School Science Lessons
Chemistry
Updated: 2009-10-10
Please send comments to: J.Elfick@uq.edu.au

Table of contents
3.1.0 Bunsen burner
3.2.0 Identify pure substances
3.9.0 Solubility and solutions
3.18.0 Separate substances from mixtures
3.28.0 Substances that gain mass when heated, copper foil
3.30.0 Substances may decompose and lose mass when heated, thermal decomposition
3.31.0 Hygroscopic, deliquescent and efflorescent chemicals

3.32.0 Prepare, collect and test gases

3.52.0 Rusting
5.3.4 Acid-base indicators
3.54.0 Crystal growth
3.55.0 Matter as particles
3.1.0 Conductivity
3.61.0 Construction materials

3.70.0 Chemical reactions
3.80.0 Energy from chemical reactions
3.84.0 Electrical energy from chemical reactions
3.91.0 Rate of reaction
3.95.0 Breakdown large molecules to small molecules
3.100.0 Building up molecules

3.1.0 Bunsen burner
3.1.1 Bunsen burner
3.1.1.1 Bunsen burner gases
3.1.2 Lighting a Bunsen burner
3.1.3 Study the Bunsen burner flame
3.1.4 Bunsen burner flame and candle flame

3.2.0 Identify pure substances
3.2 Melting point of naphthalene
3.3 Melting point of naphthalene with a capillary tube
3.4 Impurities affect the melting point of a substance
3.5 Boiling point of water
3.5.1 Boiling point of sodium chloride solution
3.6 Boiling point of inflammable liquids
3.7 Volatility of different liquids
3.8 Pressure affects the boiling point

3.9.0 Solubility and solutions
3.9 Solubility in water
3.10 Solubility and temperature, solubility of salts in water
3.11 Solubility of a substance in water at a given temperature
3.12 Solubility and particle size
3.13 Solubility and solvents
3.14 Solubility and agitation
3.15 Volume of solutions
3.16 Miscible liquids
3.17 Heat of solution
3.17.1 "Magnetic" sugar cube dissolves

3.18.0 Separate substances from mixtures
3.18 Separate tin from a mixture of tin and carbon (charcoal)
3.19 Separate by sublimation, iodine
3.20 Separate by distillation
3.21 Separate by fractional distillation of crude oil
3.22 Separate salt and sand
3.23 Solvent extraction of oil from nuts
3.24 Separate pigments from green leaves with paper chromatography
3.24.1 Separate mixed inks with paper chromatography
3.25 Gases dissolved in a water sample
3.26 Separate immiscible liquids of different density
3.27 Separate solids using density differences

3.28.0 Substances that gain mass when heated, copper foil
3.28.1 Substances that gain mass when heated, prepare magnesium oxide
3.28.2 Substances that gain mass when heated, prepare lithium oxide
3.28.3 Substances that gain mass when heated, prepare calcium oxide
3.28.4 Collect and weigh the gaseous products of a burning candle

3.30.0 Substances may decompose and lose mass when heated, thermal decomposition
3.30.1 Carbonates
3.30.2 Hydrogen carbonates (hydrogencarbonates, bicarbonates)
3.30.3 Hydroxides
3.30.4 Nitrates
3.30.5 Oxides
3.30.6 Sulfates
3.30.7 Sulfites
3.30.8 Salts with water of crystallization, hydrous salts
3.30.9 Boric acid, H3BO3
3.30.10 Oxalic acid
3.30.11 Potassium chlorate, KClO3
3.30.12 Sublimation, iodine, camphor, naphthalene
3.30.13 Silica sand
3.30.14 Ammonium dichromate
3.30.15 Ammonium chloride

3.31.0 Hygroscopic, deliquescent and efflorescent chemicals
3.31.1 Expose different salts to the air
3.31.2 Expose sodium carbonate decahydrate, washing soda, to the air
3.31.3 Tests for water with cobalt (II) chloride

3.1.1 Bunsen burner
See diagram 3.1: Bunsen burner flame | See diagram 3.1.1: Bunsen burners
Be careful! Do not turn the gas on without lighting the Bunsen burner. Gas forms an explosive mixture in air.
1. Combustion is the burning in oxygen gas of a substance to produce heat and sometimes light energy. A flame appears during combustion when a gas has such a high temperature that it emits heat and light. A flame appears only where the burning gas and oxygen gas are in contact.
2. The Bunsen burner consists of 2.1. a barrel, shaft, 2.2 an air regulator, i.e. a sleeve with a hole in it, 2.3 a jet, air mixture valve, needle valve, 2.4 a base, 2.5. a gas inlet opening.
3. Adjust the flame by opening or closing the gas tap. When the air regulator is open, the gas burns with a noisy blue flame that may be nearly invisible in strong light. If the flame rises up from the burner, turn down the gas supply.
4. When not using the Bunsen burner, either turn off the gas or close the air regulator to give a safety flame. The flame is yellow because of the incandescence of carbon particles. It is not as hot as the blue flame and leaves black soot deposits on glassware.
5. Regularly inspect gas fittings on the benches and hoses connecting Bunsen burners to gas turrets to make sure that connections are free of leaks.
Tests for leaks by dipping the part in soapy water. Be careful! Do not use a lighted match.
6. Heat flammable liquids in water baths using electrical hot plates, not Bunsen burners. Turn the gas off first at the gas tap, then at the cylinder or main supply tap.
7. Use the Bunsen burner only in a draught free area. Allow the Bunsen burner to cool before you move or store it.
8. Do not heat low melting point objects, e.g. plastics, solder, lead, over the barrel of the burner. Melted pieces may fall inside the barrel. Hold the burner at an angle. If a match is blown out, turn gas off, then light the Bunsen burner again.

3.1.1.1 Bunsen burner gases
See 16.1.1cc: LPG (liquefied petroleum gas, LP gas)
See 16.1.1a.1: Natural gas
Town gas was previously used on school laboratories, but nowadays the gas is usually natural gas, i.e. mostly methane CH4, or LPG, bottled gas, mostly propane, C3H8. In a laboratory, the pilot light should burn with a 90% blue flame. If the flame is yellow, the gas may be contaminated with condensates. Do not use such a gas but immediately inform the local gas authority.
The heating values of fuels: town gas 88 MJ / kg, natural gas 55.6 MJ / kg, LPG gas 49 MJ / kg, diesel fuel oil 38 MJ / L, kerosene 36.7 MJ / L, coke or coal 27 MJ / kg, dry split wood 12.5 MJ / kg.
CH4 (g) + 2O2 (g) --> CO2 (g) + 2H2O (g) + heat

3.1.2 Lighting a Bunsen burner
See diagram 3.1.2: Right and wrong ways to use a Bunsen burner
1. Close the air regulator, light a match, hold the match flame at the side of the barrel opening, turn the gas tap on, raise the match flame to light the gas. The gas burns with a visible yellow flame, a quiet safety flame. Hold a test-tube just above the flame. Note the carbon (soot, carbon black) that deposits on the glass. To test whether unburned carbon gives the yellow colour to the flame, sprinkle powdered charcoal on the flame and compare the yellow colours.
2. Start to open the air regulator until the gas burns with a medium blue flame with a light blue inner cone and a pale violet outer flame with a bushy appearance. The flame has an outer oxidizing zone where combustion is complete, a middle reducing zone, and an inner unburned gases zone surrounded by a blue cone. This flame is the most useful for heating. Fully open the gas regulator until you get a roaring blue flame with a light blue triangle in the centre of the blue cone.
3. Open the air regulator. Keep turning down the gas supply. The gas "blows back", "strikes back". The gas is burning inside the barrel. Turn the gas fully on and strike the gas supply rubber tube with a sharp blow from the side of your hand. If the flame does not reappear, immediately turn the gas off and leave to cool because the barrel may be hot. Then light the Bunsen burner again.

3.1.3 Study the Bunsen burner flame
See diagram 3.1.3: Burning the gas in a cone of flame
1. Hold the end of a glass tube in the centre of the cone. You can light the gas coming out of the other end of the glass tube.
2. Hold a piece of wire in different parts of each kind of flame, moving it from the bottom to the top. Find the hottest flame and the hottest place in each flame with a piece of nichrome wire or iron wire stuck into a cork for a handle. The approximate temperatures and colours for the wire are as follows: 1. <500oC, wire gives no light, flame is non-luminous 2. 500oC to 950oC, wire becomes red, then dark red, then bright red (red-hot) 3. 950oC to 1350oC, wire becomes yellow-red then becomes white 4. >1350oC, wire becomes white (white-hot). The safety flame has a similar temperature in different parts about 300oC. It is never used for heating. The medium blue flame has the hottest point at the tip of the blue cone at about 500oC. The roaring blue flame has the hottest point at the tip of the cone at about 700oC.
3. Close the air regulator. Use a wood splint or a taper to test that parts of the flame support ignition. The wood splint match is set alight in all positions in the yellow flame where no air mixes with the gas. Repeat the experiment with the air regulator open. A cone of mixed air and gas exists in the centre of the cone where the gas is not burning.
4. Turn off the gas. Push a pin at right angles through a match just below the chemical on the end of the match. Use the pin to hang the match in the barrel with the chemical end just above the rim. Open the air regulator and light the gas again. The match does not ignite inside the cone. Move the match to the outer cone of the blue flame. The match ignites.
5. Close the air regulator and light the gas. Hold a piece of copper wire gauze with tongs 3 cm above the top of the barrel. Hold a lighted match above the gauze. The gas ignites above the gauze. Lower the gauze until the flame passes through it. Repeat the experiment with an open air regulator. Light the gas and lower a copper wire gauze down on the flame. The flame remains below the wire gauze as the gauze becomes red-hot. Heat is removed from the gas air mixture by the copper gauze.

3.1.4 Bunsen burner flame and candle flame
See diagram 3.1.4: Bunsen burner flame and candle flame
Repeat the above experiments with a candle and a spirit burner. Just above the wick of a burning candle is a dark region of unburned gas. Above and around it is a yellow region containing incandescent particles of carbon undergoing combustion to form carbon dioxide. Put the candle flame under an evaporating basin. Note the deposits of carbon, soot, because of insufficient oxygen gas to complete combustion.

3.2 Melting point of naphthalene
See diagram 3.2: An approximate melting point
Put 2 cm of naphthalene flakes in a test-tube. Hold a thermometer with its bulb in the naphthalene. Use a small flame to heat the test-tube gently and watch the thermometer reading. To find the melting range, note the temperature when the naphthalene melts. Leave to cool and note the temperature when the naphthalene solidifies. To find the melting point, calculate the average of these two values. The melting point of pure naphthalene is 80.5oC.

3.3 Melting points of naphthalene with a capillary tube
See diagram 3.3: More accurate way of finding the melting point
Make a capillary tube by drawing out a glass tube over a hot flame. Put a very small amount of naphthalene in a capillary tube sealed at one end. Attach a thermometer to the capillary tube, a sealed end down. Put the thermometer and capillary tube in a container of water and slowly heat the water while stirring with the thermometer and capillary tube. Do not let water enter the capillary tube. To find the melting range, note the temperature when the naphthalene melts, leave to cool, and note the temperature when the naphthalene solidifies. To find the melting point, calculate the average of these two values.
Repeat the experiment with stearic acid, m.p. 69oC, palmitic acid, m.p. 63oC, butter, soap, 1,4-dichlorobenzene (deodorizer) m.p. 53oC, paraffin wax, m.p. 45oC - 65oC, sodium thiosulfate pentahydrate 48.3oC.

3.4 Impurities affect the melting point of a substance
Mix stearic acid with the naphthalene to make the naphthalene impure. Note changes in the melting point. Impurities lower the melting point.

3.5 Boiling point of water
See diagram 3.5: Boiling point of water
1. Pour water into a test-tube. Hold a thermometer with its bulb just under the water. Heat very slowly by moving the test-tube in and out of the flame or add boiling chips, anti-bumping granules. Heat the water gently until it boils. Record the temperature. Note the same temperature in all parts of the test-tube. Note any change in the reading if the thermometer touches the bottom of the test-tube. The water must cover the bulb of the thermometer and the bulb must not touch the sides of the test-tube.
2. Show that the boiling point of water does not depend on the size of the container. Repeat the experiment with a large container. Heat the water quickly. The water first starts to boil near the bottom and sides of the container. Note the temperature in different parts of the container. Note any change in the reading if the thermometer touches the bottom of the container. The boiling point is the same in small and large containers.

3.5.1 Boiling point of sodium chloride solution
1. A solution of sodium chloride in water boils at a higher temperature and has a lower freezing point than pure water. Use freezing points and boiling points to find the purity of substances. Use three test-tubes containing the same volume of water. Add some sodium chloride to the second test-tube. Keep adding sodium chloride to the third test-tube until no more dissolves to produce a saturated solution at that temperature. Join the test-tubes with an elastic band. Heat the test-tubes equally over a Bunsen burner. The first test-tube containing only water boils first. The second test-tube containing some sodium chloride boils next. The third test-tube containing the saturated solution of sodium chloride boils last.
2. Put a beaker containing demineralized water in a broad pan containing a concentrated salt solution. Slowly heat the broad pan and note that the demineralized water boils first.

3.6 Boiling point of inflammable liquids
See diagram 3.6: Boiling point of inflammable liquids
1. Do not use a Bunsen burner to find the boiling point of inflammable liquids, e.g. ethanol, b.p. 78.4oC and acetone, b.p. 56oC. Use an electric hot plate or use the following method. Pour 2 cm of the inflammable liquid into a test-tube in an empty container. Place a thermometer in the test-tube with its bulb in the liquid. Boil water in an electric jug or on an electrical hot plate. Pour the hot water into the container so that the level is higher than the inflammable liquid in the test-tube. Stir the inflammable liquid gently with the thermometer and read thermometer when the inflammable liquid boils. [It is not good practice to stir liquids with thermometers!]
2. Use a very small test-tube or seal one end of a piece of glass tubing, 8 cm length and 3 cm external diameter. Put the inflammable liquid into this test-tube. Put a capillary tube, sealed at one end, into the inflammable liquid with the sealed end up and the open end down in the inflammable liquid. Use a rubber band to attach the test-tube containing inflammable and capillary tube to the bulb of a thermometer. Hold the apparatus in a container of water and heat gently with an electric hot plate. When the temperature rises, bubbles slowly come out of the capillary tube. At the boiling point the bubbles suddenly come out as a steady stream. Read the temperature. Let the water cool and read the temperature again when the steady stream of bubbles stops. Calculate the boiling point as the average of the two readings.

3.7 Volatility of different liquids
Evaporation is the movement of particles from the surface of a liquid to the gas state, when below the boiling point. Volatile liquids evaporate readily at room temperatures.
1. Select liquids from the laboratory, e.g. water, methylated spirit, gasoline, mineral turpentine, kerosene (paraffin oil) household machine oil, car oil, vinegar, vanilla essence, eucalyptus oil, glycerine. Wet a 5 cm piece of absorbent paper with a liquid, Write the name of the test liquid in pencil. Attach the piece of paper to a horizontal string. Examine the paper every ten minutes, every two hours and each day.
2. Repeat the experiment with perfumes. Smell the paper every ten minutes, every two hours and each day. Some perfumes soon disappear, but others last for days. Record the relative "person-attracting" capacity for each perfume.

3.8 Pressure affects the boiling point
See diagram 3.8: Decreasing the pressure on boiling water
1. Put water in a sidearm test-tube or in a round-bottom flask with a one-hole stopper. Insert a thermometer through a hole in the stopper so that the bulb of thermometer reaches, but does not touch, the bottom of the test-tube or flask. Add boiling chips to prevent bumping. Boil the water and read the temperature. Stop heating. Connect a water pump to the sidearm or to the second hole of the two-holes stopper. When the water stops boiling, turn on the water pump to reduce the pressure. Read the temperature, heat to boiling and read the temperature again.
2. Boil water on a high mountain and note the boiling point. People who climb Mount Everest complain that they cannot get a good cup of tea.

3.9 Solubility in water
In this document the word "solution" refers to substances dissolved in water, i.e. aqueous solutions. A solvent is a liquid that dissolves another substance, the solute, to form a solution. The three ways to increase the rate at which a solid dissolves in water are as follows: 1. grinding the solid until finely divided 2. shaking the solution while the solid is dissolving, and 3. heating the solution.
Try to dissolve 5 g of different salts each in 15 mL of water in a test-tube. Attach a stopper and shake vigorously. Solubility is a characteristic of a particular substance. Classify each salt as soluble or slightly soluble or insoluble. The solubility of a salt can be expressed as the number of grams able to dissolve in 100 g of water at 20oC, e.g. ammonium chloride 37.2 g, barium chloride 35.7 g, calcium chloride 42.7 g, copper (II) sulfate 20.7 g, lead nitrate 54.4 g, magnesium sulfate 25.2 g, potassium chloride 34.0 g, potassium iodide 144.0 g, sodium hydrogen carbonate (sodium bicarbonate) 9.6 g, sodium chloride 36.0 g, sodium hydroxide 109.0 g, sodium nitrate 87.5 g.

3.10 Solubility and temperature, solubility of salts in water
The solubility of a potassium dichromate in 100 g of water varies with temperatures: 0oC - 5 g, 10oC - 7 g, 20oC - 12 g, 30oC - 20 g, 40oC - 26 g, 50oC - 34 g, 60oC - 43 g, 70oC - 52 g, 80oC - 61 g, 90oC - 70 g, 100oC - 80 g.
1. Show that a saturated solution contains less dissolved solid at a lower temperature than at a higher temperature. Make a 50 mL saturated solution of potassium dichromate or potassium nitrate at 60oC. Pour the clear solution into a clean container and keep the temperature of this container at 40oC until crystals stop forming. Pour the clear solution from this container into another clean container. Do not pour crystals into the container. Leave to cool and note more crystals forming as the solution cools.
2. Repeat the experiment with sodium chloride. This salt is unusual because the solubility hardly changes with change in temperature.

3.11 Solubility of a substance in water at a given temperature
1. Add a teaspoonful of powdered milk to water with ice floating in it, water at room temperature, warm water, boiling water. Observe the time taken for the powdered milk to dissolve in the water  at different temperatures.
2. Add sodium hydrogen carbonate (sodium bicarbonate) to 100 g of water in a container while stirring. Keep adding sodium carbonate until no more solute will dissolve. Decant the clear saturated solution. Read the temperature of the saturated solution, i.e. room temperature. Weigh a clean evaporating dish, w1. Add some clear saturated solution and weigh again, w2. Carefully evaporate the solution in the evaporating dish to dryness and weigh again, w3. The mass of the dissolved sodium hydrogen carbonate = w3 - w1. The mass of water = w2 - w1 - w3. Calculate the solubility of the sodium hydrogen carbonate as weight in grams dissolved in water at room temperature (w3 - w1) / (w2 - w1 - w3).
Repeat the experiment using water at a higher temperature.
3. The solubility of sodium bicarbonate in 100 g of water varies with temperature: 0oC - 6.9 g, 10oC - 8.15 g, 20oC - 9.6 g, 25o- 10.35 g, 30oC - 11.1 g, 40oC - 12.7 g, 50oC - 14.45 g, 60oC -16.4 g.

3.12 Solubility and particle size
Show that small particles dissolve faster than large particles.
1. Add coarse table salt to a first test-tube half filled with water. Add the same quantity of fine table salt to a second test-tube that contains the same volume of water. Shake both test-tubes equally and simultaneously. Note the amount of undissolved table salt left in each test-tube.
2. Use two equal samples of large crystals of copper (II) sulfate. Grind one sample into a fine powder. Put both samples into equal quantities of water in separate test-tubes and shake. Compare the rates at that the different samples dissolve and cause the water to change colour.

3.13 Solubility and solvents
1. Fill two test-tubes one third full with water and another with methylated spirit. To each test-tube add 1 g sodium chloride, attach a stopper and shake. Sodium chloride dissolves readily in water, but not so readily in methylated spirit.
2. Add sodium chloride crystals to a dilute solution of sugar in water and note whether the crystals dissolve. Drop crystals of potassium dichromate into the solution. Note whether the solution changes colour. Colour change shows that potassium dichromate is also dissolving. The presence of one dissolved substance does not prevent other substances dissolving in the solution. Unless the concentrations are high, one solute does not affect the solubility of other solutes in the solution.

3.14 Solubility and agitation
Prepare two equal samples of cane sugar. Put one sample of cane sugar into a test-tube half full of water. Add a stopper and shake the test-tube until all the sugar dissolves. Put the other sample of cane sugar into a test-tube. Very slowly add the same volume of water as in the first test-tube. Leave to stand. Most of the sugar has not dissolved but, if left to stand for long enough, all the cane sugar will dissolve as in the first test-tube.

3.15 Volume of solutions
1. Fill a small, narrow-necked flask with water to a level in the neck and mark this level. Add sodium chloride to the water with continual shaking until the solution is saturated and no more dissolves. Note the new level of the liquid. The volume of the solution is only slightly greater than the original volume of the water.
2. Close one end of a glass delivery tube about 30 cm long. Fix it upright, half fill it with water and mark the level of the water. Slowly add alcohol to fill the delivery tube. The water and the alcohol fill equal lengths in the tube. Shake the tube thoroughly to mix the water and alcohol. The new level of the solution in the tube shows a slight decrease in total volume.

3.16 Miscible liquids
Put 10 mL of water in three test-tubes. Add 1 mL of: 1. methylated spirit 2. glycerine and 3. kerosene. Shake each test-tube. Miscible liquids can mix in all proportions. 1. Alcohol and water are miscible. 2. Glycerine and water are miscible. 3. Kerosene and water are not miscible, immiscible.

3.17 Heat of solution
Dissolve some equal quantities of sodium hydroxide, potassium hydroxide, ammonium chloride and urea in separate test-tubes half full of water. Feel the test-tubes and note any change in the temperature. Sodium hydroxide and potassium hydroxides dissolve in water with an increase in temperature. Ammonium chloride and urea absorb heat from their surroundings when dissolving in water.

3.17.1 "Magnetic" sugar cube dissolves
Fill a large dish with water Wait until the water is absolutely still then lower a matchstick into the centre of the water. Carefully dip a sugar cube in the water near the edge of the dish. The matchstick moves towards the dissolving sugar cube. When the sugar dissolves in the surface water, the surface water becomes heavier and falls to be replaced by surface water flowing towards the sugar cube, carrying the matchstick with it.

3.18 Separate tin from a mixture of tin and carbon (charcoal)
Get tin bits by cutting a tin welding rod to pieces because 66% of the tin welding rod is tin and the rest is lead. Do not use a "tin can" because it is mostly iron with a thin layer of tin on its surface!
1. Make a mixture of tin (tin filings or small cut pieces of tin) mp 232oC and carbon (crushed charcoal) mp 3,730oC. Mix the tin bits and charcoal bits uniformly. Heat the mixture in a crucible. Stir with a splint until the tin melts and forms a liquid below the charcoal. Pour the tin onto plaster of Paris mold or other heat-proof surface. While pouring, hold back the charcoal in the crucible with a wood splint. Use melting point and melting point behaviour to identify a substance and decide if it is pure. Tin solder melts at 250oC. Carbon melts at 3,700oC.
2. Mix solder filings with powdered charcoal. Heat the mixture in a crucible. Stir with a splint until the solder melts and forms a liquid below the charcoal. Pour the liquid into a container by holding back the charcoal in the crucible.

3.19 Separate by sublimation
See diagram 3.19: Sublimation of iodine
Separate iodine from a mixture of crystals of iodine and sodium chloride. Heat the mixture in an evaporating dish with a funnel placed over it. The iodine sublimes on to the cool sides of the funnel.

3.20 Separate by distillation
See diagram 3.20.1: Distil ink | See diagram 3.20.2: Condensing the vapour
Put 10 mL of ink in a flat-bottomed conical flask. Add boiling chips to prevent bumping. Fit a stopper with a delivery tube reaching half way down a receiving test-tube or a U-tube, in a container of water. Heat the ink with a Bunsen burner flame. Drops of a colourless liquid appear in the receiving test-tube. Identify the liquid as water by its action of turning white anhydrous copper (II) sulfate to blue hydrated copper (II) sulfate. Do not allow ink to froth up or splash into the delivery tube.

3.21 Separate crude oil fractions by fractional distillation
See diagram 3.21: Collect fractions
1. Use crude oil or a substitute for crude oil, e.g. a mixture of used car oil, paraffin, thin lubricating oil, diesel oil and petroleum jelly. Use a hard-glass test-tube, or sidearm test-tube, fixed to a retort stand, a delivery tube and five small ignition tubes. Use a 0o to 360oC thermometer. Put 4 mL of crude oil in the test-tube. Add boiling chips to prevent bumping. Set up five small ignition tubes to collect the fractions. Heat the oil very gently. Collect 10 drops of distillate in the first ignition tube, then collect 10 drops of distillate successively in the other ignition tubes. The boiling point of the remaining oil will become higher as distillation proceeds and oil will then require more heat from the Bunsen burner. Arrange the fractions in order of increasing distillation temperature: 1. up to 80oC 2. 80 - 120oC 3. 120 - 180oC 4. 180 - 220oC.
2. Examine the different fractions:
2.1 The colour should change from colourless to yellow.
2.2 The viscosity should increase.
2.3 The high temperature fractions should be more difficult to ignite than the low temperature fractions.
2.4 The high temperature fractions should burn with more soot in the flame than low temperature fractions. Burn the fractions in bottle tops with the cork removed.
2.5 Note the dark residue remaining in the test-tube.

3.22 Separate salt and sand
Prepare a mixture of salt and sand. Put 2 mL of the mixture in a test-tube. Add 5 mL of water and shake until all the salt has dissolved. Pour the contents of the tube into a filter paper in a funnel over an evaporating basin. Wash the test-tube with water and add this to the filter paper. The sand will remain on the filter paper and may be dried and collected. Recover the salt from the filtrate by warming the evaporation basin to drive off the water.

3.23 Solvent extraction of oil from nuts
Put peanuts (groundnuts) or pieces of chopped coconut into a mortar. Add 20 mL of acetone or methylated spirit. Grind the nuts in the solvent as finely as possible. Pour off the liquid into a test-tube and filter into an evaporating basin. Warm the evaporating basin for 10 minutes. The solvent evaporates leaving the oil extracted from the nuts.

3.24 Separate pigments from green leaves with paper chromatography
See diagram 3.24: A chromatogram
A absorbent paper, B solution of crushed leaves
1. Collect green leaves and cut them into very small pieces. Use a mortar and pestle to grind the leaves for five minutes with a small volume of methylated spirit and clean sand until a deep green solution forms. Draw a fine pencil line 5 cm from the end of a 1 cm wide strip of absorbent paper. Suspend the absorbent paper in a test-tube without touching the bottom. Use a fine eye dropper to put one small drop of the solution on the centre of the fine pencil line and let it dry. Add more solution to the same place to make a small concentrated spot. Hang the paper strip with the lower end in the methylated spirit solvent and the spot of green solution above the solvent level. Leave the paper strip in the solvent until the methylated spirit has almost reached the top of the absorbent paper. Capillary attraction draws up the solvent. Mark the chromatogram on the paper to show a top orange band of xanthophyll and a lower green band of chlorophyll. A band of carotene is visible if the solvent is toluene.
2. Repeat the experiment with other solvents, e.g. toluene, acetone (propanone)

3.24.1 Separate mixed inks with paper chromatography
Prepare a mixed solvent from 6 parts of water, 3 parts of methylated spirit, and 1 part of ammonia solution. Put 5 mL of mixed solvent in a test-tube. Prepare mixed ink from equal quantities of red and blue ink. Put a drop of the mixed ink near one end of a 2 cm wide paper strip. Lower the paper strip so that its lower end is in the mixed solvent. Use a stopper to prevent evaporation. As the solvent moves up the paper strip, the component colours of the ink separate to form different coloured bands with red above and blue below. Try other solvents and other inks to obtain good separation of colours.
Repeat the experiment by drawing a line with a ball pen or an ink pen near the end of the paper strip.

3.25 Gases dissolved in a water sample
See diagram 3.25: Gases in water
Stand a container of water in sunlight. Bubbles of air appear. The taste of boiled water is different from the taste of tap water because boiled water has lost its dissolved oxygen. Note the temperature of a sample of water. Boil the water until no more bubbles appear. Collect the gases from the water in an inverted measuring cylinder.

3.26 Separate immiscible liquids of different density
See diagram 3.26: Separation tube
Separate two immiscible liquids of different density, e.g. kerosene (paraffin oil) and water. Use a separating funnel or make a separating funnel with a piece of wide plastic tubing fitted with a one-hole stopper and rubber tubing with a clip. Shake the mixture thoroughly in a closed container then run it into the separating funnel. Wait until a clear boundary appears between the two liquids and then run off the more dense layer into a container below.

3.27 Separate solids using density differences
In industry, a separator concentrating machine shakes mixed ores to separate the different ores. Beach sand often consists of quartz particles mixed with heavier particles such as ilmenite or zircon. Shake a mixture of sand and iron oxide to make them separate into different layers.

3.28 Substances that gain mass when heated, copper foil
Cleaned copper is red-brown. In moist air the surface turns green due to oxidation. The green surface is called a patina. It also forms on old unpolished bronze.
1. Clean a piece of copper foil with steel wool. Hold it in a flame with a pair of tongs. The black copper (II) oxide looks like carbon. To test the substance, drop dilute sulfuric acid on it, then heat it. Blue copper (II) sulfate forms. Test some powdered carbon. No colour change occurs.
2. Show that something is added to the copper from the air. Use a sensitive balance to weigh the copper before and after heating.
3. Use two identical hard-glass test-tubes with one-hole stoppers fitted with bent delivery tubes. Fix both test-tubes to a stand so that the test-tubes slope down with the ends of the delivery tubes under water in a beaker. Put copper foil in the first test-tube and heat with a hot burner flame. After two minutes, heat the empty second test-tube. Move the burner regularly between the two test-tubes until no more bubbles come out of the ends of the delivery tubes. Stop heating both test-tubes. As the test-tubes cool, they suck water up the delivery tube. The test-tube containing the copper (II) oxide sucks up more water.

3.28.1 Substances that gain mass when heated, prepare magnesium oxide
Use magnesium ribbon because magnesium powder is too reactive. Be careful! Do not heat magnesium powder.
Magnesium has density 1.74 g / cm3 and melting point 650oC, but magnesium oxide has density 3.58 g /cm3 and melting point 2800oC because the Mg2+--O2- chemical bond is stronger than the Mg--Mg bond.
1. Hold a 10 cm strip of magnesium ribbon in a pair of tongs. Place the ribbon in a Bunsen burner flame until it starts to burn. Be careful! Magnesium burns with a very bright white light. Magnesium ribbon corrodes slightly in air and burns with an intense white flame to form a white ash of magnesium oxide.
Mg + 1/2O2 --> MgO
2. Clean 25 cm of magnesium ribbon and cut into pieces 1 cm long. Put the pieces into a crucible with a lid. Weigh the crucible + lid + contents = W1. Put the crucible on a pipe clay triangle on a tripod stand. Heat gently then strongly. Use tongs to raise the lid. The magnesium darkens before it melts. When the magnesium starts to burn, put the lid back on the crucible and remove the burner. Every few seconds raise the lid slightly to let more air enter. Do not let white magnesium oxide smoke escape. When the magnesium does not burn after you raise the lid, remove the lid and heat the crucible strongly. Hold the lid ready in case the magnesium starts to burn again. Let the crucible cool. Again weigh the crucible + lid + contents = W2. Note W2 > W1. The formation of magnesium oxide causes the increase in weight.

3.28.2 Substances that gain mass when heated, prepare lithium oxide
Heat pieces of lithium metal shot on a metal spoon (deflagrating spoon). Note the violet glow when it starts to burn, then put the burning lithium in oxygen gas.

3.28.3 Substances that gain mass when heated, prepare calcium oxide
Heat a shaving of calcium metal in a crucible and heat it with a Bunsen burner for 10 -15 minutes because it is difficult to ignite.

3.28.4 Collect and weigh the gaseous products of a burning candle
See diagram 3.29: Gaseous products of burning candle
Candle wax is a mixture of different alkanes (paraffins) saturated hydrocarbons with general formula CnH2n+2that are solid at room temperature. Soda lime is a grey-white mixture of sodium hydroxide and calcium hydroxide as granules or powder that absorbs the products of combustion, carbon dioxide and water. Use soda lime instead of sodium hydroxide because soda lime is not deliquescent. Weigh a candle, C1. Weigh a U-tube containing granules of soda lime, U1. Put a candle under an inverted glass filter funnel connected to one arm of the U-tube. Attach a filter pump to the other arm to draw air through the U-tube. Light the candle and turn on the filter pump to draw air over the candle. Let the candle burn for five minutes. Extinguish the candle and disconnect the filter pump. Weigh the candle again, C2. The candle has lost weight, C1-C2. When the U-tube is cool, weigh it again, U2. The U-tube containing the soda lime has gained weight, U2-U1. The U-tube gains more weight than the candle loses weight (U1-U2) > (C2-C1) for two reasons: 1. The candle wax combines with oxygen gas in the air to form carbon dioxide gas and water. 2. The air sucked in by the filter pump contains some water vapour absorbed by the soda lime. To measure the weight of water absorbed from the air, in a control experiment, repeat the experiment for the same period, but without the candle.

3.30 Substances may decompose and lose mass when heated, thermal decomposition
See diagram 3.30: Collecting the products of heating copper sulfate crystals
Substances that remain stable after heating constantly with a Bunsen burner flame may be listed under the heading "Substances that do not decompose when heated". However, all compounds breakdown when heated to a high enough temperature. Metal compounds higher in the activity series are usually more stable than compounds of metals lower in the activity series. Salts that remain stable when heated constantly with a Bunsen burner flame are calcium sulfate, potassium chloride, potassium sulfate, sodium carbonate, sodium chloride, and sodium sulfate. Ammonium oxalate (NH4COO)2, and ammonium dichromate (NH4)2Cr2O7, decompose before melting. Ammonium sulfate (NH4)2SO4, decomposes above 280oC.

3.30.1 Carbonates mostly decompose to form carbon dioxide and a metallic oxide.
Sodium carbonate and potassium carbonate do not decompose when heated to a high temperature.
CaCO3 (s) --> CaO (s) + CO2 (g)
white calcium carbonate --> white calcium oxide + carbon dioxide
CuCO3 (s) --> CuO (s) + CO2 (g)
blue-green --> black
MgCO3 (s) --> MgO (s) + CO2 (g)
white --> white
PbCO3 (s) --> PbO (s) + CO2 (g)
white --> yellow
ZnCO3 (s) --> ZnO (s) + CO2 (g)
white --> white (cold) or yellow (hot)
Ammonium carbonate may decomposes completely without heating when exposed to the air, to give ammonia, water and carbon dioxide.
(NH4)2CO3 (s) --> 2NH3 (g) + H2O (vapour) + CO2 (g)
colourless

3.30.2 Hydrogen carbonates (hydrogencarbonates, bicarbonates) decompose to form the metal carbonate, water and carbon dioxide.
Calcium bicarbonate and magnesium bicarbonate can exist only as a solution that on heating form the metal carbonate, water and carbon dioxide.
Sodium hydrogen carbonate, NaHCO3 (sodium bicarbonate) begins to lose carbon dioxide at 50oC to form sodium carbonate. A solution of a sodium hydrogen carbonate begins to lose carbon dioxide at 20oC.
Heat sodium hydrogen carbonate crystals. The crystals lose water and carbon dioxide, and at 100oC are converted to sodium carbonate.
2NaHCO3 (s) --> Na2CO3 (s) + CO2 (g) + H2O (vapour)
colourless --> colourless
Ca(HCO3)2 (aq) --> CaCO3 (s) + CO2 (g) + H2O (vapour)
Mg(HCO3)2 (aq) --> MgCO3 (s) + CO2 (g) + H2O (vapour)
2KHCO3 (s) --> K2CO3 (s) + CO2 (g) + H2O (vapour)

3.30.3 Hydroxides decompose to form water and the metallic oxide

3.30.4 Nitrates decompose to form oxygen gas, nitrogen dioxide and a metallic oxide.
2Ca(NO3)2 (s) --> 2CaO + 4 NO2 (g) + O2 (g)
colourless --> white
2Cu(NO3)2 (s) --> 2CuO + 4 NO2 (g) + O2 (g)
blue --> black
2Pb(NO3)2 (s) --> 2PbO + 4 NO2 (g) + O2 (g)
colourless --> yellow
Lead nitrate decomposes at 470oC.
2Zn(NO3)2 (s) --> 2ZnO + 4 NO2 (g) + O2 (g)
colourless --> white (cold), yellow (hot)
Potassium nitrate and sodium nitrate first melt and then decompose to give the metal nitrite and oxygen gas. Potassium nitrate melts at 336oC. Sodium nitrate melts as 316oC.
2KNO3 (s) --> 2KNO2 (s) + O2 (g)
colourless --> colourless
2NaNO3 (s) --> 2NaNO2 (s) + O2 (g)
colourless --> yellow
Silver nitrate decomposes to give the metal, nitrogen dioxide and oxygen gas.
2AgNO3 (s) --> 2Ag (s) + 2NO2 (g) + O2 (g)
colourless --> silver
Ammonium nitrate decomposes to form water vapour and nitrous oxide, N2O (laughing gas) so the ammonium nitrate disappears.
NH4NO3 (s) --> N2O (g) + H2O (g)
colourless

3.30.5 Oxides of most metals are stable
Oxides of potassium, sodium, calcium, magnesium, aluminium, zinc, iron, lead and copper do not decompose.
Grey-black silver oxide decomposes into the metal and oxygen gas.
2Ag2O (s) --> 4Ag (s) + O2 (g)
silver oxide --> silver + oxygen
Heat zinc oxide in a crucible. Zinc oxide becomes yellow when hot and white when cold but no change in weight occurs. The substance does not decompose and does not gain anything from the air or lose anything to the air. Zinc oxide has melting point above 1,800oC.
ZnO (s) <--> ZnO (s)
white (cool) yellow (hot)
Thermal decomposition of higher oxides of lead
2PbO2 (s) --> 2PbO (s) + O2 (g)
brown lead dioxide --> yellow lead oxide + oxygen gas
2Pb3O4 (s) --> 6PbO (s) + O2 (g)
red trilead tetroxide --> yellow lead oxide + oxygen gas

3.30.6 Sulfates if heated very strongle may decompose to form the metallic oxide, sulfur dioxide and oxygen gas
Put 4 cm of crushed blue copper (II) sulfate crystals in a dry test-tube fitted with a one-hole stopper and delivery tube. Heat the dry test-tube and crystals gently. Note whether vapour collects on the cooler parts of the dry test-tube and whether any liquid collects in the receiving test-tube. Note any change of colour of the crystals from blue to white. Identity the liquid in the receiving test-tube by measuring the boiling point. When all the copper (II) sulfate crystals have become white and the dry test-tube has cooled, pour the liquid in the receiving test-tube back on the white crystals. Note whether the blue colour of the crystals is restored and if any heat is given off.
2CuSO4 (s) --> 2CuO (s) +2SO2 (g) + O2 (g)
grey white --> black
2PbSO4 (s) --> 2PbO (s) +2SO2 (g) + O2 (g)
white --> yellow
2ZnSO4 (s) --> 2ZnO (s) +2SO2 (g) + O2 (g)
white --> white (cold) yellow (hot)

3.30.7 Sulfites mostly decompose into the metal oxide and sulfur dioxide
Sulfites of sodium and potassium do not decompose when heated.
CaSO3 (s) --> CaO (s) + SO2 (g)
white --> white
MgSO3 (s) --> MgO (s) + SO2 (g)
white --> white
ZnSO3 (s) --> ZnO (s) + SO2 (g)
white --> white (cold) yellow (hot)

3.30.8 Salts with water of crystallization, hydrous salts, lose their water of crystallisation, and become anhydrous powders and then become stable
They also lose their crystalline shape and sometimes their colour.
Prepare test-tubes containing 1 cm of 1. iodine crystals 2. sodium hydrogen carbonate granules or crystals 3. silica sand 4. zinc oxide. Fix a cotton wool plug in the mouth of each test-tube to prevent loss of solid during heating, then weigh each test-tube. Heat each test-tube and cotton wool plug thoroughly and weigh it again. Note any change in weight because of the loss of water of crystallization.
(blue) copper (II) sulfate crystals + heat < = > (white) anhydrous copper (II) sulfate + water.
CuSO4.5H2O (s) --> CuSO4 (s) + 5H2O (vapour)
blue --> grey white
Na2CO3.10H2O --> Na2CO3 (s) + 10H2O (vapour)
colourless --> white

3.30.9 Boric acid, H3BO3, loses water until it decomposes to the anhydride, B2O3.

3.30.10 Oxalic acid begins to sublime at 100oC, becomes anhydrous at 189oC and when heated rapidly decomposes into carbon dioxide, carbon monoxide, formic acid and water.

3.30.11 Potassium chlorate, KClO3, decomposes above 368oC into potassium perchlorate and oxygen gas
Potassium ferricyanide, K2Fe(CN)6, decomposes before melting. Monosodium orthophosphate, NaH2PO4.H2O, and disodium orthophosphate [disodium hydrogen phosphate (V)] Na2HPO4.12H2O, lose water of crystallization.
10KClO3 <--> 6KClO4 + 4KCl + 3O2

3.30.12 Sublimation, iodine, camphor, naphthalene
Sublimation is when a solid changes directly into a gas without melting. Iodine, camphor and naphthalene do not decompose when heated but sublime from the crystal to the vapour form.
Put black, shiny crystals of iodine in an evaporating dish. Cover the dish with a piece of filter paper and stand a filter funnel upside down on the dish. Heat the dish gently. Purple vapours rise through the filter paper. As they cool in the filter funnel, shiny black crystals of iodine form again.

3.30.13 Silica sand
Silica sand consists of pieces of silicon (IV) oxide (SiO2) crystals. Heat sand in a crucible. The sand particles may break up physically, but do not break up chemically.

3.30.14 Ammonium dichromate is an orange coloured crystalline substance
It starts decomposing with sparks and gives out heat to form a green fluffy powder chromic oxide, nitrogen and water.

3.30.15 Ammonium chloride
Put ammonium chloride into the bottom of a dry test-tube and warm it over a small flame. The solid decompses to form ammonia and hydrogen chloride. Some of it recombines at the cool upper part of the test-tube as a white layer. Heat the recombined layer again and it moves further up the test-tube. This process is recombination not sublimation.
NH4Cl (s) <--> NH3 (g) + HCl (g)