average titre.

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CALCULATIONS AND CONCLUSION

Your task is to calculate the % by mass of sulfate in the original Epsom salts.

The boxes below contain all the things you need to do, but they are not in the correct order. They will help in the process. First rearrange them in the correct order. Number them appropriately

Now carry out your calculation

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Calculate the average titration volume from your

results.

Calculate the number of moles of barium ions which were

added to the conical flask (you know the concentration and the volume of the solution added).

The barium ions and the sulfate ions react in a 1:1

ratio, so work out the number of moles of sulfate ions in the

10.0mL of Epsom salts solution which was originally

added.

Calculate the total number of moles of sulfate ions in the 250mL volumetric flask (this

contains all the sulfate ions in the original sample of Epsom salts which you weighed out).

Calculate the mass of sulfate ions in the original sample of Epsom salts. From this, calculate the percentage by mass of sulfate ions

in your original sample which you weighed out.

You could also calculate the % by mass of sulfur in

the original sample of Epsom

salts.

QUESTIONS

1. Compare your results with the other members of the class. What differences do you notice

2. What are the possible sources of error in this experiment?

3. How could you minimize the error so that you have the most reliable result.

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Experiment 19 QUALITATIVE TESTS FOR ASSESSING WATER QUALITY

INTRODUCTIONWater is an essential resource for all life forms. In plants and animals, the water contained in each cell is the solvent in which cellular processes take place. Animals and plants cannot survive without water. Water is such a powerful solvent that it is never found naturally pure. It dissolves gases from the air and minerals and organic matter from rocks and soil.Wastes from human activities are often discharged into oceans and rivers and dissolve in the water, often significantly altering its quality. Scientists continually monitor water quality to ensure it is safe for human consumption and use.

It is important to monitor water for the presence of heavy metal cations because many of them are toxic to humans and other organisms. Some of the cations such as zinc and copper are needed in trace amounts but higher concentrations are harmful to humans and the environment .Cations in solution are primarily identified by precipitation reactions. One cation can interfere with the test for another. Performing the tests in a particular sequence and on the filtrates left over from previous tests helps overcome this problem.

AIMTo analyse samples of water for appearance and smell, acidity and common

dissolved anions

EQUIPMENTSamples of water from different sources such as sea water, bore water, irrigation channel water or water from a stormwater drain.

test tubes test tube rack pH meter 100mL beaker

CHEMICALS0.05 mol/L, 0.01mol/L and

0.0001 mol/L AgNO30.1mol/L, 0.2 mol/L and 0.02mol/L BaCl2universal indicator1mol/L NH4OH

ammonium molybdate [(NH4)2MoO4] reagent7mol/L Nitric Acid0.1mol/L KI0.9M NaOHAcidified potassium permanganate

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RISK ASSESSMENT


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METHODObtain and test 3 water samples from different sites such as those mentioned above.

A Appearance and smellMake the following preliminary observations of the sample and record them in the

table below.1 Look carefully at the sample collected. Note its clarity—is it clear or cloudy? Note

its colour—for example it may be brown, grey, green or colourless.2 Smell the sample. Is there any odour of petrol or oil, rotten egg gas, ammonia,

chlorine or is it odourless?

RESULTS

PRELIMINARY

OBSERVATIONS

SAMPLE 1 SAMPLE 2 SAMPLE 3

Clarity

Colour

Smell

B Acidity

Use either a pH meter or universal indicator to determine the pH. If the water is

cloudy, filter it first.

RESULTS

Sample 1 Sample 2 Sample 3

pH

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C Tests for common anions, chloride (Cl -), sulfate (SO42-), phosphate (PO43-)

Perform the tests given in Table 1 Tests for common anions. You should be familiar with these tests from previous experiments.

Table 1: Tests for common anions

ANION TESTS

Chloride

(Cl-)Add drops of AgNO3 solution: if Cl- is present a white ppt forms or the sample becomes turbid (murky)To get an estimate of concentration: If 1 drop of 0.01 mol/L Ag+ added to a 10 mL sample produces turbidity, the concentration Cl– > 0.2 ppm.If 1 drop of 0.0001 mol/L Ag+ added to a 10 mL sample produces turbidity, the concentration Cl– > 20 ppm

Sulfate

(SO42-)

Add drops of BaCl2 solution: if SO42– is present a white ppt forms or the sample becomes turbid.To get an estimate of concentration:If 1 drop of 0.2 mol/L Ba2+ added to a 10 mL sample produces turbidity, the

concentration SO42– > 0.1 ppm.

If 1 drop of 0.02 mol/L Ba2+ added to a 10 mL sample produces turbidity, the

concentration SO42– > 1 ppm.

Phosphate

(PO43-)

Add a drop of 7mol/LHNO3 followed by 10 drops of ammonium molybdate reagent: if PO43– is present a yellow ppt or turbidity forms.Even a very pale turbidity indicates a phosphate concentration well above that of clean water. A negative test indicates phosphate at ‘acceptable’ levels — say < 0.1 ppm.

RESULTS

TESTOBSERVATIONS


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CONCLUSIONS:

Sample 1

Sample 2

Sample 3

D Tests For Heavy Metal Cations, lead, copper (ll), zinc, chromium, mercury, iron

METHODTest each of the samples using the table below

CATION TESTSlead (Pb2+) Add drops of KI solution: if Pb2+ is present, a yellow ppt forms

or sample becomes turbid.copper(ll) (Cu2+) Add drops of NaOH solution: if Cu2+ present, a pale blue ppt

forms or the sample turns turbid; addition of excess NH3 solution should cause the ppt to dissolve, leaving a blue solution.

zinc (Zn2+) Add drops of NaOH solution: if Zn2+ is present, a white ppt forms or the sample turns turbid; addition of excess NaOH or NH3 should dissolve the ppt.

chromium (Cr3+) Add drops of NaOH solution: if Cr3+ is present a grey-blue ppt forms or the sample goes turbid; addition of excess NaOH should dissolve the ppt, while addition of excess NH3 should have no effect on the ppt.

mercury (Hg2+) Add drops of NaOH solution: if Hg2+ is present, a yellow or red ppt forms or the sample turns turbid. Add drops of KI solution: if Hg2+ is present, a red ppt forms which dissolves in excess KI.

Iron (Fe2+) or (Fe3+) Add drops of NaOH solution: if Fe2+ or Fe3+ is present a brown ppt will form. Add a few drops of acidified KMnO4 to a fresh sample

If Fe2+ is present the acidified KMnO4 is decolourised

If Fe3+ is present the acidified KMnO4 is not decolourised

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RESULTS

TEST OBSERVATIONS


CONCLUSIONS:

Sample 1

Sample 2

Sample 3

QUESTIONS

Write equations for the precipitation reactions which occurred.

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Experiment 20 QUANTITATIVE TESTS FOR ASSESSING WATER QUALITY

INTRODUCTION

In the last experiment you conducted a number of mainly qualitative tests to assess aspects of water quality. In this worksheet you will use other tests to assess further aspects of water quality in a more quantitative manner.

AIMTo analyse samples of water for:A suspended solids and total dissolved solidsB turbidityC hardnessD dissolved oxygen (DO)E biochemical oxygen demand (BOD)

EQUIPMENTfilter funnelburettebeakerstirring rodBunsen burnerpipe clay triangle tripod oSecchi disc or turbidity tubeConductivity metervolumetric flasksfilter paperconical flasksevaporating basin25 mL pipetteelectronic balanceBurette ClampPipette filler500mL conical flask Oxygen sensor and data logger

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CHEMICALS

Samples of water from different sources such as tap-water, local sea water, bore

water, irrigation channel or stormwater drain.

5 mL buffer (57 mL conc ammonia solution and 7.0 g NH4Cl diluted to 100 mL)

10 mL 2.0 mol/L MnSO4.4H2O solution

10 mL sodium iodide solution (alkaline iodide solution) (50 g NaOH + 15 g NaI in 100

mL)

5 mL 9 mol/L H2SO4

starch indicator (0.5 g boiled starch in 100 mL)

100 mL 0.02 mol/L EDTA

0.025 mol/L sodium thiosulfate solution

2 mL Eriochrome Black-T indicator

RISK ASSESSMENT


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METHODObtain water samples from different sites such as those mentioned above. Different

groups in the class should analyse different samples and collate all results into the

table at the end of the experiment.

A Suspended and total dissolved solids (TDS)

1 Accurately weigh a clean dry filter paper.

2 Filter 100 mL of the water sample being sure to rinse all solids from the beaker.

3 Dry the filter paper in an incubator and weigh it again. Record the results. Clear

water should give a zero or possibly a negative result.

4 Carefully weigh a clean dry evaporating basin and pour the filtrate into it.

5 Place the evaporating basin over a gentle Bunsen flame and evaporate to

dryness.

6 Reweigh the evaporating basin.

7 Alternatively, if a conductivity meter is available for measuring TDS use it after

filtering off suspended solids (step 2). You do not need to evaporate the solution

to dryness.

RESULTS

Suspended solids

Initial weight of filter paper =__________

Final weight of filter paper =__________

Weight of solids = __________

% solid =

If percentage is very low

express it as ppm __________

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Dissolved solids

Initial weight of evaporating basin =____

Final weight of evaporating basin =_____

Weight of solids = __________

% solid = __________ _________

If percentage is very low express it as ppm

By conductivity meter TDS = ____________________

B Turbidity

Depending upon the water sample, the filtering in step 2 of A may not have produced

a clear solution. The suspended matter may have been so finely divided that it

passed through the filter paper or it may be successfully trapped by the filter paper. .

A measure of this suspended matter can be obtained by measuring turbidity.

Using a calibrated turbidity tube

A calibrated turbidity tube such as that supplied by

Streamwatch can be used to measure turbidity.

1 Add water to be tested slowly to the cylinder until

the mark at the bottom of the tube just becomes

invisible.

2 Read the turbidity value from the tube.

RESULTS

Using a turbidity tube

Turbidity = ____________

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ROLAND SMITH

C Hardness

Hardness is determined by measuring the combined number of moles of Ca2+ and

Mg2+ ions in the solution and expressing the result as an equivalent mass of CaCO3

per litre. (It is assumed that all the Ca2+ plus Mg2+ is Ca2+ equivalent.)

Mg2+ ions are necessary in this titration to give the purple colour and a distinct end

point. If there is no Mg2+ in the sample, some will need to be added and then this

amount subtracted from the final results. Adding 1 mL of 0.01 mol/L MgCl2.6H2O is

equivalent to adding 1 mg/L CaCO3 and should be enough to cause a good colour

change.

1 Add 25 mL of filtered water sample to a 250 mL conical flask.

2 Add 2–3 drops of Eriochrome Black-T indicator which changes from pink to blue

when chelated or acidified.

3 Add 1 mL of buffer.

4 Titrate with EDTA until a permanent blue colour appears (that is until the purple

colour loses its last trace of red).

5 Record the volume of EDTA used.

6 Repeat steps 1–5 twice more.

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RESULTS

TRIAL NUMBER VOLUME WATER USED VOLUME EDTA USED

1

2

3

AVERAGE

i Write a balanced equation for the reaction.

ii Calculate using the average values:

Concentration of EDTA solution = _____________

Moles EDTA used in the titrations = ____________

Moles Ca2+ + Mg2+ ions = __________________

Equivalent mass CaCO3 in sample = __________

Mass CaCO3 in mg/L (ppm) = ________________

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D Dissolved oxygen (DO)

NB. DO can be measured using the chemical tests below or it can be measured using the oxygen sensor and the data logger. Thus it can be used for the DO initially and then for the DO 5 and so we can calculate the BOD.

1 Obtain four 250 mL flasks completely filled with sample water and stoppered

underwater so all air has been excluded. Place two of the flasks, still firmly

stoppered, at 20°C in the dark for 5 days. (This will be used to measure BOD.)

2 Pour a very small amount of the water out and, using a graduated pipette with the

tip immersed in the water, add 2 mL MnSO4 solution and 2 mL sodium iodide

solution to each of the samples.

3 Stopper well, trying to ensure that no air is trapped in the flask, and invert the

flask several times to mix. A brown precipitate will form.

4 Wait until the precipitate has partially settled (until the top 2 or 3 cm of the

solution is clear), then withdraw 4 mL of clear solution and add 4 mL of 9 mol/L

sulfuric acid.

5 Re-stopper then invert the flask to mix; the precipitate should disappear.

6 Accurately measure 250 mL of the stoppered sample into a 500mL conical flask

and add 1.0 mL of starch suspension solution as an indicator.

7 Titrate with the sodium thiosulfate solution until the solution becomes colourless

(blue colour just disappears).

8 Record the volume of sodium thiosulfate used.

9 Repeat the titration process with one of other samples.

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RESULTS

TRIAL NUMBER VOLUME SAMPLE USED

VOLUME THIOSULFATE USED

1

2

AVERAGE

Equations for the reaction:

Mn2+ + 2OH → Mn(OH)2(s)

2Mn(OH)2(s) + O2(aq) → 2MnO(OH)2(s)

MnO(OH)2(s) + 2I– (aq) + 4H– (aq) → I2(aq) + Mn2+(aq) + 3H2O(l)

I2(aq) + 2S2O32– (aq) → 2I– + S4O62– (aq)

i Using the above equations write the net reaction for this test.

ii Give the stoichiometric ratio between O2 and S2O32–.

moles O2 present = moles S2O32–

iii Calculate using the average values:

Volume of sample used = ____________________

Moles of thiosulfate used = __________________

Moles O2 present = ________________________

Concentration of O2 mol/L = _________________

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Concentration of O2 mg/L = __________________

E Biochemical oxygen demand (BOD)

1 Obtain the stoppered flasks that have been incubated in the dark for 5 days.

2 Find the dissolved oxygen concentration using the above technique.

RESULTS

TRIAL NUMBER VOLUME SAMPLE USED

VOLUME THIOSULFATE USED

1

2

AVERAGE

i Using average values from the 5-day samples:

Volume of sample used = ____________________

Moles of thiosulfate used = __________________

Moles O2 present = ________________________

Concentration of O2 mol/L = _________________

Concentration of O2 mg/L = __________________

ii Calculate the BOD by subtracting the concentration of oxygen in the 5-day

sample from the concentration of oxygen in the initial sample:

BOD _______________mg/L

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QUESTIONS

1 Summarise your results for this and Worksheet 12 along with class results for different samples in the table below.

TEST RESULTS


Origin of the

sample

Appearance and

smell

Acidity

Dissolved anions

Suspended solids

and total dissolved

solids

Turbidity

Hardness

Dissolved oxygen

Oxygen demand

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2 Considering the results above, which sample would be the most suitable for human consumption? Justify your reasons. What additional test(s) would you recommend before actually drinking this water?

3 Suggest reasons for differences between samples. If you consider any samples to be polluted, suggest possible sources of the pollution.

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ReliabilityReliability is the extent to which an experiment, test, or any measuring procedure yields the same result on repeated trials.

Reliability: ExampleAn example of the importance of reliability is the use of measuring devices in Olympic track and field events. For the vast majority of people, ordinary measuring rulers and their degree of accuracy are reliable enough. However, for an Olympic event, such as the discus throw, the slightest variation in a measuring device -- whether it is a tape, clock, or other device -- could mean the difference between the gold and silver medals. Additionally, it could mean the difference between a new world record and outright failure to qualify for an event. Olympic measuring devices, then, must be reliable from one throw or race to another and from one competition to another. They must also be reliable when used in different parts of the world, as temperature, air pressure, humidity, interpretation, or other variables might affect their readings.

ValidityValidity refers to the degree to which a study accurately reflects or assesses the specific concept that the researcher is attempting to measure. While reliability is concerned with the accuracy of the actual measuring instrument or procedure, validity is concerned with the study's success at measuring what the researchers set out to measure.

Validity: ExampleMany recreational activities of high school students involve driving cars. A researcher, wanting to measure whether recreational activities have a negative effect on grade point average in high school students, might conduct a survey asking how many students drive to school and then attempt to find a correlation between these two factors. Because many students might use their cars for purposes other than or in addition to recreation (e.g., driving to work after school, driving to school rather than walking or taking a bus), this research study might prove invalid. Even if a strong correlation was found between driving and grade point average, driving to school in and of itself would seem to be an invalid measure of recreational activity.

A test is valid when it measures what it’s supposed to. How valid a test is depends on its purpose—for example, a ruler may be a valid measuring device for length, but isn’t very valid for measuring volume. If a test is reliable, it yields consistent results. A test can be both reliable and valid, one or the other, or neither. Reliability is a prerequisite for measurement validity.

reliable, but not valid

not reliable, not valid

reliable and valid

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ACCURACY, RELIABILITY AND VALIDITY The Board of Studies definitions are very brief and the following expanded definitions may be of use:

a) ACCURACY: Exactness or conformity to truth. Science texts refer to accuracy in two ways: (i) Accuracy of a result or experimental procedure can refer to the percentage difference between the experimental result and the accepted value. The stated uncertainty in an experimental result should always be greater than this percentage accuracy. (ii) Accuracy is also associated with the inherent uncertainty in a measurement. We can express the accuracy of a measurement explicitly by stating the estimated uncertainty or implicitly by the number of significant figures given. For example, we can measure a small distance with poor accuracy using a metre rule, or with much greater accuracy using a micrometer. Accurate measurements do not ensure an experiment is valid or reliable. For example consider an experiment for finding g in which the time for a piece of paper to fall once to the floor is measured very accurately. Clearly this experiment would not be valid or reliable (unless it was carried out in vacuum).

b) RELIABILITY: Trustworthy, dependable. In terms of first hand investigations the Board seems to define reliability as repeatability or consistency. If an experiment is repeated many times it will give identical results if it is reliable. In terms of second hand sources reliability refers to how trustworthy the source is. For example the NASA web site would be a more reliable source than a private web page. (This is not to say that all the data on the site is valid.) The reliability of a site can be assessed by comparing it to several other sites/sources.

c) VALIDITY: Derived correctly from premises already accepted, sound, supported by actual fact. A valid experiment is one that fairly tests the hypothesis. In a valid experiment all variables are kept constant apart from those being investigated, all systematic errors have been eliminated and random errors are reduced by taking the mean of multiple measurements. An experiment could produce reliable results but be invalid (for example Millikan consistently got the wrong value for the charge of the electron because he was working with the wrong coefficient of viscosity for air). An unreliable experiment must be inaccurate, and invalid as a valid scientific experiment would produce reliable results in multiple trials.

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