Chemical Analysis Student Guide National 5 Chemistry
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 Page 2
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, Ca 2+, 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, Ca 2+ 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. Ca 2+ + 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. Page 3
The Experiment You 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) 1 mol l -1 sodium hydroxide solution (NaOH) Clamp and stand Murexide indicator Funnel 3 cm 3 dropper or 5/10 cm 3 measuring cylinder 50 cm 3 burette 25 cm 3 pipette and safety filler 100 cm 3 conical flask Safety 1 mol l -1 sodium hydroxide is corrosive. Wear goggles. Method 1. Using the funnel, fill a 50 cm 3 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 cm 3 of your water sample into a 100 cm 3 conical flask. 3. Add 2 cm 3 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 cm 3. 7. Repeat the titration until your titres agree to within 0 2 cm 3. Page 4
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Investigation A2 - Calcium in milk Introduction Milk, 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. Ca 2+ + 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. Page 6
The Experiment You will need 0 1 mol l -1 EDTA solution Murexide indicator 1 mol l -1 sodium hydroxide solution (NaOH) Clamp and stand Funnel 3 cm 3 dropper or 5/10 cm 3 measuring cylinder 50 cm 3 burette 10 cm 3 pipette and safety filler 100 cm 3 conical flask 100 cm 3 measuring cylinder Distilled water White tile Safety 1 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 cm 3 of milk to the 100 cm 3 conical flask. 3. Using the measuring cylinder, add 40 cm 3 of distilled water to the flask. 4. Add 5 cm 3 of 1 mol l -1 sodium hydroxide using a 3 cm 3 Pasteur pipette or a small measuring cylinder. 5. Add a spatula tip of murexide indicator powder. 6. Remove the funnel from the top of the burette and note the reading on the burette. 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 cm 3. 8. Repeat the titration until the titres agree to within 0 2 cm 3. Page 7
* 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. Page 8
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, Fe 3+. When potassium iodide is added iodine, I2, is formed. 2Fe 3+ + 2Iˉ 2Fe 2+ + I2 This iodine is titrated with a solution of thiosulfate, S2O3 2-. I2 + 2S2O3 2-2Iˉ + S4O6 2- 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: 2Fe 2S O 2Fe S O 3+ 2-2+ 2-2 3 4 6 iron ions thiosulfate The larger the titre of thiosulfate, the greater the mass of iron present in the food sample. Page 9
The experiment You will need Preparing the solution Sample of food or tea Access to a balance (2dp) Bunsen burner, tripod and pipe-clay triangle 2 mol l -1 nitric acid solution crucible 100 cm 3 beaker 25 or 100 cm 3 measuring cylinder 50 cm 3 volumetric flask Funnel and filter paper The titration 20 cm 3 pipette and safety filler 100 cm 3 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 cm 3 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 cm 3 standard flask and made up to the mark using distilled water. Page 10
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 cm 3 of the food extract to a conical flask. 3. Add 1 0 g of potassium iodide. The solution should now go brown. 4. Remove the funnel from the top of the burette and note the reading on the burette. 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 cm 3 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 cm 3. 8. Repeat the titration until the titres agree to within 0 2 cm 3. Colour change of iodine/starch indicator Initial colour iodine fading starch added final end-point Page 11
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 redbrown 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. Page 12
The Experiment Equipment Needed Preparing dilute samples of seawater 20 cm 3 pipette and safety filler 100 cm 3 volumetric flask Titration diluted sea water sample 250 cm 3 conical flasks 10 cm 3 and 100 cm 3 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 cm 3 sample into a 100 cm 3 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 cm 3 sample of diluted seawater into a conical flask and add about 50 cm 3 distilled water and 1 cm 3 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 lemonyellow colour. Page 13
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 cm 3. 6. Repeat the titration until the titres agree to within 0 2 cm 3. Colour change of potassium chromate indicator Initial colour silver chloride near end-point silver chromate precipitate end-point Page 14