Edexcel Chemistry Unit 3 Notes 3o4n40

Titrations 1. No unnecessary titrations. If you have done two that are close enough, that's already sufficient. close enough means two titres within o.10 cm^3. Example 21.50 and 21.60 is fine, 22.75 and 22.80 is fine but 23.25 and 23.40 is not enough. In such cases, perform another titration. BURETTE: A burette has an accuracy of 0.05 cm^3. The graduations are after every 0.10 cm^3 but you can read up to an accuracy of 0.05 cm^3. So the maximum accepted error in the titre is 0.10 cm^3 and that's how close the two readings we pick should be. Graphs 1. The scale needs your values to cover at least half of the page, mind it. 2. Sharp your pencil, use 2H pencil(not HB) for graphs if available. You can get in a good stationary shop. 3. Double check before plotting, but be real quick. One mistake in plotting and you may mess with the whole graph. 4. Line of best fit doesn't strictly mean line only, sometimes(rarely) it could be a curve too. 5. Always draw a triangle for gradient. The bigger the traingle, the more accurate is your gradient. 6. Ask me now if you don't know ideas about gradient and y-intercept clearly. 7. Always write by the axis, what it represents along with its unit. 8. Independent on x-axis, dependent is y-axis. Independent is something you are changing, depending is something you are measuring. 9. LABELLING!!! 10. Not more than 15 minutes on graph!!! 11. Even if graph is wrong, you won't lose *dherai* marks, just don't lose hopes and work out on gradient and intercept stuffs, you will get method markings. But never leave the questions, just fill the space even if you go blank. 12. When they ask you to prove some relationship like T=kV or T is directly proportional to V, write that since the graph is a straight line, the value of the K is constant and therefore the relationship is true. DON'T FORGET IT!!! 13. If you don't have enough time, don't make all points. Make two or three

points and work out on other stuffs... Believe me, it helps if you are running out of time.

Salt analysis (Ion test) 1. There is a trick with aluminium and lead(II) ions. Both of them give same observation with ammonia and sodium hydroxide. To distinguish them, you need to add HCl . If precipitate forms, it is because of lead(II) ions, otherwise it is aluminium ions. Lead chloride is insoluble. 2. Never spend more than 45 minutes on this section.

Gas test 1. You should write the OBSERVATION of the confirmatory test whenever needed. But to save time, you DO NOT need to perform it. Just use your common sense of chemistry. Example, when ZnCO3 is reacted with HCl, you know that a gas is released. What could that be? Just think...it is CO2...Now what you need to write is something like this(already given in the booklet) -"Colorless gas which turns lime water milky" Already given in the booklet, but NO2 is brown in color and to my knowledge, it is very difficult to observe it. Common sense helps a lot. Otherwise use a white paper or tile as background while observing. 2. Ammonia is pungent, very disturbing...sometimes they make you smell it. YES, they do! Don't forget to use wet red litmus to test ammonia as well. Red litmus turns into blue because ammonia is a base. 3. If you are not familiar with the smells of gases, you can confuse colorless gases which smell differently. H2S, for example, smells like rotten eggs while SO2 smells like the smoke just after you burn a match.

General Tips 1. Read carefully the names of all solutions..otherwise you will have to suffer. Sometimes you might have to start from the beginning, it will kill your confidence as well as time. Unlike physics practical, you need to be very

fast...It is very unlikely to happen that you will finish before time. 2. Don't spend time washing burettes and all...do it real quick!!! Like 1 burette washing in 10 seconds. 3. If you think some solutions have been mixed or there is impurity or even if your intuition says daal mey kuch kaaala hey, then ask them to replace it! They will not scold you like Minnu mam. 4. You make get confused with droppers. It helps making marks on them..whatever, don't use same dropper for two solutions. Believe me, it will affect your results. 5. If strong heating is directed, you should STRONG heat it, not just warming...don't just write the observation fast. 6. Some examiners might be strict, they may not provide you chemicals again until empty. Be careful in spending. 7. Use rationally pipette, burette or measuring cylinder. 8. Quote from the syllabus: "Candidates should normally record burette readings to the nearest 0.05 cm^3 and temperature readings to the nearest 0.5 C when using thermometer calibrated in 1 C intervals, and to the nearest 0.1 C when the interval is 0.2 C" 9. Talking about errors, an example: Error in any instrument is half of the smallest unit you can read on it. For instance, the stopwatch we had in the lab had the smallest unit as a 0.01s so the error in the readings read would be +/- 0.005s. Now you can find the uncertainty using the following formula, it might give you an understanding why it depends on the size of the value you measure. Ex. If I measure a value of 50s using the same stopwatch as mentioned before with the error 0.005s. Then, Uncertainty = [(error)/(my measured value)]*100 = [0.005/50]*100 = 0.01 percent. But it is much more complex due to human reaction time in stopwatchs. They won't ask you, hopefully. :)

7. Lastly, even if the practical goes horrible(like with me), don't get your hopes down. You can easily manage to get 'a' if you work good on P1 and P2. I've seen students with 'U' on P3, still they managed to get overall 'a'. :)

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Chemistry Unit 3B notes Tests of ions: Ion

Formula Test Test 1 Heat the solid in a test tube with a Bunsen burner. It should decompose producing the oxide and carbon dioxide. E.g.

Carbonate

CO3

2-

Hydrogencarbonate HCO3-

Sulphate (VI)

SO4

2-

Sulphite

SO32-

Chloride

-

Cl

Test for the gas using limewater solution.

Observations

Limewater should turn from colourless to cloudy in the presence of carbon dioxide due to precipitation of calcium carbonate.

Test 2 Add dilute HCl to the solid. Test for the gas evolved using limewater solution.

Vigorous effervescence. Limewater should turn from colourless to cloudy in the presence of carbon dioxide due to precipitation of calcium carbonate.

Test Add calcium chloride to a hydrogencarbonate solution.

No precipitate forms since calcium hydrogencarbonate is soluble.

Test Add barium chloride solution acidified with dilute HCl to the test solution.

White precipitate of barium sulphate forms.

Test Warm the sulphite with dilute HCl. The solution turns Test for gases using acidified potassium green. dichromate(VI) solution (or paper) Test 1 Add concentrated sulphuric acid to the solid chloride.

White steamy acidic fumes are seen - HCl fumes.

Test 2 Add dilute nitric acid to a solution of a chloride to acidify the solution. This White precipitate of eliminates any carbonates or sulphites. AgCl forms. Add silver nitrate to the solution. Add dilute ammonia solution. Solid dissolves. Test 1 Add concentrated sulphuric acid to the solid bromide. Bromide

Br-

Test 2' Add dilute nitric acid to a solution of a bromide to acidify the solution. This Cream precipitate of eliminates any carbonates or sulphites. AgBr forms. Add silver nitrate to the solution. Add concentrated ammonia solution. Solid dissolves. Test 1 Add concentrated sulphuric acid to the solid iodide.

Iodide

-

I

Steamy brownish acidic fumes are seen.

Purple acidic fumes are seen. The mixture turns to a brown slurry.

Test 2 Add dilute nitric acid to a solution of a iodide to acidify the solution. This Yellow precipitate of eliminates any carbonates or sulphites. AgI forms. Add silver nitrate to the solution.

Solid is insoluble.

Add concentrated ammonia solution.

Test 1 Heat solid nitrate.

Nitrate

NO3-

If group 1 solid (not Li) then will decompose to give the nitrite and oxygen. All other solid nitrates decompose to give the metal oxide, nitrogen dioxide and oxygen.

Oxygen gas is evolved that will relight a glowing splint. Brown gas is seen (NO2). Oxygen gas is also evolved and will relight a glowing splint.

Test 2 Boil nitrate solution with aluminium/Devarda’s alloy, in sodium hydroxide solution. Test vapour with red litmus paper.

Ammonium

NH4+

Test Warm ammonium compound with NaOH. Test vapours immediately using damp red litmus paper.

Litmus paper turns blue in the presence of ammonia.

NH 3 turns the litmus paper blue.

Test Dip nichrome wire in HCl. Lithium

Li+

Dip wire in solid. Heat wire in centre of flame. Observe colour of flame.

A carmine red flame is seen.

Test Dip nichrome wire in HCl. Sodium

Na+

Dip wire in solid. Heat wire in centre of flame.

A yellow flame is seen.

Observe colour of flame. Test Dip nichrome wire in HCl. Potassium

K+

Dip wire in solid.

A lilac flame is seen.

Heat wire in centre of flame. Observe colour of flame.

Magnesium

Mg2+

Test Add NaOH solution to the magnesium solid.

A white solid forms which is insoluble in excess NaOH(aq). This is Mg(OH)2(s)

Test Dip nichrome wire in HCl. Ca2+

Calcium

Dip wire in solid.

A brick red flame is seen.

Heat wire in centre of flame. Observe colour of flame. Test Dip nichrome wire in HCl. Sr2+

Strontium

Dip wire in solid.

A crimson red flame is seen.

Heat wire in centre of flame. Observe colour of flame. Test Dip nichrome wire in HCl. Ba2+

Barium

Dip wire in solid.

A apple green flame is seen.


Tests of gases: Name

Formula Test

Observations

Hydrogen

H2

Ignite gas.

Squeaky pop is heard.

Oxygen

O2

Place a glowing splint in a sample of the gas.

The glowing splint relights.

Carbon dioxide

CO2

A solution turns from colourless to Bubble gas through cloudy. A white (milky) precipitate of limewater (saturated solution calcium carbonate forms which is of calcium hydroxide) sparingly soluble.

Ammonia

NH3

Test for gas using damp red

Litmus paper turns blue.

litmus paper.

Chlorine

Cl2

Test 1 Test for gas using damp litmus paper (red or blue)

Chlorine bleaches the litmus paper very quickly.

Test 2 Test for gas using moist starch-iodide paper.

The paper turns blue-black.

Test 3 The solution turns from colourless to gas through a solution of orange. a bromide. Test 4 The solution turns from colourless to gas through a solution of brown (possibly with a black an iodide. precipitate, iodine).

Nitrogen dioxide

Sulphur dioxide

NO2

SO2

Not many tests for this gas.

The gas is brown.

Test 1 Bubble gas through a solution The solution turns from orange to of potassium dichromate (VI) green. dissolved in sulphuric acid. Test 2 Bubble gas through a solution The solution turns from purple to of potassium manganate (VII) colourless. dissolved in sulphuric acid.

Volumetric analysis: Volumetric analysis (titration) involves the reaction between two solutions. For one solution, both the volume and the concentration are known; for the other, the volume only is known. Apparatus used includes a burette, a pipette and a volumetric flask.

What is a standard solution? A solution for which concentration is accurately known. The concentration may have been found by a previous titration or by weighing the solute and making a solution of known volume. Such a solution is a primary standard solution.

How is a 250cm3 standard solution prepared?               

Make sure that the balance is clean and dry. Wipe it with a damp cloth. Place the weighing bottle on the pan and take the balance (i.e. re-zero it) Take the bottle off the balance and add solid to it. This ensures that no spillages fall on the pan. Take the balance of the weighing bottle + solid and find the balance of solid by subtraction. Replace on balance, and if the required amount is added, withdraw the mass. When you have the required amount, write its value down immediately. Wash out a 250cm3 volumetric flask three times using pure water. Transfer the solid to a 250cm3 volumetric flask using a funnel, and wash out the weighing bottle into the flask through the funnel using distilled water. Add about 100cm3 of distilled water to the flask. Stir the solution using a glass rod until all the solid visibly dissolves into a solution. Wash all remaining apparatus including the glass rod, funnel and transfer the rest of this to the flask. Make up to 250cm3 with distilled water so that the bottom of the meniscus just touches the 250cm3 mark. Attach a stopper to the flask. Shake the flask vigorously and/or invert the flask 5 or 6 times to create a solution with uniform concentration. Concentration of solution = mass of solid used/molar mass of solid x 1000/250 (units moldm-3)

Using the pipette:   



A glass bulb pipette will deliver the volume stated on it within acceptable limits only. Using a pipette filler, draw a little of the solution to be used into the pipette and use this to rinse the pipette. Fill the pipette to about 2-3cm3 above the mark. Pipette fillers are difficult to adjust accurately, so quickly remove the filler and close the pipette with your forefinger (not thumb). Release the solution until the bottom of the meniscus is on the mark. Immediately transfer the pipette to the conical flask in which you will do the titration, and allow the solution to dispense under gravity.

Using the burette:    

 

  

Making sure that the tap is shut, add about 10-15cm3 of the appropriate solution to the burette and rinse it out, not forgetting to open the tap and rinse the jet. Close the tap and fill the burette. A small funnel should be used to add the solution but be careful not to overfill the funnel. Remove the funnel, because titrating with a funnel in the burette can lead to serious error if a drop of liquid in the funnel stem falls into the burette during the titration. Bring the meniscus on to the scale by opening the tap to allow solution to through the burette. There is no particular reason to bring the meniscus exactly to the zero mark. Make sure that the burette is full to the tip of the jet. After a suitable indicator has been added to the solution in the conical flask, swirl the flask under the burette with one hand whilst adjusting the burette tap with your other hand. Add the solution in the burette to the conical flask slowly, swirling the flask all the time. As the endpoint is approached, the indicator will change colour more slowly. The titrant should be added drop by drop near to the endpoint. Repeat the titration until you have three concordant titres, i.e. volumes that are similar. This means within 0.2cm3 or better if you have been careful. Taking the mean of three tires that differ by 1cm3 or more is no guarantee of an accurate answer.

Common indicators:  

Methyl orange - yellow in alkali, red in acid & orange in neutral solutions(usually the end point of a titration) Phenolphthalein - pink in alkali, colourless in acid.

Enthalpy change measurements:        

Weigh a spirit lamp (containing a liquid alcohol) using a balance accurate to 3 decimal places. Record the mass measured. Use a measuring cylinder to put 100 cm3 of distilled water into a small beaker and clamps this at a fixed height above the spirit lamp (about 2 cm). Record the initial temperature of the water using a thermometer. Light the lamp using a burning splint. Heat the water using the spirit lamp until the temperature has gone up by about 10C. Stir the water with the thermometer the whole time. Put a cap on the spirit to stop the alcohol burning. The lid stops also stops further evaporation of the liquid alcohol. Reweigh the spirit lamp and record the mass. Calculate the enthalpy change.

Possible sources of error:     

There may be heat loss due to the apparatus used and heat may have dissipated through the insulating material --> should use a polystyrene cup and insulation like a lid. The specific heat capacity and density of water are used (and not of HCl). The masses of solid added to the acid are ignored. It is assumed that the specific heat capacity of the polystyrene cup is negligible. Some heat is lost when the hydrogen or carbon dioxide are evolved in the reactions.

Organic Procedures: Heating under reflux - This allows reactions to occur slowly, over a long period of time, without any loss of volatile liquid. The solvent evaporates and is condensed and returns to the flask.

Many organic reactions are slow and require prolonged heating   

To achieve this without loss of liquid, reaction mixtures are heated in a flask carrying a vertical condenser. This is heating under reflux; the solvent is condensed and returned to the flask, so the mixture can be heated as long as desired. To heat the round bottomed flask, either use a water bath, an oil bath or a heated plate mantle. A Bunsen burner isn’t really suitable.

Simple distillation: To separate a volatile solvent from a mixture    

Simple distillation is used where a volatile component has to be separated from a mixture, the other components of the mixture being very much volatile or non-volatile. The mixture is heated. The fraction that boils is collected within the temperature range of the fraction. (normally 1 or 2 degrees before the boiling temperature) The condenser cools the fraction so it distils and is collected in the receiving flask.

Fractional distillation: To separate mixtures of volatile liquids. 

Re-crystallisation - Used to purify a solid material by removing both soluble and insoluble impurities. The choice of solvent is important. The substance must be easily soluble in the boiling solvent and much less soluble at room temperature. This ensures the smallest possible loss of material, although some loss is inevitable with this technique.

Re-crystallisation method: 1. Dissolve the solid in the minimum amount of boiling solvent. This ensures that the solution is saturated with respect to the main solute but not with respect to the impurities, which are present in much smaller amounts. 2. Filter the hot mixture through a preheated filter funnel. This removes insoluble impurities. The hot funnel is necessary to prevent the solute crystallising and blocking the funnel. Filtration under vacuum using a Buchner funnel is often preferred, because it is fast. 3. Cool the hot filtrate, either to room temperature or, if necessary, in a bath of iced water. Rapid cooling gives small crystals, slow cooling large ones. The large crystals are often less pure. 4. Filter the cold mixture using a Buchner funnel. 5. Wash the crystals with a small amount of cold solvent. This removes any impurity remaining on the surface of the crystals. A small amount of cold solvent is used so that the crystals aren’t washed away / don’t dissolve. 6. Suck the crystals as dry as possible on the filter. 7. Transfer the crystals to a desiccator to dry. Drying between filter paper is sometimes recommended, but it is a very poor method.

Melting point determination: This is used to determine the purity of the re-crystallisation solid. Place small amount of the solid in the sealed end of a capillary tube. Place in the melting point apparatus. A sharp melting point over a small range shows purity, when compared with the set-book value of a higher melting point, that indicates an impure solid.

Organic tests:     



Collect 10 cm³ of the samples. Test the samples in the following order Alkenes – bromine water --> decolourises --> alkene Alcohols – Spatula of solid PCl5. Test fumes with damp litmus paper --> litmus red? White fumes near ammonia? --> alcohol Halogenoalkane – Add NaOH, ethanol as solvent. Shake and warm for 3 minutes. Cool and add nitric acid + silver nitrate. --> white = chloride; cream = bromide; yellow = iodide. --> Confirm with ammonia Alkane, the substance left is the alkane.

Group 1 and 2 reactions: Sulphate solubility: If a solution of any sulphate is added to a solution of a group 2 metal compound then a precipitate is likely. Group 2 ion in solution Effect of adding a sulphate solution Mg2+

No precipitate, MgSO4 is soluble

Ca2+

White precipitate of CaSO4

Sr2+

White precipitate of SrSO4

Ba2+

White precipitate of BaSO4

Hydroxide solubility: If sodium hydroxide is added to a solution of a group 2 compound then a precipitate is likely. Group 2 ion in solution Effect of adding a hydroxide solution Mg2+

Faint white precipitate of Mg(OH)2

Ca2+

Faint white precipitate of Ca(OH)2

Sr2+

Faint white precipitate of Sr(OH)2 on standing

Ba2+

No precipitate, Ba(OH)2 is soluble

Heating carbonates and nitrates: Substance

Effect of heat

Lithium and all group 2 carbonates Carbon dioxide detected Sodium and potassium carbonates No effect (except water of crystallisation may be given off) Sodium and potassium nitrates

Oxygen only gas evolved

Lithium and all group 2 nitrates

Nitrogen dioxide and oxygen evolved

Action of heat on compounds:    

Carbonates - Carbon dioxide is given off. Hydrogencarbonates - Carbon dioxide and water formed. Group 1 nitrates - Nitrite and oxygen formed. Group 2 nitrates - Oxide, brown fumes of nitrogen dioxide and oxygen formed.

Distinguishing between hydrocarbons:  

Alkane: Burn /oxidise/combust them. They will burn with a yellow flame and form CO2 and H2O (limited supply of CO). Alkene: A yellow, sootier flame is produced (due to the extra carbon and higher ratio of carbon:hydrogen).

Identifying some functional groups: 

Alkene - Add to orange bromine water. The alkene will decolourise it.



Halogenalkane - Heat with sodium hydroxide solution. Acidify with dilute nitric acid and then test with silver nitrate solution as with inorganic halides.



Alcohols or carboxylic acids containing C-OH - In a dry test tube (i.e. dry alcohol), add PCl5. Misty fumes of HCl are produced, which turns blue litmus paper red.

Distinguishing between different classes of alcohol: Primary - Add PCl5. Warm it with aqueous potassium dichromate (K2Cr2O7) and dilute H2SO4. Misty fumes are given off and the colour changed from orange to green --> aldehyde. Secondary - Misty fumes, changes to green --> ketone. Tertiary - Misty fumes, no colour change.

Halogens are toxic and harmful by inhalation, although iodine is much less so than chlorine or bromine, because it is a solid. Chlorine and bromine must always be used in a fume cupboard. Liquid bromine causes serious burns an must be handled with gloves. Ammonia is toxic. Concentrated ammonia solutions should be handled in the fume cupboard. Concentrated mineral acids are corrosive. If spilt on the hands, washing with plenty of water is usually enough, but advice must be sought. Acid in the eye requires immediate attention and prompt professional medical attention. Barium chloride solution and chromates and dichromates are extremely poisonous and so should be used in the fume cupboard/should not be inhaled. Sodium or potassium hydroxide or concentrated ammonia in the eye is extremely serious and must always receive professional and immediate attention. Sodium hydroxide and other alkali metal hydroxides are amongst the most damaging of all common substances to skin and other tissue. Wear gloves, goggles and an apron when handling these solutions in high concentrations.

General safety:     

Toxic/carcinogenic – use gloves, fume cupboard Flammable – Water baths, no naked fumes. Harmful gases – Use fume cupboard Corrosive – wear goggles/gloves Spillage of concentrated acid – wash with plenty of water.

EDEXCEL AS CHEMISTRY UNIT 3 NOTES Specification a) Recognise the results of reactions of compounds specifically mentioned in Units 1 and 2 of the specification and the results of tests for simple ions: carbonate, hydrogencarbonate, sulphate(VI), sulphite, chloride, bromide, iodide, nitrate, ammonium, lithium, sodium, potassium, magnesium, calcium, strontium and barium

Tests Ion

Formula Test

Observations

Test 1 Heat the solid in a test tube with a Bunsen burner.

Limewater should turn from It should decompose producing the oxide and colourless to cloudy in the presence of carbon dioxide due to carbon dioxide. E.g. precipitation of calcium carbonate. Carbonate

CO32Test for the gas using limewater solution. Test 2 Add dilute HCl to the solid. Test for the gas evolved using limewater solution.


Hydrogencarbonate HCO3-

Test No precipitate forms since calcium Add calcium chloride to a hydrogencarbonate hydrogencarbonate is soluble. solution.

Sulphate (VI)

SO42-



Sulphite

SO

Test Warm the sulphite with dilute HCl. Test for gases using acidified potassium dichromate(VI) solution (or paper)

The solution turns green.

Test 1 Add concentrated sulphuric acid to the solid chloride.

White steamy acidic fumes are seen - HCl fumes.

Chloride

Cl-

Bromide

Br-

23

Test 2 Add dilute nitric acid to a solution of a chloride to acidify the solution. This eliminates any White precipitate of AgCl forms. carbonates or sulphites. Add silver nitrate to the solution. Solid dissolves. Add dilute ammonia solution.

Test 1 Add concentrated sulphuric acid to the solid bromide.


Iodide

I-

Test 2' Add dilute nitric acid to a solution of a bromide to acidify the solution. This eliminates any carbonates or sulphites. Add silver nitrate to the solution. Add concentrated ammonia solution.

Cream precipitate of AgBr forms.

Test 1 Add concentrated sulphuric acid to the solid iodide.


Solid dissolves.

Test 2 Add dilute nitric acid to a solution of a iodide to acidify the solution. This eliminates any Yellow precipitate of AgI forms. carbonates or sulphites. Add silver nitrate to the solution.

Solid is insoluble.

Add concentrated ammonia solution. Test 1 Heat solid nitrate. If group 1 solid (not Li) then will decompose to give the nitrite and oxygen. Nitrate

NO3-

Oxygen gas is evolved that will relight a glowing splint.

All other solid nitrates decompose to give the Brown gas is seen (NO2). Oxygen metal oxide, nitrogen dioxide and oxygen. gas is also evolved and will relight a glowing splint. Test 2 Boil nitrate solution with aluminium/Devarda’s Litmus paper turns blue in the alloy, in sodium hydroxide solution. presence of ammonia. Test vapour with red litmus paper.

Ammonium

NH4+

Test NH Warm ammonium compound with NaOH. 3 Test vapours immediately using damp turns the litmus paper blue. red litmus paper. Test Dip nichrome wire in HCl.

Lithium

Li+

Dip wire in solid. Heat wire in centre of flame. A carmine red flame is seen. Observe colour of flame. Test Dip nichrome wire in HCl.

Sodium

Na+



Test Dip nichrome wire in HCl. Potassium

K+

Dip wire in solid.



Magnesium

Mg2+



Test Dip nichrome wire in HCl. Calcium

Ca2+

Dip wire in solid.



Test Dip nichrome wire in HCl. Strontium

Sr2+

Dip wire in solid.



Test Dip nichrome wire in HCl. Barium

Ba2+



b) Recognise the chemical tests for simple gases, to include hydrogen, oxygen, carbon dioxide, ammonia, chlorine, nitrogen dioxide and sulphur dioxide

Tests Name

Formula Test

Observations

Hydrogen

H2

Ignite gas.


Oxygen

O2



Carbon dioxide

CO2

Bubble gas through limewater (saturated solution of calcium hydroxide)

A solution turns from colourless to cloudy. A white (milky) precipitate of calcium carbonate forms which is sparingly soluble.

Ammonia

NH3

Test for gas using damp red litmus paper.


Test 1 Test for gas using damp litmus Chlorine bleaches the litmus paper very quickly. paper (red or blue)

Test 2 Test for gas using moist starch- The paper turns blue-black. iodide paper. Chlorine

Cl2 Test 3 gas through a solution of a The solution turns from colourless to orange. bromide.

Nitrogen dioxide

Sulphur dioxide

NO2

Test 4 gas through a solution of an iodide.

The solution turns from colourless to brown (possibly with a black precipitate, iodine).


The gas is brown.

Test 1 Bubble gas through a solution of potassium dichromate (VI) dissolved in sulphuric acid.

The solution turns from orange to green.

Test 2 Bubble gas through a solution of potassium manganate (VII) dissolved in sulphuric acid.

The solution turns from purple to colourless.

SO2

c) Describe the techniques used in volumetric analysis and enthalpy change measurements

Volumetric analysis Volumetric analysis (titration) involves the reaction between two solutions. For one solution, both the volume and the concentration are known; for the other, the volume only is known. Apparatus used includes a burette, a pipette and a volumetric flask.


How is a 250cm3 standard solution prepared? 

Make sure that the balance is clean and dry. Wipe it with a damp cloth.



Place the weighing bottle on the pan and take the balance (i.e. re-zero it)



Take the bottle off the balance and add solid to it. This ensures that no spillages fall on the pan.



Take the balance of the weighing bottle + solid and find the balance of solid by subtraction.



Replace on balance, and if the required amount is added, withdraw the mass.



When you have the required amount, write its value down immediately.



Wash out a 250cm3 volumetric flask three times using pure water.



Transfer the solid to a 250cm3 volumetric flask using a funnel, and wash out the weighing bottle into the flask through the funnel using distilled water.



Add about 100cm3 of distilled water to the flask.



Stir the solution using a glass rod until all the solid visibly dissolves into a solution.



Wash all remaining apparatus including the glass rod, funnel and transfer the rest of this to the flask.



Make up to 250cm3 with distilled water so that the bottom of the meniscus just touches the 250cm3 mark.



Attach a stopper to the flask.



Shake the flask vigorously and/or invert the flask 5 or 6 times to create a solution with uniform concentration.



Concentration of solution = mass of solid used/molar mass of solid x 1000/250 (units moldm-3)

Using the pipette 

A glass bulb pipette will deliver the volume stated on it within acceptable limits only.



Using a pipette filler, draw a little of the solution to be used into the pipette and use this to rinse the pipette.



Fill the pipette to about 2-3cm3 above the mark. Pipette fillers are difficult to adjust accurately, so quickly remove the filler and close the pipette with your forefinger (not thumb). Release the solution until the bottom of the meniscus is on the mark.



Immediately transfer the pipette to the conical flask in which you will do the titration, and allow the solution to dispense under gravity.

Using the burette 

Making sure that the tap is shut, add about 10-15cm3 of the appropriate solution to the burette and rinse it out, not forgetting to open the tap and rinse the jet.



Close the tap and fill the burette. A small funnel should be used to add the solution but be careful not to overfill the funnel.



Remove the funnel, because titrating with a funnel in the burette can lead to serious error if a drop of liquid in the funnel stem falls into the burette during the titration.



Bring the meniscus on to the scale by opening the tap to allow solution to through the burette. There is no particular reason to bring the meniscus exactly to the zero mark.



Make sure that the burette is full to the tip of the jet.



After a suitable indicator has been added to the solution in the conical flask, swirl the flask under the burette with one hand whilst adjusting the burette tap with your other hand.



Add the solution in the burette to the conical flask slowly, swirling the flask all the time.



As the endpoint is approached, the indicator will change colour more slowly. The titrant should be added drop by drop near to the endpoint.



Repeat the titration until you have three concordant titres, i.e. volumes that are similar. This means within 0.2cm3 or better if you have been careful. Taking the mean of three tires that differ by 1cm3 or more is no guarantee of an accurate answer.

Common indicators 

Methyl orange - yellow in alkali, red in acid & orange in neutral solutions(usually the end point of a titration)



Phenolphthalein - pink in alkali, colourless in acid.

Enthalpy change measurements 

Weigh a spirit lamp (containing a liquid alcohol) using a balance accurate to 3 decimal places. Record the mass measured.



Use a measuring cylinder to put 100 cm3 of distilled water into a small beaker and clamps this at a fixed height above the spirit lamp (about 2 cm).



Record the initial temperature of the water using a thermometer.



Light the lamp using a burning splint.



Heat the water using the spirit lamp until the temperature has gone up by about 10C. Stir the water with the thermometer the whole time.



Put a cap on the spirit to stop the alcohol burning. The lid stops also stops further evaporation of the liquid alcohol.



Reweigh the spirit lamp and record the mass.



Calculate the enthalpy change

Possible sources of error 

There may be heat loss due to the apparatus used and heat may have dissipated through the insulating material --> should use a polystyrene cup and insulation like a lid.



The specific heat capacity and density of water are used (and not of HCl).



The masses of solid added to the acid are ignored.



It is assumed that the specific heat capacity of the polystyrene cup is negligible.



Some heat is lost when the hydrogen or carbon dioxide are evolved in the reactions.

d) Describe the techniques used in simple organic preparations such as distillation and heating under reflux

Heating under reflux - This allows reactions to occur slowly, over a long period of time, without any loss of volatile liquid. The solvent evaporates and is condensed and returns to the flask.

Many organic reactions are slow and require prolonged heating 

To achieve this without loss of liquid, reaction mixtures are heated in a flask carrying a vertical condenser.



This is heating under reflux; the solvent is condensed and returned to the flask, so the mixture can be heated as long as desired.



To heat the round bottomed flask, either use a water bath, an oil bath or a heated plate mantle. A Bunsen burner isn’t really suitable.

Simple distillation To separate a volatile solvent from a mixture 

Simple distillation is used where a volatile component has to be separated from a mixture, the other components of the mixture being very much volatile or non-volatile.



The mixture is heated.



The fraction that boils is collected within the temperature range of the fraction. (normally 1 or 2 degrees before the boiling temperature)



The condenser cools the fraction so it distils and is collected in the receiving flask.

Fractional distillation To separate mixtures of volatile liquids.



Re-crystallisation - Used to purify a solid material by removing both soluble and insoluble impurities. The choice of solvent is important. The substance must be easily soluble in the boiling solvent and much less soluble at room temperature. This ensures the smallest possible loss of material, although some loss is inevitable with this technique.

Re-Crystallisation Method 

Dissolve the solid in the minimum amount of boiling solvent. This ensures that the solution is saturated with respect to the main solute but not with respect to the impurities, which are present in much smaller amounts.



Filter the hot mixture through a preheated filter funnel. This removes insoluble impurities. The hot funnel is necessary to prevent the solute crystallising and blocking the funnel. Filtration under vacuum using a Buchner funnel is often preferred, because it is fast.



Cool the hot filtrate, either to room temperature or, if necessary, in a bath of iced water. Rapid cooling gives small crystals, slow cooling large ones. The large crystals are often less pure.



Filter the cold mixture using a Buchner funnel.



Wash the crystals with a small amount of cold solvent. This removes any impurity remaining on the surface of the crystals. A small amount of cold solvent is used so that the crystals aren’t washed away / don’t dissolve.



Suck the crystals as dry as possible on the filter.



Transfer the crystals to a desiccator to dry. Drying between filter paper is sometimes recommended, but it is a very poor method.

Melting point determination This is used to determine the purity of the re-crystallisation solid. Place small amount of the solid in the sealed end of a capillary tube. Place in the melting point apparatus. A sharp melting point over a small range shows purity, when compared with the set-book value of a higher melting point, that indicates an impure solid.

Organic tests 

Collect 10 cm³ of the samples.



Test the samples in the following order



Alkenes – bromine water --> decolourises --> alkene



Alcohols – Spatula of solid PCl5. Test fumes with damp litmus paper --> litmus red? White fumes near ammonia? --> alcohol



Halogenoalkane – Add NaOH, ethanol as solvent. Shake and warm for 3 minutes. Cool and add nitric acid + silver nitrate. --> white = chloride; cream = bromide; yellow = iodide. --> Confirm with ammonia



Alkane, the substance left is the alkane.

e) recall and interpret details of the chemistry of the elements and compounds listed in Units 1 and 2 of this specification this includes the chemistry of Groups 1, 2 and 7 and the chemistry associated with the organic compounds listed in topic 2.2

Sulphate solubility If a solution of any sulphate is added to a solution of a group 2 metal compound then a precipitate is likely.

Group 2 ion in solution

Effect of adding a sulphate solution

Mg2+


Ca2+


Sr2+


Ba2+


Hydroxide solubility If sodium hydroxide is added to a solution of a group 2 compound then a precipitate is likely. Group 2 ion in solution

Effect of adding a hydroxide solution

Mg2+


Ca2+


Sr2+


Ba2+


Heating carbonates and nitrates Substance

Effect of heat

Lithium and all group 2 carbonates

Carbon dioxide detected

Sodium and potassium carbonates

No effect (except water of crystallisation may be given off)

Sodium and potassium nitrates




Action of heat on compounds 

Carbonates - Carbon dioxide is given off.



Hydrogencarbonates - Carbon dioxide and water formed.



Group 1 nitrates - Nitrite and oxygen formed.



Group 2 nitrates - Oxide, brown fumes of nitrogen dioxide and oxygen formed.

Distinguishing between hydrocarbons 

Alkane: Burn /oxidise/combust them. They will burn with a yellow flame and form CO2 and H2O (limited supply of CO).



Alkene: A yellow, sootier flame is produced (due to the extra carbon and higher ratio of carbon:hydrogen).

Identifying some functional groups 

Alkene - Add to orange bromine water. The alkene will decolourise it.



Halogenalkane - Heat with sodium hydroxide solution. Acidify with dilute nitric acid and then test with silver nitrate solution as with inorganic halides.



Alcohols or carboxylic acids containing C-OH - In a dry test tube (i.e. dry alcohol), add PCl5. Misty fumes of HCl are produced, which turns blue litmus paper red.

Distinguishing between different classes of alcohol Primary - Add PCl5. Warm it with aqueous potassium dichromate (K2Cr2O7) and dilute H2SO4. Misty fumes are given off and the colour changed from orange to green --> aldehyde. Secondary - Misty fumes, changes to green --> ketone. Tertiary - Misty fumes, no colour change.

f) interpret quantitative and qualitative results See http://www.thestudentroom.co.uk/attachment.php?attachmentid=12846 for notes on calculations.

g) devise and plan simple experiments based on the chemistry and techniques summarised in a to e above Normally the last question on the paper. It will ask you to identify certain compounds from four. E.g. titrations, how to make a standard solution, how to titrate, identify QCO3 when heating it where Q is a group 2 metal ion.

h) evaluate error in quantitative experiments see Appendix I for material available to assist centres in teaching this area Percentage error = absolute uncertainty/actual value x 100%

j) comment on safety aspects of experiments based on supplied data or recall of the chemistry of the compounds listed in Units l and 2. Safety considerations should relate to specific experiments not be of a general nature it will be assumed that students wear eye protection during all practical work.

Halogens are toxic and harmful by inhalation, although iodine is much less so than chlorine or bromine, because it is a solid. Chlorine and bromine must always be used in a fume cupboard. Liquid bromine causes serious burns an must be handled with gloves. Ammonia is toxic. Concentrated ammonia solutions should be handled in the fume cupboard. Concentrated mineral acids are corrosive. If spilt on the hands, washing with plenty of water is usually enough, but advice must be sought. Acid in the eye requires immediate attention and prompt professional medical attention. Barium chloride solution and chromates and dichromates are extremely poisonous and so should be used in the fume cupboard/should not be inhaled. Sodium or potassium hydroxide or concentrated ammonia in the eye is extremely serious and must always receive professional and immediate attention. Sodium hydroxide and other alkali metal hydroxides are amongst the most damaging of all common substances to skin and other tissue. Wear gloves, goggles and an apron when handling these solutions in high concentrations.

General safety



Toxic/carcinogenic – use gloves, fume cupboard



Flammable – Water baths, no naked fumes.



Harmful gases – Use fume cupboard



Corrosive – wear goggles/gloves



Spillage of concentrated acid – wash with plenty of water.

Chemistry 3B Sulphate solubility If a solution of any sulphate is added to a solution of a group 2 metal compound then a precipitate is likely. Group 2 ion in solution

Effect of adding a sulphate solution

Mg2+


Ca2+


Sr2+


Ba2+


Hydroxide solubility If sodium hydroxide is added to a solution of a group 2 compound then a precipitate is likely. Group 2 ion in solution

Effect of adding a hydroxide solution

Mg2+


Ca2+


Sr2+


Ba2+


Heating carbonates and nitrates Substance

Effect of heat

Lithium and all group 2 carbonates

Carbon dioxide detected

Sodium and potassium carbonates

No effect (except water of crystallisation may be given off)

Sodium and potassium nitrates




Action of heat on compounds • • • •

Carbonates - Carbon dioxide is given off. Hydrogencarbonates - Carbon dioxide and water formed. Group 1 nitrates - Nitrite and oxygen formed. Group 2 nitrates - Oxide, brown fumes of nitrogen dioxide and oxygen formed.

Distinguishing between hydrocarbons •

Alkane: Burn /oxidise/combust them. They will burn with a yellow flame and form CO2 and H2O (limited supply of CO).

•

Alkene: A yellow, sootier flame is produced (due to the extra carbon and higher ratio of carbon:hydrogen).

Identifying some functional groups • • •

Alkene - Add to orange bromine water. The alkene will decolourise it. Halogenalkane - Heat with sodium hydroxide solution. Acidify with dilute nitric acid and then test with silver nitrate solution as with inorganic halides. Alcohols or carboxylic acids containing C-OH - In a dry test tube (i.e. dry alcohol), add PCl5. Misty fumes of HCl are produced, which turns blue litmus paper red.

Distinguishing between different classes of alcohol Primary - Add PCl5. Warm it with aqueous potassium dichromate (K2Cr2O7) and dilute H2SO4. Misty fumes are given off and the colour changed from orange to green --> aldehyde. Secondary - Misty fumes, changes to green --> ketone. Tertiary - Misty fumes, no colour change.

Specification f) interpret quantitative and qualitative results. g) devise and plan simple experiments based on the chemistry and techniques summarised in a to e above Normally the last question on the paper. It will ask you to identify certain compounds from four. E.g. titrations, how to make a standard solution, how to titrate, identify QCO3 when heating it where Q is a group 2 metal ion. h) evaluate error in quantitative experiments see Appendix I for material available to assist centres in teaching this area Percentage error = absolute uncertainty/actual value x 100% j) comment on safety aspects of experiments based on supplied data or recall of the chemistry of the compounds listed in Units l and 2. Safety considerations should relate to specific experiments not be of a general nature it will be assumed that students wear eye protection during all practical work. Halogens are toxic and harmful by inhalation, although iodine is much less so than chlorine or bromine, because it is a solid. Chlorine and bromine must always be used in a fume cupboard. Liquid bromine causes serious burns an must be handled with gloves. Ammonia is toxic. Concentrated ammonia solutions should be handled in the fume cupboard. Concentrated mineral acids are corrosive. If spilt on the hands, washing with plenty of water is usually enough, but advice must be sought. Acid in the eye requires immediate attention and prompt professional medical attention. Barium chloride solution and chromates and dichromates are extremely poisonous and so should be used in the fume cupboard/should not be inhaled. Sodium or potassium hydroxide or concentrated ammonia in the eye is extremely serious and must always receive professional and immediate attention. Sodium hydroxide and other alkali metal hydroxides are amongst the most damaging of all common substances to skin and other tissue. Wear gloves, goggles and an apron when handling these solutions in high concentrations.

General safety • • • • •

Toxic/carcinogenic – use gloves, fume cupboard Flammable – Water baths, no naked fumes. Harmful gases – Use fume cupboard Corrosive – wear goggles/gloves Spillage of concentrated acid – wash with plenty of water.

tests Ion

Formu Test la

Observations

Test 1 Heat the solid in a test tube with a Bunsen burner. It should decompose producing the oxide and carbon dioxide. E.g. Carbonate

CO32-

Hydrogencarbona HCO3te Sulphate (VI)

Sulphite

Chloride

Bromide

Iodide

SO42-

2-

SO3

Cl-

Br-

I-

Limewater should turn from colourless to cloudy in the presence of carbon dioxide due to precipitation of calcium carbonate.

Test for the gas using limewater solution. Test 2 Add dilute HCl to the solid. Test for the gas evolved using limewater solution.


Test Add calcium chloride to a hydrogencarbonate solution.

No precipitate forms since calcium hydrogencarbonate is soluble.



Test Warm the sulphite with dilute HCl. Test for gases using acidified potassium dichromate(VI) solution (or paper)

The solution turns green.

Test 1 Add concentrated sulphuric acid to the solid chloride.

White steamy acidic fumes are seen HCl fumes.

Test 2 Add dilute nitric acid to a solution of a chloride to acidify the solution. This eliminates any carbonates or sulphites. Add silver chloride to the solution. Add dilute ammonia solution.

White precipitate of AgCl forms.

Test 1 Add concentrated sulphuric acid to the solid bromide.


Test 2' Add dilute nitric acid to a solution of a bromide to acidify the solution. This eliminates any carbonates or sulphites. Add silver chloride to the solution. Add concentrated ammonia solution.

Cream precipitate of AgBr forms.

Test 1 Add concentrated sulphuric acid to the solid iodide.


Solid dissolves.

Solid dissolves.

Test 2 Add dilute nitric acid to a solution of a iodide to acidify the solution. This eliminates any carbonates or sulphites. Yellow precipitate of AgI forms. Add silver chloride to the solution.

Solid is insoluble.

Add concentrated ammonia solution. Test 1 Heat solid nitrate.

Nitrate

NO3-

If group 1 solid (not Li) then will decompose to give the nitrite and oxygen.

Oxygen gas is evolved that will relight a glowing splint.

All other solid nitrates decompose to give the metal oxide, nitrogen dioxide and oxygen.

Brown gas is seen (NO2). Oxygen gas is also evolved and will relight a glowing splint.

Test 2 Boil nitrate solution with aluminium/Devarda’s alloy, in sodium hydroxide solution. Test vapour with red litmus paper.

Ammonium

NH4+

Litmus paper turns blue in the presence of ammonia.

NH Test 3 Warm ammonium compound with NaOH. turns the litmus paper blue. Test vapours immediately using damp red litmus paper. Test Dip nichrome wire in HCl.

Lithium

Li+

Dip wire in solid. Heat wire in centre of flame. A carmine red flame is seen. Observe colour of flame. Test Dip nichrome wire in HCl.

Sodium

Na+

Dip wire in solid. Heat wire in centre of flame.


Observe colour of flame. Test Dip nichrome wire in HCl. Potassium

K+

Dip wire in solid.


Heat wire in centre of flame. Observe colour of flame. Magnesium

Mg2+



Calcium

Ca2+

Test


Dip nichrome wire in HCl. Dip wire in solid. Heat wire in centre of flame. Observe colour of flame. Test Dip nichrome wire in HCl. Sr2+

Strontium

Dip wire in solid.


Heat wire in centre of flame. Observe colour of flame. Test Dip nichrome wire in HCl. Ba2+

Barium

Dip wire in solid.



Tests Name

Formu Test la

Observations

Hydrogen

H2

Ignite gas.


Oxygen

O2



Carbon dioxide

CO2

Bubble gas through limewater (saturated solution of calcium hydroxide)

A solution turns from colourless to cloudy. A white (milky) precipitate of calcium carbonate forms which is sparingly soluble.

Ammonia

NH3

Test for gas using damp red litmus paper.


Test 1 Test for gas using damp litmus paper (red or Chlorine bleaches the litmus paper very quickly. blue)

Chlorine

Cl2

Test 2 The paper turns blue-black. Test for gas using moist starch-iodide paper. Test 3 gas through a solution of a bromide.

The solution turns from colourless to orange.

Test 4 gas through a solution of an iodide.

The solution turns from colourless to brown (possibly with a black precipitate, iodine).

Nitrogen dioxide

Sulphur dioxide

NO2

SO2


The gas is brown.

Test 1 Bubble gas through a solution of potassium The solution turns from orange to green. dichromate (VI) dissolved in sulphuric acid. Test 2 Bubble gas through a solution of potassium The solution turns from purple to colourless. manganate (VII) dissolved in sulphuric acid.

Volumetric analysis Volumetric analysis (titration) involves the reaction between two solutions. For one solution, both the volume and the concentration are known; for the other, the volume only is known. Apparatus used includes a burette, a pipette and a volumetric flask.


How is a 250cm3 standard solution prepared? • • • • • • • • • • • • • •

Make sure that the balance is clean and dry. Wipe it with a damp cloth. Place the weighing bottle on the pan and tare the balance (i.e. re-zero it) Take the bottle off the balance and add solid to it. This ensures that no spillages fall on the pan. When you have the required amount, write its value down immediately. Replace on balance, and if the required amount is added, withdraw the mass. Wash out a 250cm3 volumetric flask three times using pure water. Transfer the solid to a 250cm3 volumetric flask using a funnel, and wash out the weighing bottle into the flask through the funnel. Add about 100cm3 of distilled water to the flask. Stir the solution using a glass rod. Wash all remaining apparatus including the glass rod, funnel and transfer the rest of this to the flask. Make up to 250cm3 with distilled water so that the bottom of the meniscus just touches the 250cm3 mark. Stopper the flask. Shake the flask vigorously and/or invert the flask 5 or 6 times to dissolve the solid. Concentration of solution = mass of solid used/molar mass of solid x 1000/250 (units moldm-3)

Using the pipette • • •

•

A glass bulb pipette will deliver the volume stated on it within acceptable limits only. Using a pipette filler, draw a little of the solution to be used into the pipette and use this to rinse the pipette. Fill the pipette to about 2-3cm3 above the mark. Pipette fillers are difficult to adjust accurately, so quickly remove the filler and close the pipette with your forefinger (not thumb). Release the solution until the bottom of the meniscus is on the mark. Immediately transfer the pipette to the conical flask in which you will do the titration, and allow the solution to dispense under gravity.

Using the burette •

Making sure that the tap is shut, add about 10-15cm3 of the appropriate solution to the burette and rinse it out, not forgetting to open the tap and rinse the jet.

• • • • • • • •

Close the tap and fill the burette. A small funnel should be used to add the solution but be careful not to overfill the funnel. Remove the funnel, because titrating with a funnel in the burette can lead to serious error if a drop of liquid in the funnel stem falls into the burette during the titration. Bring the meniscus on to the scale by opening the tap to allow solution to through the burette. There is no particular reason to bring the meniscus exactly to the zero mark. Make sure that the burette is full to the tip of the jet. After a suitable indicator has been added to the solution in the conical flask, swirl the flask under the burette with one hand whilst adjusting the burette tap with your other hand. Add the solution in the burette to the conical flask slowly, swirling the flask all the time. As the endpoint is approached, the indicator will change colour more slowly. The titrant should be added drop by drop near to the endpoint. Repeat the titration until you have three concordant titres, i.e. volumes that are similar. This means within 0.2cm3 or better if you have been careful. Taking the mean of three tires that differ by 1cm3 or more is no guarantee of an accurate answer.

Common indicators • •

Methyl orange - yellow in alkali, red in acid (orange at end point) Phenolphthalein - pink in alkali, colourless in acid.

Enthalpy change measurements • • • • • • • •

Weigh a spirit lamp (containing a liquid alcohol) using a balance accurate to 3 decimal places. Record the mass measured. Use a measuring cylinder to put 100 cm3 of distilled water into a small beaker and clamps this at a fixed height above the spirit lamp (about 2 cm). Record the initial temperature of the water using a thermometer. Light the lamp using a burning splint. Heat the water using the spirit lamp until the temperature has gone up by about 10C. Stir the water with the thermometer the whole time. Put a cap on the spirit to stop the alcohol burning. The lid stops also stops further evaporation of the liquid alcohol. Reweigh the spirit lamp and record the mass. Calculate the enthalpy change

Possible sources of error • • • • •

There may be heat loss due to the apparatus used and heat may have dissipated through the insulating material --> should use a polystyrene cup and insulation like a lid. The specific heat capacity and density of water are used (and not of HCl). The masses of solid added to the acid are ignored. It is assumed that the specific heat capacity of the polystyrene cup is negligible. Some heat is lost when the hydrogen or carbon dioxide are evolved in the reactions.

Many organic reactions are slow and require prolonged heating • •

To achieve this without loss of liquid, reaction mixtures are heated in a flask carrying a vertical condenser. This is heating under reflux; the solvent is condensed and returned to the flask, so the mixture can be heated as long as desired.

•

To heat the round bottomed flask, either use a water bath, an oil bath or a heated plate mantle. A Bunsen burner isn’t really suitable.

Simple distillation To separate a volatile solvent from a mixture • Simple distillation is used where a volatile component has to be separated from a mixture, the other components of the mixture being very much volatile or non-volatile. • The mixture is heated. • The fraction that boils is collected within the temperature range of the fraction. (normally 1 or 2 degrees before the boiling temperature) • The condenser cools the fraction so it distils and is collected in the receiving flask.

Fractional distillation To separate mixtures of volatile liquids. • Re-crystallisation - Used to purify a solid material by removing both soluble and insoluble impurities. The choice of solvent is important. The substance must be easily soluble in the boiling solvent and much less soluble at room temperature. This ensures the smallest possible loss of material, although some loss is inevitable with this technique.

Re-crystallisation method 1. Dissolve the solid in the minimum amount of boiling solvent. This ensures that the solution is saturated with respect to the main solute but not with respect to the impurities, which are present in much smaller amounts.

2. Filter the hot mixture through a preheated filter funnel. This removes insoluble impurities. The hot funnel is necessary to prevent the solute crystallising and blocking the funnel. Filtration under vacuum using a Buchner funnel is often preferred, because it is fast.

3. Cool the hot filtrate, either to room temperature or, if necessary, in a bath of iced water. Rapid cooling gives small crystals, slow cooling large ones. The large crystals are often less pure.

4. Filter the cold mixture using a Buchner funnel. 5. Wash the crystals with a small amount of cold solvent. This removes any impurity remaining on the surface of the crystals. A small amount of cold solvent is used so that the crystals aren’t washed away / don’t dissolve.

6. Suck the crystals as dry as possible on the filter. 7. Transfer the crystals to a desiccator to dry. Drying between filter paper is sometimes recommended, but it is a very poor method.

Melting point determination This is used to determine the purity of the re-crystallisation solid. Place small amount of the solid in the sealed end of a capillary tube. Place in the melting point apparatus. A sharp melting point over a small range shows purity, when compared with the set-book value of a higher melting point, that indicates an impure solid.

Organic tests • • •

Collect 10 cm³ of the samples. Test the samples in the following order Alkenes – bromine water --> decolourises --> alkene

• •

•

Alcohols – Spatula of solid PCl5. Test fumes with damp litmus paper --> litmus red? White fumes near ammonia? --> alcohol Halogenoalkane – Add NaOH, ethanol as solvent. Shake and warm for 3 minutes. Cool and add nitric acid + silver nitrate. --> white = chloride; cream = bromide; yellow = iodide. --> Confirm with ammonia Alkane, the substance left is the alkane.

Techniques Separating insoluble impurities from a soluble substance (Removing sand and impurities from salt solution)

Separating a mixture of immiscible liquids (Separating a mixture of water and hexane) Water and hexane are immiscible forming 2 separate layers and are separated using a separating funnel Separating a solvent from solution Simple distillation

Separating a liquid from a mixture of miscible liquids Fractional distillation Separates mixtures of miscible liquids with different Bt’s, using a fractionating column increasing efficiency of redistillation process, packed with inert material(glass beads) increasing surface area where vapour may condense. - When mixture is boiled vapours of most volatile component(lowest Bt) rises into the vertical column where they condense to liquids. - As they descend they are reheated to Bt by the hotter rising vapours of the next component. - Boiling condensing process occurs repeatedly inside the column so there is a temperature gradient. - Vapours of the more volatile components reach the top of the column and enter the condenser for collection

Boiling under reflux Where reagents volatile - condenses vapours and returns reagents to flask, prevents loss of reactants/products, prolonged heating for slow reactions - For preparation of aldehyde/carboxylic acid from alcohol (1)Reason for heating the mixture but then taking the flame away (1)provide Ea, exothermic/prevent reaction getting out of control

Separating mixtures of similar compounds in solution (Separating dyes present in a sample of ink) Chromatogram Chromatography Different components of the dye spread out at different rates Using a square sheet of filter paper, spots of dye solutions are put along the baseline The filter paper is coiled into a cylinder and placed in a tank containing a small volume of solvent The lid is replaced on the tank, solvent rises up the filter paper When the solvent nearly reaches the top of the filter paper, the filter paper is removed and position of solvent marked. Dyes A & B are either pure substances or a mixture of dyes not separated with the solvent used Dye C is composed of A & B as the spots correspond Colourless substances can be separated and seen by spraying/dipping the filter paper into a locating agent which colours the spots produced

Separating a solid which sublimes, from a solid which doesn’t sublime Given a mixture of Ammonium chloride(sublimes) and sodium chloride(doesn’t sublime) Heat the mixture. Ammonium chloride turns directly to vapour but the sodium chloride remains unchanged When the vapour is cooled solid ammonium chloride collects free from sodium chloride A pure substance has a definite Mt, presence of impurities causes the substance to melt over a range of temperatures Best method of separation of (1) Oil and water (2) Alcohol and water (3) Nitrogen from liquid air (1) Separating funnel(2) Fractional distillation (3) Fractional distillation Mixture Compound - Proportions of the different elements can be varied - Different elements have to be present in fixed proportions - Properties are those of the elements making it up - Properties different from properties of elements making it up - Elements can be separated by simple methods - Difficult to separate into the elements which make it up - No energy gained or lost when the mixture is made - Energy usually given out/taken in when compound is formed Sub-atomic particles Protons, neutrons and electrons which makes up the atom Particle An atom, molecule, ion, electron or any identifiable particle RTP Room temperature and pressure Electron A negatively charged particle, with negligible mass occupying the outer regions of all atoms Immiscible Unable to mix, dissolve in each other, to form a homogenous mixture Miscible Soluble in each other (aq) Substance dissolved in water to form an aqueous solution State symbols Physical state of the reactants at RT Aqueous(aq) Solvent Substance in which other substances are dissolved Solute Substance dissolved in another substance(solvent)to form a solution Chemical species Collection of particles Distilled water Water that has been purified by distillation Ion When number of protons and electrons are different Atom The smallest part of an element that can exist on its own Molecule 2 or more atoms bonded together Element A pure substance which can’t be split up by chemical reaction Compound Combination of elements in fixed proportions via synthesis In formation of a compound from ions the charges balance out Physical properties: Mt, Bt, hardness • Compounds ending in –ate –ite contain oxygen, greater proportion of oxygen in the compound ending in –ate

Sodium sulphate Na2SO4 Sodium sulphite Na2SO3 • Compounds with prefix per– contain extra oxygen Sodium oxide Na2O Sodium peroxide Na2O2 • Compounds with prefix thio– contain a sulphur atom in place of an oxygen atom Sodium sulphate Na2SO4 Sodium thiosulphate Na2S2O3 Metalloid Element which has properties between metals and nonmetals - Ions in an ionic compound are tightly held together in a regular lattice, lattice energy is required to break it up and melt the substance A metal high in the reactivity series has stable ores and the metal can be obtained only by electrolysis A metal middle in the reactivity series doesn’t form stable ores and can be extracted by reduction reactions (often with carbon) A metal low in the reactivity series, if present in unstable ores can be extracted by heating Decomposition Splitting up of a compound (Thermal decomposition - decomposition of a compound by heating) Combustion is the reaction of a substance with oxygen, total mass of products is greater than the mass of the substance burned, difference being the mass of oxygen combined Sublimation of an element/compound is a transition from solid to gas with no intermediate stage When a change of state takes place the temperature remains constant despite a continuing supply of energy. Latent heat is the energy which is not being used to raise the temperature and supplies particles with the extra energy they require as the state changes(given out when the reverse changes take place) Fe(s) + CuSO4(aq)  FeSO4(aq) + Cu(s) Blue solution turns colourless and brown copper is deposited A displacement reaction where a more reactive metal replaces a less reactive metal in a compound Electrolysis of HCl(aq): 2HCl(aq)  Cl2(g) + H2(g) Cathode: 2H+(aq) + 2e– H2(g) Anode: 2Cl –(aq)Cl2(g)+2e– Heat Ammonium chloride NH4Cl(s) NH3(g) + HCl(g) Cool A stopper from a bottle of (conc)NH3(aq) held near a stopper from a bottle of (conc)HCl acid gives a dense white smoke of NH4Cl Synthetic Pathways(Series of reactions built up to convert one functional group into another) • Reactions of functional groups assumed to be the same whether molecules are simple or complicated • Synthesis of the product molecule possible because in any reaction of a functional group a product is formed capable of conversion into other molecules

– 5.31 × 4.18 × ΔT mass of G used (g)

–1 Would the result for ∆H be more accurate if the temperature of the solution were known to 3dp? ∆H = kJ mol Yes, temperature would then be known to a comparable precision to the other factors in the equation Plan an experiment to investigate concentration on rate of a reaction Mg(s) + 2HCl(aq) → MgCl (aq) + H (g) 2 2 • Apparatus diagram/description/addition of Mg to acid in appropriate container • Weigh Mg/cut measured length and measure volume of acid, measure volume of H with time/time how long it takes for Mg to 2 ’dissolve’ • repeat with different concentration(s) HCl • repeat with same mass/length Mg and same volume HCl • one axis labelled concentration/volume H and other axis time/t axis changed to 1/t 2 if measured volume H at least two curves showing increasing rate with conc 2 if measured time to dissolve Mg one line showing time decreasing with conc or 1/t increasing with conc • Acid irritant so wear gloves or hydrogen explosive - no naked flame or Build up of pressure in syringe - attach plunger with string – – (NH ) CO reacts with both 1moldm ³(dil)HNO and 1moldm ³(dil)KOH in the ratio 1:2 42 3 3 Devise an experiment to determine which of the two reactions is the more exothermic • Prepare solutions of known concentration of the solid

• Suggestion of apparatus used e.g. lagged calorimeter or low mass polystyrene cup • Use same volume of each solution • Measure maximum temperature change • Improve reliability of results, repeat experiment • Possible sources of error identified • Reaction with the greater temperature change is the more exothermic • Since ammonia evolved use a fume cupboard M = Group 1 M CO (s) + 2HCl(aq) → 2MCl(aq)+CO (g)+H O(l) 2 3 2 2 Plan an experiment, results of which used to calculate the relative molecular mass of the carbonate and identify M 3 At temp of experiment 1 mole of CO occupies a volume of 24dm 2 G a s s y rin g e Relative atomic mass: Li = 7, C = 12, O = 16, Na = 23, K = 39, Rb = 85, Cs = 133

• Add M CO + acid and stopper flask/use suspended test tube in large flask 2 3 • Mass of M CO and (final)vol of CO /when effervescence stops record volume of gas in syringe 2 3 2 • Explain conversion volume of CO to moles by correct use of 24 2 • Use of moles M CO = mass M CO ¸ M M CO to find M • Hence find identity of M 2 3 2 3 r 2 3 r • Escape of gas before bung replaced/solid did not all react/CO soluble in acid • Eye protection-acid hazard 2 • Should not affect identification since even if M is slightly wrong it will still correspond to nearest Group 1 metal atomic mass r –1 Plan an experiment to identify an acidic compound, molar mass for an acid estimated to be 88 ± 2 g mol What simple test would allow butenoic acid to be distinguished from the other two? Devise a plan based upon a quantitative experiment that would allow the other two to be distinguished CO O H –1 –1 CH = CH CH COOH 86 g mol CH CH CH COOH 88 g mol –1 2 2 3 2 2 HC O O H 90 g mol 3 structures proposed for this acid Alkene detected using bromine water, only 1 would show decolourisation Fixed mass of acid, Standard named alkali, Controlled method of addition, To an identified end point of a named suitable indicator, Calculation of mole ratio (1)A sample of NaCl was thought to contain an impurity of Ba(NO3)2 A student suggested a flame test (a)(i)Suggest why a flame test on the mixture would not be a satisfactory way of detecting the presence of barium ions in the sample (i) • yellow/stronger/persistent Na flame • Obscures/Ba flame NOT “makes it difficult to distinguish between the two colours” (ii)Suggest a reagent that could be used to produce a ppt of a barium compound from a solution of the sample (ii) • (conc)H SO (solution) of any soluble sulphate (MgSO , (NH )2SO , Na SO ) 2 4 4 4 4 2 4 (2)(a)X(has OH group)decolourised cold potassium manganate(VII) acidified with (dil)H SO acid, structure of X suggested by this? 2 4 (a)carbon double bond (b)Complete oxidation of X with potassium dichromate(VI) solution and (dil)H2SO4 acid produces Y C H O structure for Y? 4 6 H

H

C

C

H C

C

H

C H

2

C H

CO

CH

3

(1 )

H

H

H

H

C

C

C

C

H

(2 )

O H H O H (b) H (c)Structural formula for X?(c) H X must be a secondary alcohol because ketone formed on oxidation carboxylic acid is not formed

+ – (a)CH H Br + H O → C H OH + H + Br 4 9 2 4 9 (b)Suggest why ethanol was used in the experiment (b)Solvent/silver nitrate solution in water and bromobutane immiscible (c)Suggest a reason for the use of a water bath (c)Reaction slow at RT/increases rate/flammable (4)Describe tests you would use to distinguish between the following pairs of compounds including results (a)NaNO and ammonium nitrate NH NO (a) Flame test, sodium salt gives yellow colour, ammonium salt gives no colour 3 4 3 (5)(a)1-bromobutane Bt102 °C may be prepared by the reaction C H OH + NaBr + H SO → C H Br + NaHSO + H O 4 9 2 4 4 9 4 2 (a)Describe how you would use distillation apparatus to give a sample of pure 1-bromobutane (a)Heat mixture(slowly), collect only distillate produced at around 102 °C at Bt of 1-bromobutane (b)Suggest 2 reasons why the actual yield was much lower than the max yield (b)side reactions, reaction incomplete, product lost in purification/transfers (3)(a)Write an ionic equation for the hydrolysis of 1-bromobutane by water

M a ss lo s s /g

2 .0 0 (6)CaCO (s) + 2HCl(aq) → CaCl (aq) + H O(l) + CO (g) 3 2 2 2 Experiment CaCO HCl acid 3 1 .5 0 1 RT Small pieces 3 –3 50cm of 1moldm 2 Small pieces 3 –3 50cm of 1moldm heated to 1 .0 0 E x p t. 80°C 3 RT One large 3 –3 50cm of 1moldm piece 0 .5 0 4 RT Small pieces 3 –3 50cm of 2moldm (a)(i)Explain why there is a loss in mass as the reaction proceeds (i)CO (g) evolved 5 10 15 2 T im e /m in (ii)Explain the shape of the curve drawn for Experiment 1 Results of Experiment 1 (ii)Reaction(fast at first then)slows down/gives off less CO per 2 min when line is horizontal, the reaction has finished/after 6 or 7 minutes/when 1 g of CO lost 2 (b)Draw curves on the graph to represent the results you would expect for Experiments 2, 3 and 4. Label the curves 2, 3 and 4 (b)Experiment 2 steeper than 1 and same mass loss Experiment 3 less steep than 1 and same mass loss/reaction incomplete Experiment 4 steeper than 1 and horizontal at twice mass loss (7)Suggest one appropriate safety precaution that should be taken as ethanedioic acid is toxic (7)Safety pipette filler 2+ 2+ (8)In an experiment to find ∆Η, zinc CuSO4(aq) in a plastic cup Zn(s) + Cu (aq) → Zn (aq) + Cu(s) Suggest reasons why a series of temp readings is taken rather than simply initial and final readings (8) Reason 1 Any fluctuations in temperature smoothed out / minimises reading error/allows line of best fit to be drawn Reason 2 Able to allow for cooling effect/able to calculate more accurate temperature change/need to find highest temperature

Another Test List (If you find above one hard to read) Flame test 1 Clean end of platinum/nichrome wire with(conc)HCl, burning off impurities in a roaring bunsen flame until there’s no persistent flame colouration 2 Moisten the end of the clean wire with (conc)HCl and then dip into the sample to be tested 3 Hold the sample at the edge of a roaring bunsen flame Lithium Carmine red Calcium Brick red Sodium Yellow Strontium Crimson Potassium Lilac Barium Apple green Gas Test Ammonia NH3 Pungent smell, Moist litmus paper red  blue, (conc)HCl at mouth of bottle, white smoke forms Carbon dioxide CO2 through lime water, turning lime water milky CaCO3(s) + H2O(l) + CO2(g) Ca(HCO3)2(aq) Limewater CO2 test Ca(OH)2(aq) + CO2(g)  CaCO3(s) White ppt + H2O(l) Chlorine Cl2 Swimming pool smell, moist litmus paper blue  red  bleached Hydrogen H2 Lighted splint, burns with squeaky pop Hydrogen chloride HCl Moist litmus paper blue  red Nitrogen(IV)oxide NO2 Brown gas, acrid smell, moist litmus paper blue  red Oxygen O2 Glowing splint, relights Water vapour H2O White anhydrous copper(II) sulphate white  blue CuSO4(s) + 5H2O(l) CuSO4.5H2O(s) Sulphur dioxide Cation Ammonium

SO2

NH4+ H+

Copper(II) Cu2+ Iron(II)

Fe2+

Or dry blue cobalt chloride paper blue  pink Acrid smell, moist litmus paper blue  red Or potassium dichromate(VI) solution/paper from orange  green Test Add NaOH(aq) Add NH4OH(aq) • Heat, ammonia evolved, moist litmus paper red  blue • Moist litmus paper blue  red • Add a carbonate, gas through lime water, CO2 evolved turning lime water milky • A little, blue ppt forms • A little, blue ppt forms • In excess, insoluble • In excess, dissolves and a dark blue solution forms • A little, green ppt forms • A little, green ppt forms

1

Iron(III)

Fe3+

Calcium

Ca2+

Magnesium

Mg2+

Aluminium

Al3+

Lead

Pb2+

Zinc

Zn2+

Anion Carbonate pH>10 CO32– universal indicator Hydrogen carbonate HCO3– pH 8-9 Chloride Cl– Bromide

Br–

Iodide

I–

Nitrate

NO3–

Sulphate

SO42–

Sulphite

SO32–

• In excess, insoluble • In excess, insoluble • A little, brown ppt forms • A little, brown ppt forms • In excess, insoluble • In excess, insoluble • A little, milky suspension forms • A little, milky suspension forms • In excess, insoluble • In excess, insoluble Distinguish Mg from Ca through flame test • A little, milky suspension forms • A little, milky suspension forms • In excess, insoluble • In excess, insoluble • A little, white ppt forms • A little, white ppt forms • In excess, dissolves giving colourless solution • In excess, insoluble No ppt with (dil)H2SO4/cold(dil)HCl/(dil)KI/(dil)Na2S Sodium sulphide • A little, white ppt forms • A little, white ppt forms • In excess, dissolves giving colourless solution • In excess, insoluble White ppt with (dil)H2SO4 White ppt with cold(dil)HCl Yellow ppt with (dil)KI Black ppt with (dil)Na2S Sodium sulphide • A little, white ppt forms • A little, white ppt forms • In excess, dissolves giving colourless solution • In excess, dissolves giving colourless solution Test Add (dil)HCl(aq) gas through lime water, CO2 evolved turning lime water milky Or add group II ions, white ppt or heat/add boiling water, no gas evolved Add (dil)HCl(aq) gas through lime water, CO2 evolved turning lime water milky Or add metal ions no ppt but heating causes white ppt to form or Heat/add boiling water, CO2 evolved Acidify with (dil)HNO3(aq) Add AgNO3(aq) Add (dil)NH3 to ppt White ppt AgCl forms Ppt dissolves leaving colourless solution Acidify with (dil)HNO3(aq) Add AgNO3(aq) Add (conc)NH3 to ppt Cream ppt AgBr forms Ppt dissolves leaving colourless solution Acidify with (dil)HNO3(aq) Add AgNO3(aq) Add (conc)NH3 to ppt Yellow ppt AgBr forms Ppt insoluble Add NaOH(aq) Add Devarda’s alloy (powdered Zn, Al) Heat & hold moist red litmus at mouth of test tube NH3 evolved, litmus paper red  blue Add Barium nitrate Ba(NO3)2(aq)/chloride BaCl2(aq) Add HCl(aq) White ppt Insoluble Add Barium nitrate Ba(NO3)2(aq)/chloride BaCl2(aq) Add HCl(aq) White ppt Dissolves Or add (dil)HCl(aq) Heat SO2 evolved turning potassium dichromate(VI) solution/paper from orange  green

Practical List 1. Make a salt and calculate the percentage yield (hydrated nickel sulfate) 2. make a salt and calculate the percentage yield (ammonium iron(II) sulfate) 3. carry out and interpret results of simple test tube reactions 4. measuring some enthalpy changes 5. finding the enthalpy of combustion of an alcohol 6. finding an enthalpy change that cannot be measured directly 7. reaction of alkanes 8. reaction of alkenes 9. experiment to find the effect of electrostatic force on jets of liquid 10. solubility of simple molecules in different solvents 11. thermal decomposition of group 2 nitrates and carbonates 12. flame tests on compounds of group 1 and 2 13. simple acid-base titrations 14. oxidation of metal and non-metallic elements and ions by halogens 15. disproportion reactions with cold and hot alkali 16. iodine/thiosulfate titration and the determination of purity of potassium iodate(V) 17. reactions between halogens and halide ions/some reactions of the halides 18. factors that influence the rate of chemical reactions 19. effect of temperature, pressure and concentrations on equilibrium 20. reactions of alcohols 21. preparation of organic liquid (reflux and distillation) 22. preparation of a halogenoalkane from an alcohol 23. reactions of the halogenoalkanes.

GCE Chemistry

guide

Edexcel Advanced Subsidiary GCE in Chemistry (8CH01) Edexcel Advanced GCE in Chemistry (9CH01) Internal Assessment of Practical Skills Issue 2 May 2008

Edexcel, a Pearson company, is the UK’s largest awarding body, offering academic and vocational qualifications and testing to more than 25,000 schools, colleges, employers and other places of learning in the UK and in over 100 countries worldwide. Qualifications include GCSE, AS and A Level, NVQ and our BTEC suite of vocational qualifications from entry level to BTEC Higher National Diplomas, recognised by employers and higher education institutions worldwide. We deliver 9.4 million exam scripts each year, with more than 90% of exam papers marked onscreen annually. As part of Pearson, Edexcel continues to invest in cuttingedge technology that has revolutionised the examinations and assessment system. This includes the ability to provide detailed performance data to teachers and students which helps to raise attainment.

This specification is Issue 2. Key changes are sidelined. We will inform centres of any changes to this issue. The latest issue can be found on the Edexcel website: www.edexcel.org.uk

References to third party material made in this specification are made in good faith. Edexcel does not endorse, approve or accept responsibility for the content of materials, which may be subject to change, or any opinions expressed therein. (Material may include textbooks, journals, magazines and other publications and websites.) Authorised by Roger Beard Prepared by Sarah Harrison All the material in this publication is copyright © Edexcel Limited 2008

Contents

Introduction

1

1 — The scheme of assessment

3

1.1 Activities

3

1.2 Conditions under which assessments are to be carried out

3

1.3 Materials allowed when carrying out assessment activities

4

1.4 Health and safety

4

2 — Activity b: Qualitative observation

5

2.1 Recording observations

5

2.2 Making inferences from observations

6

3 — Activity c: Quantitative measurement

7

3.1 Accuracy

7

3.2 Errors

8

3.3 Calculations

8

3.4 Graphs

9

3.5 Assumed laboratory skills

9

4 — Activity d: Preparation

11

5 — Activity c+d: Multi-stage experiment

12

Reference section

13

A — Inorganic compounds and elements

13

1

Appearance

13

2

Flame tests

14

3

Heating

14

4

Recognition and identification of common gases

15

5

Action of dilute acids

15

6

Tests for oxidizing and reducing agents

16

7

Hydrogen peroxide solution

16

B — Precipitates

17

1

Sodium hydroxide solution

17

2

Ammonia solution

18

3

Barium chloride solution

18

4

Silver nitrate solution

19

5

Concentrated sulfuric acid

19

C — Organic compounds

20

1

Appearance

20

2

Solubility

20

3

Ignition

20

4

Chemical tests

21

D — Spectroscopy

23

1

Mass spectrometry

23

2

Infrared spectroscopy

23

3

Nuclear magnetic resonance (nmr)

23

Introduction This guide has been written for students to help them in preparing for the tasks that are part of the internal assessment of practical skills. Teachers may also find this information helpful when preparing students for the assessment activities. The material contained in the booklet does not extend the specification content, but aims to help students to succeed in the assessment activities by: •

explaining in more depth what is required in carrying out the activities, making observations and measurements with appropriate precision and recording these methodically

•

advising them how to interpret, explain, evaluate and communicate the results of the activities clearly and logically using the relevant chemical knowledge, understanding and appropriate specialist vocabulary.

guide (Internal Assessment of Practical Skills) — Edexcel AS/A GCE in Chemistry (8CH01/9CH01) — Issue 2 — May 2008 © Edexcel Limited 2008

1

2


1 — The scheme of assessment

1.1

Activities

At both AS and A2 Level the skills being assessed in the activities are: a

General practical competence (GPC)

(verification)

b

Qualitative observation

(14 marks)

c

Quantitative measurement

(14 marks)

d

Preparation

(12 marks)

The maximum mark available at both AS and A2 Level is 40. Students will have their practical skills assessed by carrying out the assessment activities on a number of occasions throughout the course. The assessment activities are set by Edexcel. For activity a, the teacher is required to confirm that students have completed a range of practicals over the whole year and developed their laboratory skills. Students must have carried out at least five practicals in class. The practicals that the students complete must cover the three areas of physical, organic and inorganic chemistry. These five practicals can be either core practicals, or suitable alternatives. The marks for activities b, c and d are awarded following the teacher’s application of the mark schemes that accompany each assessed activity. Students need only carry out one exercise for each of activities b, c and d. However they may carry out more than one. In this case, only the highest mark for each activity will count towards the final mark out of 50. At A2 there is the option of completing a multi-stage experiment at A2, which consists of activities c and d together in a longer practical, which is worth 26 marks. Teachers have the option of marking these activities or having them marked by Edexcel.

1.2

Conditions under which assessments are to be carried out

The practical assessment activities must be carried out under controlled conditions that guarantee that students produce individual work. This includes those activities that involve the processing of results. Students must not consult with each other during the activities. Some activities can be completed in a laboratory session of approximately one hour. When an activity has to be carried over to a following session students must not remove any materials, including results and instruction sheets, from the laboratory. Instead, these must be collected by the teacher and reissued at the beginning of the next session when the activity is to be completed.


3

1.3

Materials allowed when carrying out assessment activities

Each activity has a cover sheet and a student brief giving full instructions for carrying out the activity and the questions based on it. During the activity students must not refer to books and notes although the data booklet is required for some exercises. This guide must not be used as a reference when carrying out the internal assessment tasks.

1.4

Health and safety

Students must follow the health and safety rules which normally operate in their chemistry laboratories, including the following: •

eye protection must always be worn

•

laboratory coats must be worn when appropriate

•

plastic gloves must be worn when supplied for a particular exercise

•

all substances should be regarded as being potentially toxic and hazardous

•

HazChem labels (eg flammable) should be read and appropriate precautions (eg keep liquid away from flame) taken

•

all substances spilled on the skin should be rinsed off immediately

•

chemicals must never be tasted

•

gases and vapours should never be smelt unless the question instructs the students to do so, and then this should be done only with great care.

4


2 — Activity b: Qualitative observation Activity b will be assessed by tasks that include a number of tests to be carried out, usually on a test tube scale. The instructions for each test will include details on quantities to be used, whether heat is required and whether excess reagent should be added. The reagents used for the tests will be limited to those included in the specification. Similarly, unknowns will be limited to compounds containing the ions, elements and organic compounds with the functional groups listed in the specification. As part of a task to identify an organic unknown, spectroscopic data may be included for students to analyse. Other physical data such as melting temperatures may also be given. These qualitative observation tasks will change each year.

2.1

Recording observations

After a student has carried out each test, as instructed in the task, they must communicate the results by giving a brief description of what has been observed. Possible changes that can occur during tests are listed below, along with examples of what students should write in the observation boxes. Possible change

Example of observation

a colour change in solution

yellow solution turns orange

the formation of a precipitate

white precipitate (ppte is allowed) is formed

a precipitate dissolves in excess reagent

precipitate dissolves in excess to form a green solution

a gas is evolved

bubbles of gas or effervescence

tests on a gas

the gas turned damp red litmus paper blue

flame tests

yellow flame

a solid dissolves

dissolves to give a blue solution

a reaction is exothermic

the mixture becomes hot

a reaction is endothermic

the mixture feels cold

a coloured solution loses its colour

the yellow solution turns colourless

There are a number of common mistakes which students make when recording their observations. The following should be avoided. •

Referring to ‘layers’ in test tubes. Almost certainly the presence of layers is due to inadequate mixing of the reagents. The exception to this is when an organic liquid is mixed with an aqueous solution, or with water, in which case ‘layers’ may be a valid observation.

•

Describing colours with elaborate adjectives such as brown-black or blue-green. Marks are awarded for simple descriptions of the colours of solutions and precipitates, such as black or blue, even if there is a trace of a second colour.

•

Stating that a gas is evolved without making an observation such as bubbles or effervescence.


5

•

Using clear instead of colourless for a solution. All solutions are clear even if coloured, eg copper sulfate solution is clear but coloured blue, whereas sodium sulfate solution is colourless.

2.2

Making inferences from observations

The purpose of asking students to make inferences from their observations is to test their knowledge, understanding and evaluation of the chemistry which leads to the observations. Activities include asking students to identify precipitates or gases formed in a test and which are recorded as an observation. Also, students could be asked to identify the unknown compound following a series of tests. In some cases the test may be enough to enable students to suggest the identity of a particular ion or functional group, but in others a number of possibilities may exist as a result of a single test. Example 1 The addition of aqueous ammonia to an inorganic compound produces a green precipitate. The inference from this test alone should be iron(II) hydroxide, chromium(III) hydroxide or nickel(II) hydroxide. Example 2 In an organic analysis, a compound produces an orange precipitate with 2,4−dinitrophenylhydrazine. The expected inference would be that the unknown compound is an aldehyde, ketone or carbonyl compound. If a following test shows that the compound is an aldehyde then the inference still stands, as this was valid on the basis of the 2,4−dinitrophenylhydrazine test. If more is known about the compound, for example its molecular formula and the fact that it cannot be oxidised, then it may be that the compound can be identified as a result of the 2,4−dinitrophenylhydrazine test.

6


3 — Activity c: Quantitative measurement Activity c will be assessed through tasks that involve using apparatus to make measurements and to process the results, to draw conclusions and evaluate the procedure. Quantitative activities could involve: •

a volumetric analysis task

•

a simple thermochemistry task

•

a task to follow the rate of a reaction.

3.1

Accuracy

Unless an activity instructs students differently, they should assume that readings from equipment and apparatus should be made with the following precision. Apparatus

Precision of reading

pipette

one volume only of 25.0 cm3

burette

each volume to the nearest 0.05 cm3 mean titre to 0.05 cm3 or to the second decimal place

measuring cylinder

a 100 cm3 size to 5 cm3, a 10 cm3 size to 1 cm3

balance

readings should be made to 0.01 g or 0.001 g depending on the precision of the balance

timers

normally read to the nearest second.

thermometers

0 to 100°C thermometers should be read to the nearest 1.0°C 0 to 50°C thermometers to an accuracy of at least 0.5°C

The following points are important. •

When students record readings they should include the appropriate number of decimal places. For example a burette reading of exactly 24.7 cm3 should be recorded in a results table as 24.70 cm3.

•

When titres have to be averaged, the mean should be expressed to either the nearest 0.05 cm3 or to the second decimal place. Eg if a student records four titres as listed below the mean should be calculated as:

•

26.50 + 26.25 + 26.60 + 26.65 = 26.5 cm3 4 If the student decides to ignore the second titre and to average the remaining three: 26.50 + 26.60 + 26.65 = 26.583 3 which should be recorded as 26.60 cm3 or as 26.58 (to the second decimal place).


7

•

Students should try to obtain at least two titres within 0.20 cm3, or better, of each other and average these to obtain a mean titre. Students should make it clear which titres have been used to obtain a mean.

•

A final result should only be given to the same number of significant figures as is suggested in the exercise. For example a task to find a ΔH value may involve a weighing of 1.56 g and a temperature rise of 7.5°C. A student who calculates a value of ΔH from these figures may obtain a value of 195.6843 on a calculator but this should be finally recorded as 200 kJ mol−1, although 196 kJ mol−1 may also be acceptable.

•

Units should always be included with a quantitative result.

A significant proportion of the marks awarded for a quantitative measurement activity will be for accuracy. These marks will be awarded by comparing the student’s results with an expected value.

3.2

Errors

Students should appreciate that any piece of equipment (burette, pipette, thermometer, balance) used in a quantitative exercise has an uncertainty associated with its use. Even if the equipment is used carefully, the uncertainty leads to an error in the reading and in the final result. Eg a balance has an uncertainty of 0.01 g when read to the second decimal place. A reading of 2.64 g recorded in an experiment has an error of:

0.01 × 100% = 0.38% 2.64 The following points will apply.

the uncertainty in the equipment × 100% the reading

1

The error in the reading =

2

Students should:

•

calculate the error involved in using a particular piece of equipment and state what effect this has on the overall accuracy of the activity

•

understand that the percentage error would be affected by the magnitude of the quantity being measured. There will be a greater error in weighing a mass of 2.64 g than in weighing 8.64 g using the same balance.

3

Only a simple treatment of errors is needed and students will not be asked to combine errors.

3.3

Calculations

Usually calculations will be structured. Students will be taken/guided through a series of steps leading to a final answer. Since most of the marks for these steps will be for use of a correct method, rather than for the numerical answer, it is important that students include their workings even if these seem to be trivial. Marks cannot be awarded for an incorrect answer without workings, but a correct method followed by an incorrect answer can often receive credit. Units, if appropriate, should always be included with a quantitative result. 8


3.4

Graphs

For some activities students will need to treat their readings graphically. The following are some useful points for students to consider when drawing graphs. •

Put the dependent variable, the quantity being measured (eg temperature), on the y−axis. And the pre-determined quantity (eg volume of solution) on the x−axis.

•

Choose the scales so that the results are spread out as far apart as the size of the grid allows, but this should not be at the expense of using a sensible scale, eg using 1 cm on the axis to represent 3 or 4 units might spread the readings better than using 1 cm to represent 5 units, but the scale would be hard to read.

•

The origin (0,0) does not need to be included on either scale if it is not relevant,

•

For example if temperature readings between 21.0°C and 39.0°C are to be plotted there is no need to begin the y−axis at 0. Rather it could be scaled from 20.0°C to 40.0°C.

•

Clearly label the axes with the quantity being plotted (eg time) and its units (eg minutes).

•

the points plotted with a continuous straight line or smooth curve. Since the readings are all subject to experimental error the line drawn may not necessarily through every point. Points should never be ed by a series of short, straight lines.

3.5

Assumed laboratory skills

The instructions for carrying out a quantitative measurement activity will include the essential points for the particular task. For example: •

which chemicals to use

•

the quantities needed

•

the sequence of steps in the method

•

the readings to be taken.

The assessment activities will assume that students have developed a range of routine practical skills in their course leading up to the exercise. Instructions may not, therefore, include every step needed to gain accurate readings. Examples of normal laboratory procedures which may not be referred to in the instructions include: •

burettes, pipettes and measuring cylinders should be rinsed with the solution they are to contain

•

a pipette filler should always be used with a pipette

•

conical flasks and volumetric flasks should be rinsed out with distilled water

•

a burette should be read at eye level

•

a thermometer bulb must be held in the centre of a solution when temperature readings are taken

•

care should be taken not to lose drops of a solution when thermometers or stirring rods are removed from it


9

•

containers may need to be labelled if this has not already been done

•

care should be taken as to where apparatus is placed on the bench, eg temperature measurements should not be taken in apparatus standing next to a Bunsen burner, in a patch of sunlight or in a strong draught.

10


4 — Activity d: Preparation Activity d is assessed through tasks in which students follow a procedure to prepare and, in some cases, purify an inorganic or organic compound. Marks are awarded for the student's ability to follow laboratory procedures and to use apparatus competently and safely. Preparations can include some of the following laboratory procedures: •

distillation

•

heating under reflux

•

filtration, including under reduced pressure

•

purification by washing in a separating funnel

•

solvent extraction

•

drying

•

boiling and melting temperature determination

•

crystallization and recrystallization.

Following some preparations students will be asked to calculate the maximum mass of product and a percentage yield.


11

5 — Activity c+d: Multi-stage experiment Instead of completing activities c and d separately, students could complete a multi-stage experiment, which combines the two. This would be a longer practical which covers all of the same aspects of c and d mentioned above. This would most likely be completed in two laboratory sessions and students would have to hand in their notes and results at the end on one session, and receive these back when they complete the practical in the second session.

12


Reference section

A — Inorganic compounds and elements 1 Appearance •

•

Coloured crystalline substances are usually hydrated salts of transition metals. The colour of the solid substance may provide evidence for its identity. Colour

Possible transition metal ions in salt

blue

copper(II)

pale green

iron(II)

green

copper(II), chromium(III), nickel(II)

purple

manganate(VII)

brown

iron(III)

pink

manganese(II)

yellow

chromate(VI)

orange

dichromate(VI)

The colours of transition metal ions in dilute, aqueous solution are shown in the table below. Colour

Possible identity

blue

copper(II)

green

iron(II), chromium(III), nickel(II)

brown / yellow

iron(III)

pale pink

manganese(II)

yellow

chromate(VI)

orange

dichromate(VI)

purple

manganate(VII)

colourless

zinc(II)


13

2 Flame tests •

*

To carry out a flame test, a clean nichrome wire is used to mix a sample of a solid with one drop of concentrated hydrochloric acid. The wire is held in progressively hotter parts of a non-luminous Bunsen flame. Flame colour

Inference

yellow

sodium ion

lilac

potassium ion

yellow-red*

calcium ion

red*

lithium or strontium ion

pale green

barium ion

Further tests would be needed to distinguish these ions.

3 Heating •

14

Gases or vapours may be evolved on heating a solid compound. Gas or vapour

Possible source

carbon dioxide

carbonates of metals other than group 1

oxygen

group 1 nitrates (other than Li)

oxygen and nitrogen dioxide

nitrates (other than Na or K)

water

hydrated salts


4 Recognition and identification of common gases

*

Gas

Observations

oxygen

colourless gas which relights a glowing splint

carbon dioxide

colourless gas which gives a white precipitate with limewater (calcium hydroxide solution)

ammonia

colourless gas which turns moist red litmus paper blue and forms white smoke with hydrogen chloride

nitrogen dioxide

brown gas*

hydrogen

colourless gas which ignites with a ‘pop’

hydrogen chloride

steamy fumes on exposure to moist air, acidic and forms white smoke with ammonia

chlorine

pale green gas which bleaches moist litmus paper

bromine

brown gas*

iodine

purple vapour

water vapour

turns blue cobalt chloride paper pink

Bromine dissolves in organic solvents to form a brown solution whereas nitrogen dioxide is insoluble.

5 Action of dilute acids •

When dilute sulfuric or hydrochloric acid is added to a substance a gas may be evolved or there may be a colour change in the solution. Action of acid

Possible source

carbon dioxide evolved

carbonate

hydrogen evolved

a metal

yellow solution turns orange

chromate(VI) to dichromate(VI)

sulfur dioxide evolved and pale yellow precipitate formed

thiosulfate


15

6 Tests for oxidizing and reducing agents Reducing agents usually: decolourise aqueous acidified potassium manganate(VII) and may also turn aqueous, acidified potassium dichromate(VI) from orange to green. Reducing agents include: iron(II) ions iodide ions hydrogen peroxide. Oxidizing agents usually: liberate iodine as a brown solution or black solid from aqueous potassium iodide. Iodine solution gives a blue-black coloration with starch. Oxidizing agents include: acidified manganate(VII) ions acidified dichromate(VI) ions hydrogen peroxide copper(II)ions aqueous chlorine aqueous bromine.

7 Hydrogen peroxide solution Aqueous hydrogen peroxide (H2O2) can act as both an oxidizing and a reducing agent often with the evolution of oxygen, although this may be unreliable.

16

Observation on adding H2O2

Inference

brown precipitate

manganate(VII), brown precipitate is MnO2

purple solution is decolourised

manganate(VII) in acid solution

pale green solution turns yellow

iron(II) to iron(III) in acid solution

green precipitate turns brown

iron(II) hydroxide to iron(III) hydroxide

green alkaline solution goes yellow

chromium(III) to chromate(VI)

brown solution or black precipitate

iodine from iodide in acid solution


B — Precipitates When two aqueous solutions are mixed together and an insoluble compound is formed this is known as a precipitate not a suspension. The observation that a precipitate is formed should always be accompanied by the colour of the precipitate, even if this is white. Some reagents should be added until they are in excess. This may result in a precipitate forming then dissolving in excess reagent.

1 Sodium hydroxide solution •

When dilute sodium hydroxide (NaOH) solution is added to a solution containing a metal ion a precipitate of the insoluble hydroxide, eg Mn(OH)2, is usually formed. Precipitates which are amphoteric hydroxides will dissolve in excess sodium hydroxide to give a solution containing a complex ion, eg [Cr(OH)6]3−.

•

Students should assume that aqueous sodium hydroxide should be added until it is in excess even if this is not explicitly stated in the instructions.

Observation on adding dilute NaOH

Observation on adding excess dilute NaOH

Likely ion

green precipitate

precipitate dissolves to a green solution

chromium(III)

off-white precipitate which darkens on exposure to air

precipitate is insoluble

manganese(II)

green precipitate which turns brown on exposure to air


iron(II)

brown precipitate


iron(III)

green precipitate


nickel(II)

blue precipitate


copper(II)

white precipitate

precipitate dissolves to a colourless solution

zinc(II)

white precipitate


magnesium, barium, strontium, calcium

no precipitate

sodium, potassium

no precipitate but ammonia evolved on warming

ammonium


17

2 Ammonia solution •

Dilute aqueous ammonia (NH3), when added to a solution containing a cation, will form the same hydroxide precipitate as dilute sodium hydroxide solution, eg Mn(OH)2. Excess aqueous ammonia may dissolve the precipitate to form a complex ion, eg [Cu(NH3)4(H2O)2]2+.

•

Students should assume that aqueous ammonia must be added until it is in excess.

Observation on adding dilute aqueous NH3

Observation on adding excess dilute aqueous NH3

Likely ion

green precipitate

precipitate is soluble to give green solution

chromium(III)

off-white precipitate


manganese(II)

green precipitate turning brown

precipitate insoluble

iron(II)

brown precipitate

precipitate insoluble

iron(III)

green precipitate

precipitate dissolves to give blue solution

nickel(II)

blue precipitate

precipitate dissolves to give deep blue solution

copper(II)

white precipitate

precipitate dissolves to give colourless solution

zinc(II)

white precipitate


magnesium

3 Barium chloride solution •

Aqueous barium chloride forms precipitates of insoluble barium salts with a number of anions but is usually used as the test for the sulfate, SO42−, ion. Aqueous barium chloride is usually used with dilute hydrochloric acid. Anion

Precipitate

Addition of dilute HCl

colour

formula

sulfate

white

BaSO4


sulfite

white

BaSO3

precipitate dissolves

carbonate

white

BaCO3

precipitate dissolves with effervescence

If dilute hydrochloric acid is added to the anion solution before aqueous barium chloride then only the sulfate will form as a precipitate.

18


4 Silver nitrate solution •

Aqueous silver nitrate is commonly used to test for the presence of halide ions in solution. Anions which would interfere with the test (eg carbonate) are removed by adding dilute nitric acid before the aqueous silver nitrate.

•

The identity of a halide may be confirmed by the addition of aqueous ammonia, (NH3), both dilute and concentrated.

•

Silver halides which dissolve in ammonia do so to form a colourless solution of the complex ion, [Ag(NH3)2]+. Anion

Precipitate

Addition of aqueous NH3

colour

formula

dilute

concentrated

chloride

white

AgCl

soluble

bromide

cream

AgBr

soluble in excess

soluble

iodide

yellow

AgI

insoluble

insoluble

5 Concentrated sulfuric acid •

When a few drops of concentrated sulfuric acid (H2SO4) are added to a solid halide the observed reaction products may be used to identify the particular halide ion present. This is a potentially hazardous reaction.

•

It must be carried out on a small scale and in a fume cupboard.

•

The products in brackets will not be observed since they are colourless gases. The halide ion may be identified without the need to test for these gases. No attempt should ever be made to detect these gases by smell. Halide

Observations on adding concentrated H2SO4

Observed reaction products

chloride

steamy fumes, vigorous reaction

HCl

bromide

steamy fumes, brown vapour, vigorous reaction

HBr, Br2 (SO2)

iodide

steamy fumes, black solid, purple vapour, yellow solid, vigorous reaction

HI, I2, S, (H2S)


19

C — Organic compounds Students will always be told if a compound, or mixture of compounds, to be identified is organic. Often the molecular formula, or the number of carbon atoms in a molecule, of a compound will be given. Chemical tests may be followed by spectroscopic information.

1 Appearance Simple organic compounds are usually colourless liquids or white solids. It is unlikely that appearance alone will provide firm evidence for identification.

2 Solubility Solubility of compound

Possible identity

pH of solution

Possible identity

dissolve in water

simple alcohols, simple carboxylic acids, propanone, simple aldehydes, simple amines and their salts

above 7

amines

below 7

carboxylic acids, phenols

dissolve in dilute acid but may not dissolve in water

amines

dissolve in aqueous alkali but may not dissolve in water

carboxylic acids, phenols

3 Ignition Igniting an organic unknown on a crucible lid may help in identifying it.

20

Observation

Possible inferences

burns with a smoky flame

aromatic, unsaturated eg alkene

burns with a clean flame

saturated low molar mass compound

no residue

most lower molar mass compounds


4 Chemical tests The details of how these tests are to be carried out will be included in the instructions to students in the assessment activities. Test

Observation

Inference

warm with acidified potassium dichromate(VI)

orange to green solution

primary or secondary alcohol, aldehyde

yellow solution is decolorised

alkene

if white precipitate also formed

phenol

precipitate:

halogenoalkanes:

white

C−Cl

cream

C−Br

yellow

C−I

phosphorus(V) chloride

steamy fumes of HCl that turn damp blue litmus paper red

OH group in alcohols and carboxylic acids

2,4−dinitrophenylhydrazine solution

orange precipitate

C=O group in aldehydes and ketones

boil with Fehling’s or Benedict’s solution

blue solution gives red precipitate

aldehyde

warm with ammoniacal silver nitrate

silver mirror

aldehyde

sodium or potassium carbonate or hydrogencarbonate solution

effervescence

carboxylic acid

add a small piece of sodium

effervescence (bubbles), sodium dissolves, white solid formed

alcohol, phenol or carboxylic acid

warm with carboxylic acid and a few drops of concentrated sulfuric acid

ester smell, eg glue-like

alcohol

shake with bromine water

warm with aqueous sodium/potassium hydroxide, acidify with dilute nitric acid then add aqueous silver nitrate

(Tollens’ reagent)


21

Test

Observation

Inference

sodium nitrite and dilute hydrochloric acid followed by an alkaline solution of phenol in ice-cold conditions

orange precipitate

aromatic amine

iodine in alkaline solution

pale yellow precipitate

methyl ketone or ethanal,

C CH3 O methyl secondary alcohol or ethanol

CH CH3 OH

22


D — Spectroscopy Modern instrumentation plays a vital role in the characterisation and identification of molecules and ions. The sections which follow will enable interpretation of the data likely to be presented to students.

1 Mass spectrometry Simplified mass spectra will be given which may be interpreted in two main ways. 1

The value of the compound’s relative molar mass may be obtained from the m/e value of M+, the molecular ion. This will have the highest value of m/e. This need not necessarily be the ‘base’ peak, which is simply the most abundant ion. Questions will be set in such a way that students will not be confused by the presence of a line due to the (M + 1) ion.

2

The fragmentation pattern of the spectrum gives useful information about the structure of the molecule. For example, a peak at m/e 29 is likely to be due to the presence of a C2H5 group in the molecule. Students are reminded that, when asked, they should give displayed structures for fragments, which must carry a positive charge also.

2 Infrared spectroscopy This is a very powerful non-destructive technique which provides information regarding the nature of covalent bonds within the molecule. Students should look at the most intense absorptions to quickly gain structural clues. Table 1 provides sufficient details to enable the principal bands to be assigned. Students should that absorption frequency is affected by the chemical environment and that absorption may take place outside the range given. The connection between structures should be recognised. For instance, an alcohol [O-H] stretch will be accompanied by a [C-O] stretch. Please see the Data booklet for specific IR spectroscopy data.

3 Nuclear magnetic resonance (nmr) Nuclear magnetic resonance spectra may be included as part of an investigation into structure. Spectra will be high resolution with possible spin–spin coupling displayed. Please see the Data booklet for specific nmr data.

1714sb130508S:\LT\PD\\GCE in Chemistry Intrnl Asst of Practical Skills guide Issue 2.doc.1-28/1


23

May 2008 For more information on Edexcel and BTEC qualifications please visit our website: www.edexcel.org.uk Edexcel Limited. ed in England and Wales No. 4496750 ed Office: One90 High Holborn, London WC1V 7BH. VAT Reg No 780 0898 07

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