School Science Lessons
Topic 8 Heat sources, Bunsen burner, candles, combustion, spirit burner, heat chemicals, substances that decompose / lose mass when heated
2014-11-18
Please send comments to: J.Elfick@uq.edu.au

Table of contents
8.1.1 Candles, (Chemistry)
8.1.1a Candles, (Physics)
8.6.0 Combustion, conditions for combustion, to ignite, ignition point
8.2.0 Elements combine with oxygen gas when heated in air
8.0.1 Heat sources
8.1.2 Spirit burner, alcohol lamp
8.4.0 Heat substances that decompose when heated, but may be reformed
3.30 Heat substances that decompose and lose mass when heated
8.2.4 Thermal decomposition of acids
8.2.9 Thermal decomposition of sucrose crystals, C12H22O11
8.2.9.1 Invert sugar, C12H24O12, HFCS

8.1.1 Candles, (Chemistry)
16.1.1.1 Alkanes, (CnH2n+2), paraffins (See: Candle, paraffin wax)
8.1.9 Aluminium foil below candle flame
6.35 Burn candles in closed containers
2.44 Candle flame (Primary)
3.29 Collect and weigh the gaseous products of a burning candle
6.36 Cooling candle wax (Primary)
8.1.0 Candle, paraffin wax
8.1.5 Candle flame consists of burning vapours
8.1.4 Candle flame forms carbon dioxide
8.1.6 Candle flame forms water
8.1.2a Carbon, soot, from a candle flame
8.1.7 Dark region of a candle flame
8.1.12 Egg in a candle flame
8.1.10 Floating tea candle
8.1.3 Hottest part of a candle flame
8.1.13 Melt candle wax
8.1.1 Parts of a candle flame
8.1.11 Prepare beeswax candles
3.34.1.4 Pouring tests for carbon dioxide
3.77 Reactions of magnesium with carbon dioxide, sparkler experiment
Re-lighting candles, (Happy birthday candles you cannot blow out!), "Trick Candles", "Magic Candles"
16.4.2.4 Rocking candle, balancing candle, burn a candle at both ends
6.6.7 Tests for absorption of oxygen during plant respiration
8.1.8 Test gases from the wick

8.1.1a Candles, (Physics)
Atmospheric moisture: 37.15
Burn candles in closed containers: 6.35
Candela, (candlepower), Luminous intensity, (candela, cp): 6.3.1.7
Candela, candlepower, lumen, lux, Luminance and illuminance: 4.104
Candle burner See diagram 3.2.0.2
Candle burns below water level: 22.5.8
Candle flame on a turntable, 15.2.21
Candle in a bottle, candle in dropped jar, 16.2.1.10
Coanda effect, a spoon touches a water stream (See 4.): 13.2.27
Cold candle in a parabolic mirror, 28.3.9
Convection box with two chimneys, smoke house: 23.6.6
Convection box: 37.13
Convection disc, heat snake, revolving picture lamp: 23.6.4
Convection heat snake 3.39
Cooling candle wax, candle "lava" (Primary): 6.36
Copper coil candle snuffer: 23.7.6
Float a lighted candle: 4.204
Heat snake (convection), (Primary): 3.39
Image of candle in a glass of water: 28.2.23
Light travels in straight lines, (Primary): 5.10
Metal wire in candle flame shows electric field: 31.7.24
Mirrors at an angle with a candle: 28.2.20
Parallel mirrors with a candle: 28.2.21
Rocking candle, balancing candle, burn a candle at both ends: 16.4.2.4
Rotate a candle: 15.5.9
Smoke moves up and down, (Primary): 4.12
Soap film minimal surfaces (See 3.) 19.3.22
Sound waves extinguish a candle flame: 26.4.9
Sources of light, candlepower: 32.5.8.1
Trace convection currents with a lighted candle: 4.28
Trace convection currents: 37.14

6.35 Burn candles in closed containers:
6.35.1 Burn a floating candle under a jar
6.35.2 Burn a long candle standing under a jar
6.35.3 Burn long and short candles under a jar
6.35.4 Burn a candle, lighted inside a jar
6.35.5 Burn one, two and three happy birthday cake candles inside a jar
6.35.6 Burn a candle in a falling plastic bottle
6.35.7 Burn a tall candle under a jar and a short candle under an identical jar
3.29 Collect and weigh the gaseous products of a burning candle
6.35.8 Test for carbon dioxide from a burning candle

8.6.0 Combustion, conditions for combustion, to ignite, ignition point
8.6.0 Combustion, conditions for combustion, to ignite, ignition point
Experiments
8.6.1 Burn (Monopoly money, fake money) bank notes, ethanol
3.5.13 Autoignition temperature
8.6.0 Combustion, conditions for combustion, to ignite, ignition point
8.6.3 Carbon dioxide is a product of combustion
15.3.12 Conditions necessary for rusting
3.5.0 Fire safety
3.52.1 Mass of iron and temperature increases during rusting
3.52.3 Metals can prevent rusting
15.3.4 Need for oxygen for rusting
3.52.2 Oxygen gas combines with iron during rusting
8.6.2 Oxygen gas is necessary for combustion
8.6.5 Respiration is a form of combustion
17.3.4 Spontaneous combustion, sugar with potassium chlorate

8.2.0 Elements combine with oxygen gas when heated in air
8.2.0 Elements combine with oxygen gas when heated in air
3.29 Collect and weigh the gaseous products of a burning candle
8.2.11 Heat aluminium foil to form aluminium oxide
8.2.14 Heat calcium metal to form calcium oxide
8.2.12 Heat copper foil to form copper (II) oxide
2.21 Heat different substances, (Primary)
8.2.13 Heat iron, nail, wire, filings, to form iron (II) oxide
8.2.10 Heat lithium metal to form lithium oxide
8.2.16 Heat magnesium ribbon to form magnesium oxide
8.2.15 Heat sulfur to form sulfur dioxide
13.3.6 Reactions of magnesium oxide

8.1.0 Heat sources
8.1.0 Heat sources
Gases, (household gas, laboratory gas)
3.42 Burn different substances (Primary)
5.43 Burn to make carbon (Primary)

8.1.2 Spirit burner, alcohol lamp
SC7001 Burner, Alcohol, Glass, metal Cap, wick, 10 pack, "Scientrific" (commercial website)
SC7005 Burner, Alcohol, Spare Wicks, 10 pack, "Scientrific" (commercial website)
16.5.1 Methylated spirits, (methylated spirit, spirits, denatured alcohol
2.20 Spirit burner (alcohol lamp), (Primary)
8.3.1 Spirit burner, alcohol lamp, methylated spirit burner, alcohol burner, alcohol lamp

8.4.0 Heat substances that decompose when heated, but may be reformed
8.4.0 Heat substances that decompose when heated, but may be reformed
8.2.2 Heat cobalt chloride crystals, Tests for water with cobalt (II) chloride
8.2.1 Heat copper (II) sulfate-5-water crystals, Tests for water
8.3.5 Heat calcium sulfate, gypsum

6.35 Burn candles in closed containers
See diagram 6.35: Candles burning in closed containers
When you light a candle wick, some heat from the burning wick melts the wax at the top of the candle. The melted wax rises in the wick because of capillary forces and some of that wax evaporates to form a vapour that rises then burns with a bright flame. The flame of a burning candle in a closed container will go out after some time, not because all the oxygen is used up, but because the hot carbon dioxide and water vapour produced by the ignition accumulates down from the top of the container, displaces other gases, including oxygen, and finally stifles the flame.

6.35.1 Burn a floating candle under a jar
See diagram 6.35: Candles burning in closed containers
1. Fix a short candle to a cork. Put water in the trough so that the water level is about half the length of the candle. Put water in the trough. The short candle on the cork floats in the water.
2. Measure the depth of the water in the trough. Measure the length of the jar. The jar is full of air so you are really measuring how much air in the jar. Light the candle.
3. Place an inverted jar over the candle quickly so that the mouth of the jar is under water.
4. While the candle is alight, gas bubbles come out from under the jar and rise through the water in the trough. The candle flame heated the air in the jar causing the air to expand. Lighter gases go to the top of the jar because gravity exerts a stronger pull on denser gases. As these colder, denser gases move downward, lighter gases take their place at the top of the jar.
5. The candle flame gets smaller, splutters, then goes out because some oxygen was converted to carbon dioxide gas. The blackened wick is carbon.
6. The volume of gases in the jar decreases because the gases in the jar cool, some water vapour condenses inside the jar and on the wick and some carbon dioxide dissolves in the water. The gases in the jar cool and contract to a smaller volume than before so the air pressure in the jar becomes less than the atmospheric pressure. The air pressure on the surface of the water in the trough pushes the water up into the jar. The water level rises in the jar and drops slightly in the trough.
7. The decrease of volume inside the jar is about 20%. Some people think that the candle flame has "consumed" all the oxygen and that the amount of oxygen consumed = the decrease in volume of gases in the jar, thus showing that air contains 20 % oxygen.

6.35.2 Burn a long candle standing under a jar
See diagram 6.35: Candles burning in closed containers
1. Stand a long candle in a trough. Put 3 rubber plugs on the bottom of the trough so that the rim of a jar may rest on them and allow water to enter under the jar.
2. Put water in the trough so that the water level is about half up the length of the candle. Measure the depth of the water in the trough. Measure the length of the jar.
3. Light the candle. Quickly place an inverted jar over the lighted candle so that the mouth of the jar is under water, resting on the three rubber plugs.
4. While the candle is alight, gas bubbles come out from under the jar and rise through the water in the trough.
5. The candle flame gets smaller, splutters, then goes out.
6. The water level rises in the jar and drops slightly in the trough.
7. Some water has condensed inside the jar and on the wick.
8. The decrease of volume inside the jar is about 20%.

6.35.3 Burn long and short candles under a jar
See diagram 6.35.3: Burn long and short candles under a jar
1. Attach a tall candle and a short candle to the bottom of a trough. Add water to the trough and note the water level. Add ink or cochineal to colour the water. Light both candles. Put an inverted jar over the burning candles. The tall candle extinguishes first then the short candle. Hot gas and less dense products of combustion including carbon dioxide gas and water vapour have filled the jar from the top down to extinguish the tall candle flame first. The cloud of hot carbon dioxide and water vapour gets thicker and so descends to reach the tall candle flame first. Some hot gases push out under the rim of the jar to form bubbles around the jar in the trough. When the candles are extinguished, the hot gases cool and contract to form a partial vacuum and the water level rises inside the jar. This experiment shows that in the above experiments the lighted candle did not "use up the oxygen in the jar" because when the tall candle goes out the short candle keeps burning for a while. Candles will not burn when air has lost about 30% of its oxygen.

6.35.4 Burn a candle, lighted inside a jar
See diagram 6.35.4: Light a candle inside a jar
This is a more accurate way of doing the experiment because the candle is not already burning when the inverted jar is placed over it. So you are comparing the volumes of room temperature air before combustion and after combustion.
1. Use adhesive tape to attach the head of a match to a candle wick.
2. Place an inverted jay over the candle.
3. Light the match by focussing sunlight on it.
4. Leave the apparatus to cool to room temperature.
5. Measure the increase in water level inside the inverted jar. The increase is about 5%, not 20%.

6.35.5 Burn one, two and three happy birthday cake candles inside a jug
See diagram 6.35.5: Burn candles over water
1. Fix the candles to Plasticine at the bottom of the tray then do the experiment as before
2. The air in the jar is less heated with smaller candles so you probably do not see bubbles coming out from under the jar. The jars contained about the same amount of oxygen but the jar with the three candles showed the greatest water rise because more heat was produced by the candles in it.

6.35.6 Burn a candle in a falling plastic bottle
See diagram 6.35.6: Burning candle in falling plastic bottle
Attach a candle to the screw top of a plastic bottle. Darken the room. Light the candle, attach the screw top, and drop the plastic bottle. The candle flame heat the air around the wick but the hot air falls with the same speed as the candle. So the wick remains surrounded by the hot air and the products of combustion of the candle wax vapour. Immediately after the plastic bottle is dropped,  the candle flame becomes dimmer as a blue spherical disc. The flame receives only a small supply of oxygen by diffusion because the carbon dioxide remains around it and is replaced by the oxygen when it rises. Hot carbon dioxide, and some water vapour,  accumulate at the top over the candle until the flame is extinguished.

6.35.7 Burn a tall candle under a jar and a short candle under an identical jar
See diagram 4.9: Burning candles over water | See diagram 6.35.7: Burn short and tall candles in separate jars
1. Burn two candles over water. Attach a tall candle and a short candle to the bottom of a trough. Add water to the trough and note the water level. Light both candles. Put a large jar upside down over the candles. The tall candle extinguishes first then the short candle. Hot gas products of combustion, including carbon dioxide gas, have filled the jar from the top down to extinguish the candle flames. Some hot gases push out under the rim of the jar to form bubbles around the jar in the trough. When the candles are extinguished, the hot gases cool and contract to form a partial vacuum, and the water level rises inside the jar
2. In the jar over the tall candle, the water rises almost immediately, mostly before the flame is extinguished. In the jar over the short candle, the water hardly rises until the flame dies, then rises rapidly. When the candle goes out, most of the cooling of gases occurs at the top of the jar, where the air is hottest and contact with the glass is greatest. The taller candle begins to dim earlier, caused by the smothering cloud of carbon dioxide, so the flame gives off less heat. So the rate of cooling is largely determined by the temperature of the air at the top of the jar. The flames of shorter candles are farther away from the top, so they take longer to heat up the air up there. By the time a short candle completely goes out, much carbon dioxide has been produced, mostly lying close to the water. After the air cools, there is more opportunity for the carbon dioxide to dissolve into the water than with a tall candle. This may contribute to the rapid rise after a short candle is extinguished, because the air is losing carbon dioxide.

6.35.8 Test for carbon dioxide from a burning candle
Cut off each end of a plastic drink bottle to make a cylinder. Place the cylinder vertically around the candles. Pour sodium bicarbonate solution then tartaric acid solution into the water around the candles. The acid reacts with the base to form bubbles of carbon dioxide gas. As the cylinder fills with carbon dioxide gas the short candle flame then the long candle flame will be extinguished as the carbon dioxide gas displaces the air upwards. Try to relight the candle with a match or taper. The flame is extinguished when it reaches the carbon dioxide layer. Make a loop with a piece of wire, dip it in a soap or detergent solution and blow a small bubble so that it falls gently into the cylinder. The bubble will stop falling when it reaches the carbon dioxide gas layer. Light a fireworks "sparkler" and place the lighted end in the cylinder.
BE CAREFUL! The sparkler continues to burn because it contains magnesium powder that reacts with carbon dioxide gas.

Tiny black specks of carbon form on the inside of the cylinder. When the sparkler has finished burning, you can relight the candle with a match because all the carbon dioxide has reacted with the magnesium in the sparkler.

8.0.1 Heat sources
See 16.1.1cc: LPG (liquefied petroleum gas, LP gas)
See diagram 3.2.0.0: Candle flame | See diagram 3.2.0.2: Candle burner | See diagram 3.1.4.6: Bunsen burner and candle flame
A flame is a region where a gas emits light because of the high temperature. Burning, i.e. combustion, needs oxygen gas, is exothermic process and has reaction products are carbon dioxide and water. Spontaneous combustion does not need external heat energy to start it, e.g. white phosphorus in air. Combustible substances catch fire easily, e.g. paper. You can smother a flame to cut off the oxygen gas supply and put out the fire. Water is used to put out fire because it reduces the temperature of substances below its ignition point. However, the temperature of burning oil is too high for an oil fire to be extinguished by water. Most fire extinguishers either reduce the ignition temperature or cut off the oxygen gas supply.
1. Light a match by striking the match along a roughened surface. The match head contains red phosphorus. Striking causes friction to generate heat to raise the temperature of red phosphorus to ignition temperature.
2. Heat water in a paper cup. The paper will not catch fire because the water keeps the temperature of the paper lower than its ignition point.
3. Study the flame of a Bunsen burner and a candle. A flame is the region where combustion occurs. The colour of the flame depends on the temperature and the substance burning. Hydrocarbon flames are either blue or yellow. A blue flame is not luminous. It occurs because of complete burning of hydrocarbons with plenty of oxygen gas. The blue flame does not leave any residue or any other gases. A yellow flame occurs when there is insufficient oxygen gas. It is a luminous flame. The temperature is lower than the blue flame and it leaves black soot and other residues. A candle contains wax made from petrochemicals. The wick is lighted, and the flame melts the wax. The evaporated wax rises and catches fire. As the vapours rise higher, they stay longer in the hot regions of the flame and start burning completely with oxygen gas.
4. The candle flame has three regions. The inner zone appears black, contains unburned wax vapours and is the least hot region of the candle. The middle zone is where the wax vapours start burning giving a yellowish flame of partially burnt gases because of insufficient gases for complete combustion. The flame in  the luminous region but not very hot. The outer zone is where the wax vapours have enough oxygen gas to burn completely. The flame appears blue and the temperature is very high, up to 1 100oC or more.
5. Study a flame in a gas stove. Rapid combustion releases a large amount of heat in a short time, e.g. Lighting LPG gas in a kitchen stove.
6. Study a match flame. Below the ignition temperature, e.g. white phosphorus 35oC, a combustible substance in oxygen gas will not catch fire. The wood used in the match has a similar ignition temperature.

8.1.0 Candle, paraffin wax
See diagram 3.2.0.0: Candle flame | See diagram 3.2.0.2: Candle burner | See diagram 3.1.4.6: Bunsen burner and candle flame
Candle wax is a mixture of different alkanes that are solid at room temperature. Candles are usually made of paraffin wax that is a residue from the distillation of petroleum. With enough air, the wax burns to form carbon dioxide and water. With insufficient air, the wax burns to form carbon monoxide and smoke containing carbon. The teardrop-shaped flame is called a diffusion flame because oxygen gas diffuses in from the air to the combustion region and hydrocarbon vapour diffuses out wards form the wick. Heat radiated from the burning wick melts the wax drawn up the wick by capillarity. The melted wax vaporizes to form a cloud of hydrocarbon molecules that diffused into the flame and are broken down into small molecules by the intense heat of the flame. The smaller molecules react with oxygen. The smoke from the flame contains carbon particles (soot) water vapour and various products of the reactions of the hydrocarbon particle with oxygen gas.
1. Cut the top off one of a clear plastic soft drink bottle and fill it with water. Float a candle on the water. Light the candle. A cup of molten wax forms around the wick. As the candle flame burns the wax melts and moves up the wick by capillarity then is converted to a vapour by the heat of the flame. The vapour rises and burns to form more flame. The ascending current of air, produced by the heat of the candle, keeps the outside edge cool, and forms a cup for the melted wax around the wick. The rising vapour draws up cold air containing oxygen gas.
2. See the shape of the flame with three regions:
2.1 The innermost part is a dark area, the shape of the flame around the wick. It is not luminous and consists of the vapour from the molten wax.
2.2 The coloured part of the flame is orange-yellow to blue near the bottom. It is where some combustion occurs.
2.3 The outer, almost colourless region of the flame is where most combustion occurs because more air (oxygen gas) is available. Blow out the candle then ignite again the vapour quickly with a lit match. The flame will go down and ignite again the candle. Complete combustion of the wax hydrocarbon should produce carbon dioxide and water only but the candle flame is not hot enough to allow complete combustion so a mixture of gases and tiny specks of black carbon (soot) forms. The glowing carbon particles glow and are the main emitters of candle light. Hold a white plate above the flame to see the black soot. Suspend a suspended spiral of paper above the candle flame. The spiral turns because of the force of the rising hot gases from the candle flame.
3. Relight a candle. Light a match, then blow out the candle, keeping the match lit. Then immediately bring the burning match close to the smoking candle wick and observe closely. Note when the candle flame reignites.
4. Repeat this experiment with a cold candle that has not been recently burning. Wax vapour still exists in the space between the hot wick and the match flame. Candle wax, or paraffin, is a mixture of high molecular weight saturated hydrocarbons consisting mostly of long chains of (-CH2-) units.
The simplest hydrocarbon, methane, burns as follows:
CH4 + 2O2 --> CO2 + 2H2O
A single (-CH2-) unit burns as follows:
2CH2 + 3O2 --> 2CO2 + 2H2O

8.1.1 Parts of a candle flame
See diagram 3.2.0.0: Candle flame
1. A candle contains wax made from petrochemicals. The wick is lighted, and this melts the wax. The evaporated wax rises and catches fire. As the vapours rise higher, they stay longer in the hot regions of the flame and start burning completely with oxygen gas. The candle flame has three regions. The inner zone appears black, contains unburned wax vapours and is the least hot region of the candle. The middle zone is where the wax vapours start burning giving a yellowish flame of partially burnt gases because of insufficient gases for complete combustion. The flame is a luminous region but not very hot. The outer zone is where the wax vapours have enough oxygen gas to burn completely. The flame appears blue and the temperature is very high.
2. Hold a piece of white cardboard behind the flame so that you can see each part clearly. The candle flame has three regions. Each region has the shape of the flame around the wick.
2.1 The innermost region closest to the wick consists of vapours from the molten wax. It is dark in colour because air cannot reach that region, so the gases are not burning.
2.2 The second region is bright yellow to orange to blue near the bottom. It forms much light. Incandescent soot particles cause some orange and yellow glow. The red area near the centre of the flame is about 800oC. The outer orange and yellow areas are hotter than this region. Some combustion occurs in this region.
2.3 The third region, the outer rim of the flame, is practically colourless. It is a very faint blue colour and is the hottest part of the candle flame. The blue colour shows that oxygen is mixing with the wax molecules. Most of the combustion occurs in this region. Complete combustion of the paraffin hydrocarbons should produce carbon dioxide and water only, but the candle flame may not be hot enough to produce complete combustion, so intermediate substances form. Tiny black specks in this region are particles of carbon (soot) that glow on ignition and emit most of the light from the candle.
C + O2 --> CO2 + light energy

8.1.2a Carbon, soot, from a candle flame
The soot deposited is the carbon used in the manufacture of inks and motor tires. Whenever fuels, e.g. kerosene (paraffin oil) or coal or wood, burn with insufficient oxygen, similar deposits of carbon (soot) can be seen.
1. Hold a glass rod in the centre of the flame. The rod becomes coated with a sooty black film called lamp black (carbon black). Carbon deposits on the glass rod because not enough oxygen is available for complete combustion.
2. Hold a wire gauze heating mat over the candle flame. The wire gauze cools the flame by conduction and carbon, soot, deposits.
3. Sprinkle flour on the candle flame. The flour particles sparkle as they catch fire to leave specks of carbon.
4. Bend lemon peel near a candle flame. The squeezed peel emits oil and water. Some of the oil burns in the flame to leave carbon particles.

8.1.3 Hottest part of a candle flame
Push a piece of cardboard sideways into the flame. The outside of the flame forms a sooty ring as it scorches the cardboard.
So the hottest part of a candle flame is near the edge of the flame where there is enough oxygen to burn all the vaporized candle wax completely. The burning gases rise to the tip of the flame so that is the hottest part. The flame near the wick is yellow because there is not enough oxygen for complete combustion and this causes the deposit of black (unburnt) carbon on the wick.

8.1.4 Candle flame forms carbon dioxide
Place a glass funnel over the candle flame.
1. Hold a lighted match in the hot air coming out of the stem of the funnel. The match goes out.
2. Fix a test-tube over the stem of the funnel to collect some hot air. Invert the test-tube, add limewater, seal the end of the test-tube and shake it. The limewater turns cloudy, indicating carbon dioxide.

8.1.5 Candle flame consists of burning vapours
1. Blow out a candle flame then quickly insert a lighted taper into the rising vapours. The candle lights again.
2. Use an L-shaped glass tube to lead vapours from a burning candle into a cool beaker. A grey-white vapour condenses into a solid.

8.1.6 Candle flame forms water
1. Hold a very cold beaker over a candle flame. Water droplets form inside the beaker.
2. Sprinkle ice cubes with salt then wrap them in aluminium foil. Hold the foil bundle over a candle flame and note the water droplets forming on the aluminium foil.

8.1.7 Dark region of a candle flame
Hold a glass tube so that it slants upwards and the bottom end is as close as possible to the wick. Light a match and hold it close to the gases coming out of the end of the tube. Gases burn at the end of the glass tube. These gases have come from the dark region of the flame where there is not enough air to burn them.

8.1.8 Test gases from the wick
Light the candle, let it to burn for five seconds and then blow out the flame. Immediately, light a match and hold it near the smoke, vapour trail, coming from the wick. A flame will race back along the vapour trail and reignite the candle. This shows that the gases from the wick are flammable.

8.1.9 Aluminium foil below candle flame
Cut a slot in a piece of aluminium foil and slide it just below the base of the flame and above the melted wax. The flame dies down or becomes extinguished because the foil conducts away the heat so you cannot ignite the gases.

8.1.10 Floating tea candle
A tea candle has about 3 cm diameter, 1.4 cm height and weighs about 10 g. Some people put them in a cut down plastic drink bottle to serve as a cheap lantern that is not blown out by the wind. Float a lighted tea candle in water. The flat top of the candle wax forms a cup of molten wax around the wick. The burning candle should balance symmetrically when floating and the cup of molten wax is also symmetrical. As the wick burns, the wax nearby melts and molten wax is drawn up through the wick by capillarity. As the molten wax nears the flame it evaporates and the vapour rises and ignites. The ascending current of air above the flame keeps the outside edge of the candle wax cool forming a cup for the molten wax around the wick. The draft of ascending hot gases draws up cooler air alongside the body of the candle and supplies oxygen to maintain the burning of the vapour.

8.1.11 Prepare beeswax candles
Household candles, votive candles for churches and birthday cake candles are usually made of paraffins. However, specialist suppliers sell different kinds of candle wax and wicks so you can make your own novelty candles, e.g. candle paraffin with specific melting points, different waxes for different lights, e.g. beeswax and different odours for aromatherapy.
Heat some beeswax in a tin can floating in hot water. Put a piece of white cotton thread in the melted wax for a wick. Use a fork to swirl the wick through the wax then place it to run through the centre and stick out the top by 1 cm. Let the wax cool until solid. Light the beeswax candle and compare the flame with the flame of the other candles. Beeswax comes from bee honeycomb. It is mainly an ester of palmitic acid, C15H31COOC30H61. If you can make candles with the same shape and weight from different waxes, you can compare their flames and rates of burning.

8.1.12 Egg in a candle flame
Hold an egg near the top of a candle flame. The egg becomes covered in black soot. Put the egg in a dish of water. The egg now looks like a silver mirror. A layer of air as bubbles has formed between the soot and the shell of the egg. Light reflects back from the bubbles. Leave the egg in the water. Gradually all the bubbles dissolve and the egg looks black again.

8.1.13 Melt candle wax
Most candle waxes melt at about 60oC. Do not melt candle wax over direct heat because the vapour may ignite. If it ignites, smother the flames with a lid, fire blanket, sodium carbonate powder, or moist towel, but do not use water. Melt candle wax in a heat resistant container in gently boiling water, e.g. in an electric frying pan or over a hot plate.

8.2.0 Elements combine with oxygen gas when heated in air
1. Put a small quantity of sulfur on a deflagrating spoon and set it alight with a Bunsen burner flame. Note the appearance of the burning sulfur and then lower it into a test-tube containing oxygen gas. Be careful! The sulfur dioxide produced has a very irritating odour and may cause distress to people who suffer from asthma. When you heat sulfur, it melts, turns brown, and burns with a blue flame. It burns more vigorously in oxygen gas than in air. During the burning it combines with oxygen gas to form the compound sulfur dioxide:
sulfur (s) + oxygen (g) --> sulfur dioxide (g)
2. Repeat the experiment using the following: 2.1 Iron (steel wool), 2.2 Magnesium (Do not look at the burning magnesium. The light may injure the eyes.) 2.3 Carbon. In each case, note whether the substance burns more rapidly in oxygen gas than in air.

8.2.1 Heat copper (II) sulfate-5-water crystals, test for water
See 3.80: Exothermic reactions give out heat energy
1. Use anhydrous copper (II) sulfate to test for the presence of water. Heat the crystals gently in a test-tube until they change from blue to white. Water vapour collects on the side of the test-tube. Cool the test-tube. Put some condensed water vapour on the white substance, anhydrous copper (II) sulfate. The copper (II) sulfate turns blue again. This is an example of a reversible change. The return of the blue colour is also a test for water.
(In this direction heat enters the reaction. --->) (<--- In this direction heat leaves the reaction.)
CuSO4.5H2O (s) <--> CuSO4 (s) + 5H2O (l)

2. Heat copper (II) sulfate crystals to make it lose its water of crystallization and leave anhydrous copper (II) sulfate as a white powder. The lost water appears as drops on the inner surface of the upper part of the test-tube. Test the drops for the presence of water with blue cobalt (II) chloride paper. Transfer the anhydrous copper (II) sulfate to another test-tube and add a drop of water. The blue hydrated salt forms again.

3. Put a finger width of copper sulfate in a test-tube. Use a test-tube holder to keep the test-tube horizontal and heat the copper sulfate over the spirit burner flame. To avoid overheating, move the test-tube in the flame or move the flame up and down under the test-tube. Observe the copper sulfate crystals turning white and water condensing on the cooler parts of the test-tube. Repeat the experiment by heating a finger width of copper sulfate crystals in an evaporating basin. Heat the crystals slowly and stir the powder with the glass rod until all the blue colour has just disappeared. Do not heat more because the white powder will darken. Leave the evaporating basin to cool. Divide the white powder into three parts:
3.1 To one part, in a test tube, add one drop of methylated spirit,
3.2 To the second part, in a test tube, add white spirit (dry cleaning fluid, C7 to C12 hydrocarbons),
3.3 To the third part, remaining in the evaporating basin, hold the evaporating basin in the palm of your hand, and add water.
Describe what you see and what you can feel. Only the water turns the powder blue and gives out heat that you can feel in your hand.
Store and label the dry copper sulfate crystals.

8.2.2 Tests for water with cobalt (II) chloride
1. Test for the presence of water with blue cobalt (II) chloride paper. Soak paper in anhydrous cobalt (II) chloride and store in a desiccator. Heat cobalt (II) chloride-6-water crystals. The reaction forms the dark blue anhydrous cobalt (II) chloride with the loss of water. Add water to anhydrous cobalt chloride. The solution becomes pink. Evaporate the pink solution to form purple crystals.
[In this direction, heat enters the reaction. -->]
CoCl2.6H2O (s) [pink] <--> CoCl2 (s) [blue] + 6H2O (l)
[<-- In this direction, heat leaves the reaction.]

2. To add water to the cobalt chloride crystals, grasp the cool test-tube and add water, drop by drop. The blue crystals turn pink, and the test-tube feels hot. When water is added, heat is given out. Cobalt chloride combines with the water and becomes as before being heated. The colour change when adding water is used as tests for the presence of water.

3. Use this chemical to test for the presence of water. Dissolve 5 g of cobalt (II) chloride in 100 mL deionized water. Cut strips of absorbent paper 5 cm × 1 cm and soakthem in the cobalt (II) chloride solution. Remove the strips, drain and spread flat them on trays. Place the trays in an oven at 100oC until the strips are blue. Put strips in a bottle containing dry silica gel (blue in colour) or anhydrous calcium chloride. Keep the bottle sealed, preferably in a desiccator. If the paper turns pink, heat it again as described above until it turns blue again. Do not handle the paper with fingers as moisture from the skin will affect it.

4. Heat pink cobalt chloride crystals. The crystals turn blue and water condenses on the cooler part of the test-tube. Store and label the dry cobalt chloride crystals, CoCl2.

8.2.4 Thermal decomposition of acids
1. Decomposition of boric acid
Boric acid, H3BO3, loses water until, above 170oC, it decomposes to the anhydride, metaboric acid, B2O3, a white, cubic crystalline solid.
H3BO3 --> HBO2 + H2O
Above 300oC, metaboric acid, loses more water and forms tetraboric acid, (pyroboric acid), H2B4O7:
4HBO2 --> H2B4O7 + H2O

2. Decomposition of carbonic acid
Carbonic acid + heat --> carbon dioxide + water
H2CO3 + heat --> CO2 + H2O

3. Decomposition of nitric acid
Nitric acid + heat --> nitric oxide + oxygen + water
(Decomposition occurs at 83 oC, catalysed by light, but below 83oC the colourless pure nitric acid is brown if nitric oxide is already dissolved in it.)
4HNO3 + heat --> 4NO2 + O2 + 2H2O

4. Decomposition of oxalic acid
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.

5. Decomposition of sulfuric acid
Sulfuric acid + heat --> sulfur trioxide + water
H2SO4 + heat --> SO3 + H2O

6. Decomposition of tartaric acid
Heat tartaric acid, (CHOHCOOH)2
Tartaric acid contains the elements carbon, hydrogen, and oxygen, so a residue of carbon is left on heating, and steam forms.

8.2.9 Thermal decomposition of sucrose crystals, C12H22O11
1. Heat sugar on a tin lid or in an old spoon. Note if any gases evolve and if any colours change. Note the residue left after much heating. Steam forms, and a black residue of charcoal (carbon) forms. Sugar is a carbohydrate, a compound of carbon, hydrogen, and oxygen. The last two elements are usually in the ratio of two to one as in water. So when sugar is heated, water as steam forms, leaving a residue of black carbon.
2. Heat sucrose to form carbon. Heat sugar on a metal lid. The substance melts to form a liquid, which soon turns brown. If it is cooled at this stage, the brown solid obtained is called caramel. When heated more strongly the sugar decomposes, giving off inflammable vapours and leaving a black mass of sugar charcoal on the lid.

8.2.9.1 Invert sugar, C12H24O12, HFCS
Prepare invert sugar by heating sucrose + water + cream of tartar or citric acid (fresh lemon juice) to boil for 20 minutes at 114oC in a hydrolysis reaction. It is used in confectionary, sorbets and ice cream as a substitute for corn sugar to prevent crystallization  and so give a smooth mouth feel, hygroscopic humectant to keep products moist,  assists browning caramelization (Maillard reaction), enhances aromas.
C12H22O11 + H2O --> C6H12O6 + C6H12O6
sucrose + water --> glucose + fructose
Invert sugar is 50% glucose + 50% fructose,
HFCS, high fructose corn syrup
HFCS 55, 55% fructose, used in soft drinks
HFCS 42, 42% fructose, used in processed foods

8.2.10 Heat lithium metal to form 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.

8.2.11 Heat aluminium foil to form aluminium oxide
Aluminium foil, (al-foil, alu-foil, "Reynolds wrap'), has thickness usually < 0.2 mm and is shiny on one side and matte on the other side due to the rolling process of manufacture.
Heat a piece of aluminium cooking foil or a "silver" milk bottle top. Describe what happens to the aluminium foil. You may not see any changes because aluminium does not change colour when heated. The melting point is 655oC to 660oC. When white hot,  it slowly forms a coating of aluminium oxide, alumina. Do not heat aluminium powder. If not pure, it may explode.
4Al (s) +3O2 (g) --> 2Al2O3 (s)
aluminium + oxygen --> aluminium oxide

8.2.12 Heat copper foil to form copper (II) oxide
Cleaned copper is brown-red. 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. Heat a narrow strip of copper foil, for half a minute, using the test-tube holder, so that only a small part of the foil is in the flame. Describe what happens to the metal. The metal does not melt. The heated part turns black. The spirit burner flame is not hot enough to melt the copper. The part of the metal in the flame becomes covered with black copper oxide.
2. Tests for copper (II) oxide formation
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.
2Cu + O2 --> 2CuO
copper (s) + oxygen (g) --> copper oxide (s)
3. 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.
4. Show that something is added to the copper from the air. Use a sensitive balance to weigh the copper before and after heating.
5. 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.

8.2.13 Heat iron, nail, wire, filings, to form iron (II) oxide
1. Heat an iron nail. Describe what happens to the metal. The metal does not melt. The heated part turns black. The spirit burner flame is not hot enough to melt the iron. The part of the metal in the flame becomes covered with oxide.
2. Repeat the experiment by heating fine iron wire. Describe what happens to the metal wire. The wire quickly gets red hot and melts. The iron is so thin that it gets hot enough to melt.
3. Repeat the experiment by heating iron filings. Drop a finger width of the iron filings in the spirit burner flame or a Bunsen burner flame. Describe what happens to the iron filings. Some iron filings burn in the flame, like sparklers. Very small particles of iron
become so hot that they burn. These particles combine with oxygen gas very fast to form iron oxide.
Fe + O2 --> FeO
iron + oxygen --> iron (II) oxide

8.2.14 Heat calcium metal to form 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.

8.2.15 Heat sulfur to form sulfur dioxide
Use a test-tube with a stopper to heat sulfur in a fume cupboard or well-ventilated area. When heated to the melting point, sulfur usually ignites and forms sulfur dioxide gas that may distress people suffering from asthma.
1. When you heat sulfur, it melts, turns brown, then burns with a blue flame. Sulfur combines with oxygen gas to form sulfur dioxide.
2. Heat sulfur in air
Be careful! The gases that form have an irritating odour and may cause distress to people who suffer from asthma.
Put sulfur in a combustion spoon and set it alight with a flame. Observe the burning sulfur and then lower it into a test-tube containing oxygen gas. The gases turn moist blue litmus red.
S(s) + O2 (g) ---> SO2 (g)
3. Heat sulfur gently in a crucible in a fume cupboard. The solid melts to form a clear yellow liquid. Melting sulfur without also igniting it is difficult. Ignition is indicated by the formation of a blue flame on the surface, and the formation of acrid sulfur dioxide gas.
4. Put sulfur powder in a large test-tube. Clamp the test-tube in a horizontal position and insert a piece of coiled copper wire 3 cm from the sulfur. Heat the sulfur and copper alternately for five minutes with a strong Bunsen burner flame, with most of the heat on the copper. The sulfur vapours blacken the copper and changes its electrical properties.
5. Put a small quantity of sulfur on a deflagrating spoon and set it alight with a Bunsen burner flame. Note the appearance of the burning sulfur and then lower it into a test-tube containing oxygen gas.
Be careful! The sulfur dioxide produced has a very irritating odour and may cause distress to people who suffer from
asthma. When you heat sulfur, it melts, turns brown and burns with a blue flame. It burns more vigorously in oxygen gas than in air. During the burning it combines with oxygen gas to form the compound sulfur dioxide:
sulfur (s) + oxygen (g) --> sulfur dioxide (g)

8.2.16 Heat magnesium ribbon to form 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 2 800oC because the Mg2+-- O2- chemical bond is stronger than the Mg -- Mg bond.
1. Polish 3 cm of magnesium ribbon with emery paper, then use tongs to hold it in a flame. When the magnesium ignites, hold it out of the flame and over an evaporating basin. Do not look directly at the burning magnesium because it emits a very bright light. Describe the way the metal burns. The magnesium takes fire and burns with a white, dazzling flame, leaving a white ash. Magnesium burns more easily than iron, forming white magnesium oxide, the ash. Do not heat magnesium powder because it may explode!
2Mg (s) + O2 (g) --> 2MgO (s)
magnesium + oxygen --> magnesium oxide
2. The magnesium ash appears lighter than the original magnesium ribbon. To test this observation, weigh a clean dry crucible with lid, add a 15 cm coil of polished magnesium ribbon, then weigh the crucible with lid + magnesium. Use a Bunsen burner flame to heat the crucible on a pipe clay triangle. Occasionally raise the lid slightly with tongs to allow air to enter the crucible. When burning ceases, heat the crucible for a short time without the lid, then leave the crucible to cool with the lid on. Weigh the crucible with lid + ash. The weight of the crucible with lid + ash > weight of crucible with lid + magnesium ribbon, because oxygen from the air had combined with the magnesium to form magnesium oxide.
3. Repeat the experiment with the coil of magnesium covered with 1 cm thickness of clean dry salt, sodium chloride. After heating, the magnesium does not change in weight because the layer of salt prevented the magnesium from contacting oxygen in the air.
4. 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 + O2 --> MgO
5. 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.

13.3.6 Reactions of magnesium oxide
Magnesium oxide is a metallic oxide, and is therefore basic. Magnesium oxide has a low solubility in water. Dilute hydrochloric acid reacts rapidly with aqueous magnesium hydroxide, but slowly with solid magnesium oxide. Magnesium oxide dissolves slowly in water. Phenolphthalein is an indicator that shows changes in alkalinity of the solution. An equilibrium is established between solid magnesium oxide and dissolved magnesium ions. The addition of acid disrupts the equilibrium by removing hydroxide ions from the solution. An equilibrium is established between solid magnesium oxide and dissolved magnesium ions. The addition of acid disrupts the equilibrium by removing hydroxide ions from the solution. Equilibrium is restored by slow dissolving of more magnesium oxide. Addition of larger drops or higher concentration of acid causes a larger initial excess of acid in the solution. Because the reaction of acid with the solid magnesium oxide is slow, it will take a much longer time for the pink colour to return to the mixture. The magnesium oxide formed from combustion of magnesium ribbon forms a hard mass with a small surface area for reaction. The rate of reaction with acid, and the rate of solution of the solid to form an alkaline solution, would be increased by crumbling the ash. Magnesium oxide is used as a component of refractory crucibles, fire bricks, insulation, rubber compounds, magnesia cements and boiler scale compounds, as a reflector in optical instruments and a white colour standard. Magnesium oxide fume, airborne magnesium oxide and magnesia fume is an odourless white opaque smoke. Solid magnesium oxide is a hygroscopic fine white powder. It is slightly soluble in water, pH of a saturated aqueous solution is 10.3. It is slightly soluble in water, and soluble in dilute acids and ammonium salts. It reacts violently with strong acids and halogens. Breathing freshly generated magnesium oxide fume can irritate the eyes and nose

8.3.1 Spirit burner, alcohol lamp, methylated spirit burner, alcohol burner, alcohol lamp
See diagram 3.2.0.1: Spirit burners
1. Use a small jar with a screw-on metal lid. Invert the metal lid on a block of wood and use a large nail to punch a hole in the centre. Note that when you replace the lid on the jar, the projecting metal around the hole will point upward. Make a wick by tearing  away a strip from old cloth. Use a pencil to push the end of the wick through the hole in the lid, starting from the smooth side of the hole. Half fill the jar with methylated spirits. Screw the lid tightly on the jar and light the wick. The metal lid will conduct heat away from the flame so that the flame does not go below the lid into the the methylated spirits.
2. Use a small bottle with a screw metal cap as a simple spirit burner (an alcohol lamp). Punch a hole in the centre of the metal cap. Enlarge the hole so that a metal tube 4 cm long fits into the hole. Push the tube 1 cm into the bottle. Make a wick from cotton waste or a cotton bath towel. Put the wick in the bottle and pull it up through the tube. Fill the bottle with methylated spirit. Make a simple tripod stand with tin snips to cut away the sides of a tin can. The wick should protrude about 3 cm from the cap and fit tightly into the wick holder. The wick holder should fit tightly into the burner. Use only methylated spirit or absolute alcohol (ethanol) as the fuel in the spirit burner.
3. Place the spirit burner on a metal tray or where it cannot be knocked over, i.e. not within "elbow radius" of the user. Keep the  container of methylated spirit stored in another room. To fill the spirit burner, remove the screw cap containing the wick and use a filter funnel to three quarters fill the glass reservoir. Replace the screw cap, screw it down tightly, and wipe the spirit burner dry of methylated spirit. Wash the filter funnel.
4. To put out the spirit burner (extinguish the flame) place a dry test-tube over it so that the rim of the test-tube touches the cap of the spirit burner, or use the glass / ceramic caps are fitted to some spirit burners, to extinguish the flame. The spirit burner flame is almost invisible so be sure that the flame is really extinguished before handling or moving the spirit burner.
5. Students should not be allowed to lift the spirit burner or remove it from the bench. However, they may move the spirit burner by sliding to move it to a safer or more convenient position.
6. Make an alcohol lamp, spirit lamp, from an ink bottle. Use an ink bottle with a screw-on metallic cap, a metallic sheet of 2.5 cm × 4 cm, alcohol, and  a wick made up of wasted cotton or cotton bath towel of length twice the height of the ink bottle. Drill a hole with a nail in the centre of the cap of the bottle. Use a file to enlarge the hole to diameter 10 mm and use a hard round object, e.g. a round file, to burnish the hole. Roll the small metallic sheet into a cylinder. The outer diameter of the cylinder is equal to the inner diameter of the hole on the cap of the bottle. Push the cylinder about 1 cm into the hole on the cap. If possible solderm, the cylinder on the cap, and solder the cracks between the cylinder and the cap. Insert the wick into the cylinder on the cap and leave a part of its length outside of the cap and trim that part well. Fill fuel into the bottle, but not full. Screw the cap on the bottle tightly to prevent evaporation.

8.3.5 Heat calcium sulfate, gypsum
CaSO4.2H2O + heat --> CaSO4H2O + 1 H2O (steam)
gypsum + heat --> calcium sulfate hemihydrate (plaster of Paris, CaSO4.nH2O)

8.4.0 Heat substances that decompose when heated, but may be reformed
When the products of the reaction cannot escape, the reactants and the products remain in contact and their concentrations do not change. A reversible reaction may occur. Then at equilibrium, the rate of the forward reaction = the rate of the reverse reaction, reversible change. This is shown by the arrow symbol <---->.

8.6.0 Combustion, conditions for combustion, to ignite, ignition point
Combustion is the burning, usually in oxygen gas, of a substance releasing heat energy and, sometimes, light energy. Ignition temperature is the temperature at which the substance ignites, e.g. sulfur must reach a temperature of about 400oC before it will burn.
To ignite means to makes something intensly hot by the action of fire, to heat something to the point of combustion or chemical change, to set fire to something, to catch fire and begin to burn, to cause an electric arc, to start combustion in the cylinder of an internal combustion engine.
The ignition point is the temperature at which the rate of reaction is high enough to produce more heat than is lost to its surroundings. When the heat energy accumulates it increases the rate of reaction, so a fire occurs.
1. Put small quantities (that can be put on your little finger nail) of sulfur, magnesium and carbon on a lid of a jam tin. Put the lid on a tripod and heat the centre of the lid with a Bunsen burner flame. Each chemical should receive equal heating. Note the order in which the different substances ignite. In this experiment the order of ignition temperatures should be as follows: sulfur, magnesium, carbon.
2. Repeat the experiment with small quantities of paper, wood and coke (petroleum coke). Heat the centre of the lid and note the order in which any of the materials catch alight. Your ignition temperature order should be as follows: paper, wood, coke.
3. Put some kerosene in a small tin and ignite it with a Bunsen burner flame. With the kerosene still burning, float the tin on a mixture of ice and water. The kerosene stops burning because the ice water mixture removes heat from the burning substance and cools it below its ignition temperature. Firemen use water in the same way to control fire and put out fires.

8.6.1 Burn (Monopoly, fake money) bank notes, ethanol
1. Soak a (Monopoly) bank note, e.g. ten pounds or ten dollars, in 50 mL of water. Use tongs to hold it in the yellow flame of a Bunsen burner. It does not ignite unless all the water evaporates then the paper bank note can reach ignition temperature of about 230oC (Fahrenheit 451o!).
2. Soak a (Monopoly) bank note, e.g. ten pounds or ten dollars, in 50 mL of ethanol. Use tongs to hold it in the yellow flame of a Bunsen burner. It ignites, b.p. 78.5oC.
3. Soak a (Monopoly) bank note, e.g. ten pounds or ten dollars, in a mixture of 25 mL water, 25 mL ethanol and 2 g sodium chloride to colour the flame. Use tongs to hold it in the yellow flame of a Bunsen burner. The ethanol ignites but the paper banknote does not ignite because it is still wet with water.

8.6.2 Oxygen gas is necessary for combustion
1. Ignite a small coil of magnesium wire in a crucible. Pour sand on it while it is still burning. When you cut off the supply of oxygen gas, the burning stops.
2. Play a candle flame oil to the bottom of an evaporating basin. What forms on the basin? The deposit is carbon that has not been burned to gaseous products because not enough oxygen gas was available for complete combustion. You may see similar deposits of carbon or soot when you burn fuels like kerosene, coal and wood with insufficient oxygen.
3. Put some wood shavings and a piece of wood on a metal lid. Heat the lid. Which ignites first? Since oxygen and a solid fuel can interact only at the surface of the solid, the greater the surface area of the solid the more likely the combustion of the solid is to occur. In coal grinding plants the wood shavings that had the greater surface area ignited before the piece of wood. Explosive and spontaneous combustion involving solids occur when the solids are finely divided like a powder and well mixed with air or oxygen.
Explosions have occurred in coal grinding plants and flour mills when the coal dust or flour has been well-mixed with air.
4. Put two lighted candles on a bench. Simultaneously, cover one with a small jar and the other with a larger jar. The candle in the larger jar burns longer.

8.6.3 Carbon dioxide is a product of combustion
1. Put some limewater in a test-tube. Put some carbon on a deflagrating spoon and ignite it. While it is burning, lower it into the test-tube just above the limewater. When burning stops, cover the test-tube and shake it. The limewater now has a milky colour, a test to identify the gas carbon dioxide gas.
2. Repeat the experiment with small quantities of fuels, e.g. wood, coal, and kerosene. All the common fuels are mixtures and contain compounds of carbon.

8.6.5 Respiration is a form of combustion
See diagram 9.155: Respiration of soaked peas
1. Set up flask 1 containing potassium hydroxide, flask two containing limewater, flask 3 containing snails or other small animals, flask 4 containing limewater. Equip each flask with a two-hole stopper. Connect flask 2 to flask 3, and flask 3 to flask 4 with delivery tubes. Connect flask 4 to an air pump and air can enter flask 2 through an open glass tube (not as in the diagram). Notice inlet tubes in each flask reach down to the bottom of the flask. The openings of the outlet tubes are just below the bottom of the stoppers. The air pump draws air through the flasks and through the limewater in flask 2 and flask 4. After some time the limewater in flask 4 turns milky.
2. Repeat the experiment with germinated peas. Both animals and plants produce carbon dioxide as they respire.
3. Put some germinating seeds in a thermos flask and leave a second thermos flask empty. Fit each thermos flask with a cork and a thermometer. Record the temperatures of each flask daily for a few days. The temperature in the thermos flask containing germinated seeds is higher. Heat is produced in respiration. The respiration reaction is exothermic.