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National 5Chemistry

Chemical Analysis

Student Guide

Contents

Page 3 Investigation A1 – “Calcium analysis of water”

Page 6 Investigation A2 – “Calcium analysis of milk”

Page 12 Investigation B – “Iron in tea and cereals”

Page 15 Investigation C – “Chloride in seawater”

National 5 Chemistry

Practical Assignment Chemical Analysis: Student

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 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

An indicator called murexide is used which changes from pink to purple when the endpoint is reached.

In this experiment, the larger the titre of EDTA, the higher the concentration of calcium ions present in the water sample.



The ExperimentYou will need

0·01 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

Safety1 mol l-1 sodium hydroxide is corrosive. Wear goggles.

Method

1. 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.



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

Ca2+ + Na2C10H14N2O8 Ca C10H14N2O8 + 2Na+

calcium ions EDTA calcium compound sodium ions

An indicator called murexide is used which changes from pink to purple when the endpoint is reached.

In this experiment, the larger the titre of EDTA, the higher the concentration of calcium ions present in the milk sample.



The ExperimentYou will need0·1 mol l-1 EDTA 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 10 cm3 pipette and safety filler

100 cm3 conical flask 100 cm3 measuring cylinder

Distilled water White tile

Safety1 mol l-1 sodium hydroxide is corrosive. Wear goggles.

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 pink to 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 can sometimes be difficult to see. It is easiest to do a test run first so you can have a ‘target’ colour in front of you to compare.



Investigation B – “Analysis of Iron in foods”Background

Iron is an essential nutrient in our diets. It is needed for many things but especially because it forms the heart of the haemoglobin protein that allows our blood to carry oxygen around the body.

Many foods contain iron and breakfast cereals have iron added to them to increase their nutritional value – sometimes as fine iron filings. Tea leaves are also a good source of iron.

The quantity of iron present can be measured by titration.

Nitric acid is added to a sample of food, releasing iron(III) ions, Fe3+.

When potassium iodide is added iodine, I2, is formed.

2Fe3+ + 2Iˉ 2Fe2+ + I2

This iodine is titrated with a solution of thiosulfate, S2O32- .

I2 + 2S2O32- 2Iˉ + S4O62-

Starch solution is added as an indicator and changes from blue/back to colourless when the endpoint is reached.

The overall reaction in this experiment can be shown as:

iron ions thiosulfate

The larger the titre of thiosulfate, the greater the mass of iron present in the food sample.



The experiment

You 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

Safety

2 mol l-1 nitric acid is corrosive. Wear goggles.

Method

Preparing the solution

1. Accurately weigh about 2.0g of a dry food sample into a crucible and roast it in a fume cupboard for several minutes until all the food has turned to ash and no more smoke is coming off.

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 using the filter paper and funnel it.

5. 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.




Colour change of iodine/starch indicator



Initial colour iodine fading starch added final end-point

Investigation C – “Chloride in sea water”Background

For many of us, our earliest memory of the sea is getting a mouthful of seawater and finding out that it tastes horribly salty.

The concentration of salt in sea water can vary from place to place. Near the mouth of a river, where fresh water enters the sea the concentration of salt will be less. In very hot areas, where a lot of water is lost from the sea by evaporation, the salt concentration will be higher.

One of the most common ions found in seawater is the chloride ion, Clˉ. The concentration of chloride ions can be measured by titrating a sample of sea water using silver(I) nitrate solution.

AgNO3(aq) + Clˉ(aq) AgCl(s) + NO3ˉ(aq)

Potassium chromate is added to indicate the endpoint of the titration as it will form a red-brown precipitate when the endpoint is reached.

In this experiment, the larger the titre of silver(I) nitrate, the higher the concentration of chloride ions present in the water sample.



The Experiment

Equipment 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

Preparing dilute samples of sea water

* If the water 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.

Safety

0·1 mol l-1 silver nitrate is and irritant.1 mol l-1 potassium chromate is also an irritant and can cause health hazards with long-term exposure. Wear gloves and goggles. Both chemicals are harmful to the environment and should not be poured down the sink.

Titration

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

2. Pipette a 10·0 cm3 sample of diluted seawater into a conical flask and add about 50 cm3 distilled water and 1 cm3 of potassium chromate indicator.

3. Remove the funnel from the top of the burette and note the reading on the burette.4. Titrate the solution in the flask using the 0·1 mol l-1 silver nitrate solution from the

burette. 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.

5. The endpoint of the titration is identified as the first appearance of a red-brown colour of silver chromate. Read the burette to the nearest 0·1 cm3.






Colour change of potassium chromate indicator

Initial colour silver chloride near end-point silver chromate precipitate end-point

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