Topic 8 Combustion, burning, effect of heat on substances
Updated 2008-03-19 L
Please send comments to: J.Elfick@uq.edu.au
See also: Interesting websites
See also: History of this document
Table of contents
8.1.0 Heat sources, light a match
8.2.0 Elements combine with oxygen when heated in air
8.3.0 Substances that may lose mass when heated
8.4.0 Substances that decompose when heated, but may be reformed
8.6.0 Conditions for combustion, rusting
3.1.0 Bunsen burner
3.29 Collect and weigh the gaseous products of a burning candle
8.1.1.0 Candle, paraffin wax
8.1.1 1 Parts of a candle flame
8.1.1.2 Soot from a candle flame, carbon
8.1.1.3 Describe dark region of flame
8.1.1.4 Test gases from wick
8.1.1.5 Hold aluminium foil below flame
8.1.2.1 Prepare a beeswax candle
8.1.2.2 Burn two candles, burning candle over water
8.1.2.3 Burning candle rocks to and fro
8.1.2.4 Egg in a candle flame
8.1.3.1 Spirit burner, describe spirit burner flame, methylated spirit
2.20 Spirit burner (Primary)
3.42 Things that burn (Primary)  
5.43 Burn to make carbon (Primary)
8.1.4.3 Bunsen burner flame can melt copper wire
8.2.0 Elements combine with oxygen when heated in air
3.28 Substances that gain mass when heated, copper foil
3.28.1 Substances that gain mass when heated, magnesium ribbon
2.29 Collect and weigh the gaseous products of a burning candle
8.2.1.1 Test formation of copper (II) oxide
8.2.3.0 Sulfur
8.2.3.1 Heat sulfur in air
8.3.0 Substances that may lose mass when heated
3.30 Substances that decompose and lose mass when heated
8.3.1 Heat different substances
8.3.2 Heat oxides
8.3.3 Heat nitrates
8.3.4 Heat carbonates
8.3.5 Heat sulfates
8.3.6 Heat manganates
8.4.0 Substances that decompose when heated, but may be reformed
8.4.1 Heat ammonium chloride crystals
8.4.2 Heat limewater (calcium carbonate)
3.31.3 Tests for water with cobalt (II) chloride
8.4.4 Heat copper (II) sulfate-5-water () crystals, test for presence of water
8.4.5 Water of crystallization
8.6.0 Conditions for combustion, rusting
3.52 Conditions necessary for rusting
3.52.1 The mass of iron and its temperature increases during rusting
3.52.2 Oxygen combines with iron during rusting
3.52.3 Metals can prevent rusting
8.6.1 Ignition temperature
8.6.1.1 Burn (Monopoly, fake money) bank notes, ethanol
8.6.2 Oxygen is necessary for combustion
8.6.3 Carbon dioxide is a product of combustion
8.6.5 Respiration is a form of combustion
8.1.0.0 Heat sources
A flame is a region where a gas emits light because of the high temperature. Burning, i.e. combustion, needs oxygen, 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 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 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 a not luminous and occurs because of complete burning of hydrocarbons with plenty of oxygen. The flame does not leave any residue or any other gases. A yellow flame occurs when there is insufficient oxygen. It is a luminous flame. The temperature is lower than the blue flame and leaves black soot and other residues. 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. 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 to burn completely. The flame appears blue and the temperature is very high.
4. 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.
5. Study a match flame. Below the ignition temperature, e.g. white phosphorus 35oC, a combustible substance in oxygen will not catch fire. The wood used in the match has a similar ignition temperature
8.1.1.0 Candle, paraffin wax
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 diffuses in form 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.
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.
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 yellow - orange 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) 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 that burns as follows:
CH4 + 2O2 -> CO2 + 2H2O
A single -CH2- unit burns as follows:
2CH2 + 3O2 -> 2CO2 + 2H2O
8.1.1.1 Parts of a candle flame
See diagram 3.1.4: Bunsen burner flame and candle flame
The candle flame has three parts: 1. The region closest to the wick is dark in colour because air cannot reach that region, so the gases are not burning. 2. The second region is bright yellow orange and forms much light. The 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. 2. The third region, the outer rim of the flame, is practically colourless, a very faint blue, and is the hottest part of the candle flame. The blue colour shows that oxygen is mixing with the wax molecules.
8.1.1.2 Soot from a candle flame, carbon
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.
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.
8.1.1.3 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 tube. These gases have come from the dark region of the flame where there is not enough air to burn them.
8.1.1.4 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 coming from the wick. This shows that the gases from the wick are flammable.
8.1.1.5 Hold aluminium foil below a 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.
8.1.2.1 Prepare a beeswax candle
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. 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.
8.1.2.2 Burn two candles, burning candle over water
See diagram 4.9: Burning candle over water | See diagram 20.1.5: Burning candle over water
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. 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. Cut off each end of a plastic drink 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.
8.1.2.3 Burning candle rocks to and fro
Cut wax away from around the wick at the bottom of the candle so that you have the same length of wick sticking out of each end of the candle. Push a nail or knitting needle through the middle of the candle so that the candle will balance when you place the nail across the sides of two beakers. Put the apparatus in the sink or, to catch candle drips, put a piece of aluminium foil under it if on the table. Simultaneously light both ends of the candle. The burning candle rocks up and down. When you light both ends, one end is sure to burn faster than the other end so it loses more candle wax and becomes lighter than the other end that then tilts downwards. The other end then burns faster, becomes lighter then tilts upwards. The tilted down ends burn faster because the flame becomes closer to the wax. The candle rocks because its centre of gravity, originally through the axis of the nail or needle, moves away from the end burning faster. The centre of gravity continually moves from one side of the axis to the other.
8.1.2.4 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 lookslike 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.
8.1.3.1 Spirit burner, describe spirit burner flame, methylated spirit
See diagram 22.1.1: Simple heating devices
1. 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.
2. Describe the spirit burner flame.
8.1.4.3 Bunsen burner flame can melt copper wire
See diagram 3.1
Bunsen burner flame can melt lead, m.p. = 327oC, and zinc, m.p. = 419oC. Many people think the temperature of a flame cannot exceed 500oC. Copper melts at 1085oC. It is not usually considered possible to melt copper with a Bunsen burner flame. The temperature of different parts of a non-luminous flame, a blue flame, with the air holes fully open, vary. The hottest part of the flame is at the tip of the central cone. The central core of the flame contains a mixture of unburnt gas with air. The intense blue region surrounding the central core is the main zone of combustion, in which the gaseous hydrocarbon fuel reacts with the oxygen, forming short-lived gases. The lighter blue outer flame is where these short-lived gases are completely oxidized to carbon dioxide and water. Copper metal reacts readily with oxygen from air when heated strongly, forming a coating of black copper oxide, CuO. Under reducing conditions, black copper oxide is reduced readily to metallic copper. When heated strongly, but below its melting temperature, copper glows with a bright red heat. To show that the maximum temperature reached, you can melt copper use pieces of copper wire with three different thickness found by stripping the insulation from electrical flex.
1. Light a Bunsen burner, turn the flame to maximum height, and open the air holes so that the flame is completely blue. Hold a piece of thick copper wire with tongs and probe the flame with the wire 1. starting from the bottom 2. around the sides and tip of the central cone, and 3. around the outer blue flame. At each place, record the appearances of the copper, e.g. black, orange, red-hot, or tending to melt.
2. Repeat the process with a thinner piece of copper, then with a very thin piece of copper wire.
3. Reduce the flow of gas and repeat the procedures with a smaller blue flame. The flame has six separate zones: Zone 1 is the core of unburnt gas and air at the base of the flame. Zone 2 is the bright blue region of intense combustion surrounding the core: zone 2A is the tip of the central core, zone 2B is the region at the sides of the central core. Zone 3 is the outer region of the flame where combustion becomes complete: zone 3A is at the top of the flame, zone 3B is the outer part of the sides of the flame. Zone 4 is the region just outside the flame. At each zone observe the appearance of the copper, e.g. black, orange, red-hot, tending to melt, three different thickness of wire. You can melt fine copper wire in the flame but not thick copper wire.
8.2.0.0 Elements combine with oxygen 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. 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 than in air. During the burning it combines with oxygen to form the compound sulfur dioxide:
sulfur(s) + oxygen(g) - -> sulfur dioxide(g)
2. Repeat the experiment using 1. steel wool, iron and 2. magnesium. Do not look at the burning magnesium. The light may injure your eyes. 3. carbon. In each case note whether the substance burns more rapidly in oxygen than in air.
8.2.1.1 Tests formation of copper (II) oxide
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.
8.2.3.0 Sulfur
When you heat sulfur, it melts, turns brown, then burns with a blue flame. Sulfur combines with oxygen to form sulfur dioxide.
8.2.3.1 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. The gases turn moist blue litmus red.
S(s) + O2(g) ---> SO2(g)
8.3.1 Heat different substances
Heat salts for the same period in a crucible and note the results.
8.3.2 Heat oxides
Put small quantities of zinc oxide and copper (II) oxide in separate small test-tubes then heat gently. Note that zinc oxide changes to a yellow colour on heating, but changes back to white on cooling.
8.3.3 Heat nitrates
Repeat the experiment with copper nitrate, lead (II) nitrate, potassium nitrate, and sodium nitrate. The brown gas that forms is nitrogen dioxide. Note that sodium nitrate and potassium nitrate do not give off a brown gas on heating as they only break down to the nitrite and oxygen.
8.3.4 Heat carbonates
Repeat the experiment with basic copper (II) carbonate, calcium carbonate, lead (II) carbonate, magnesium carbonate, and sodium carbonate. Copper (II) carbonate decomposes to release carbon dioxide but sodium carbonate does not decompose on heating.
8.3.5 Heat sulfates
Repeat the experiment with copper (II) sulfate, magnesium sulfate, sodium sulfate, and zinc sulfate. Note that sodium sulfate does not decompose on heating.
8.3.6 Heat manganates
3.30 Substances that decompose and lose mass when heated
Put 0.5 g potassium permanganate in a hard-glass test-tube then weigh it. Fit a loose plug of cotton wool in the mouth of the test-tube to prevent loss of solid during heating. Heat the test-tube. Weigh the test-tube again. A loss of mass occurs because of the decomposition of potassium permanganate with the release of oxygen.
8.4.0 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 <---->
8.4.1 Heat ammonium chloride crystals
Heat ammonium chloride crystals in a test-tube. Hold an open bottle of concentrated aqueous ammonia solution near the mouth of the test-tube - white fumes show that ammonia forms in the reaction.
Let the test-tube cool. Solid ammonium chloride forms again.
NH4Cl(s) <----> NH3(g) + HCl(g)
8.4.2 Heat limewater (calcium carbonate)
See diagram 12.16.3: Heat different carbonates | See diagram 3.34.1: Limewater test for carbon dioxide
Pass carbon dioxide through limewater or blow through it. A milky suspension of calcium carbonate forms. Pass more carbon dioxide through the solution. The solution becomes clear again because soluble calcium hydrogen carbonate forms. Heat the solution again and it becomes clear because insoluble calcium carbonate forms again.
CaCO3(s) + CO2(g) + H2O(l) <---> Ca(HCO3)2(aq.)
8.4.4 Heat copper (II) sulfate-5-water crystals, test for water
See also 3.80: Reactions that give out heat energy, exothermic reactions
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 he reaction. --->) CuSO4.5H2O(s) <---> CuSO4(s) + 5H2O(l) (<--- In this direction heat leaves the reaction.)
8.4.5 Water of crystallization
1. Put 1 g of steel wool in a test-tube and add 10 mL of dilute sulfuric acid. Heat the mixture in a beaker of hot water until all the steel wool has dissolved. Put a lighted paper into the test-tube to show that the gas given off is hydrogen. Filter the solution while hot and allow to cool. If crystals do not form on cooling, add alcohol to cause crystallization. Pour off the liquid and dry the crystals between absorbent paper. Observe the shape of the crystals of green salt iron sulfate.
iron(s) + sulfuric acid(aq) --> hydrogen(g) + iron sulfate(aq)
2. Put some dry crystals prepared in a small test-tube and heat very gently. Water vapour is given off and then condenses in the cool top of the test-tube. The crystals were dry so the water was in the crystals, called the water of crystallization. The shape and character of the crystals changed so water is an integral part of the crystals.
3. Heat separately crystals of sodium chloride, and magnesium sulfate to test whether they contain water of crystallization. and magnesium sulfate contain water of crystallization but sodium chloride does not.
4. Using a dilute solution of cobalt chloride as ink and a matchstick as a pen, write a word on a piece of paper. When this dries the writing becomes almost invisible. Now warm the paper over a small flame to see the word again. In some novelty cards the colour of the word changes with the weather. In fine weather, it is blue and in wet weather it is pink.
CoCl2(s) [blue] + 6H2O(l) <---> CoCl2.6H2O(s) [pink]
5. Leave samples of some of the following crystals exposed to the air, but not to the sun. Examine them each day for a week. The crystals are efflorescent, i.e. they lose water of crystallization and the anhydrous residue on the surface looks like tiny flowers: calcium chloride CaCl2.H2O, sodium sulfate Na2SO4.10H2O (Glauber's salt), sodium carbonate Na2CO3.H2O (washing soda), iron sulfate FeSO4.7H2O slow, copper (II) sulfate (bluestone), calcium sulfate CaSO4.2H2O (gypsum), magnesium sulfate MgSO4.7H2O (Epsom salts), lead acetate (CH3COO)2Pb.3H2O, borax, zinc sulfate sodium thiosulfate (in dry air only).
6. Weigh about 250 g of silica gel on a beam balance. Leave it balanced for some time and note any change in weight. Silica gel is used to keep instruments dry because it absorbs moisture from the air. Silica gel, wool, and biscuits absorb moisture from the air although they do not dissolve in it, called hygroscopic.
8.6.0 Conditions for combustion
Combustion is the burning, usually in oxygen of a substance releasing heat energy and, sometimes, light energy.
8.6.1 Ignition temperature
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. The ignition temperature is the temperature at which the substance ignites, e.g. sulfur must reach a temperature of about 400oC before it will burn. 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. 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.
8.6.1.1 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 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, 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 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. D. 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. 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
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 2-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.