topic08

School Science Lessons
Topic 8 Heat sources, candle, spirit burner, Bunsen burner, heat chemicals, combustion
Updated 2009-09-17
Please send comments to: J.Elfick@uq.edu.au
See: Interesting websites

Table of contents
8.1.0 Heat sources
8.1.1 Candle
8.1.2 Spirit burner, alcohol lamp
8.1.3 Bunsen burner
8.2.0a Heat chemicals, linked experiments
8.2.0 Elements combine with oxygen gas 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 and ignition temperature

8.1.0 Heat sources
3.42 Things that burn (Primary)
5.43 Burn to make carbon (Primary)

8.1.1 Candle
2.44 Candle flame (Primary)
6.35 Burn candle over water, candle burning in inverted jar over water (Primary)
6.36 Candle "lava" (primary)
6.36.1 Cooling candle wax
3.39 Convection heat snake
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 Dark region of a candle flame
8.1.1.4 Test gases from the wick
8.1.1.5 Aluminium foil below candle flame
8.1.1.6 Floating tea candle
8.1.2.1 Prepare beeswax candles
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

8.1.2 Spirit burner, alcohol lamp
2.20 Spirit burner (alcohol lamp) (Primary)
8.1.3.1 Spirit burner, alcohol lamp, methylated spirit

8.1.3 Bunsen burner
3.1.0 Bunsen burner
8.1.4.1 Bunsen burner, flame
8.1.4.2 Bunsen burner, flame, which part is hottest
8.1.4.3 Bunsen burner, flame can melt copper wire
13.2.1 Bunsen burner, flowing air can do work, application of Bernoulli's law
2.11 Bunsen burner, "Gas-pak "
13.2.1 Flowing air can do work, application of Bernoulli's law

8.2.0a Heat chemicals, linked experiments
8.2.01 Heat copper sulfate crystals
8.2.02 Heat cobalt chloride crystals
8.2.03 Cobalt chloride invisible ink (CoCl₂.6H₂O)
8.2.04 Lemon juice invisible ink (citric acid, COOHCH₂C(OH)COOHCH₂COOH.H₂O, C₆H₈O₇.H₂O)
8.2.05 Heat copper carbonate CuCO₃.Cu(OH)₂.H₂O, basic copper carbonate
8.2.06 Heat crystals to find water of crystallization
8.2.07 Heat cane sugar (sucrose, C₁₂H₂₂O₁₁)
8.2.08 Heat tartaric acid [(CHOHCOOH)₂,]
8.2.09 Heat iron sulfate crystals (FeSO₄.7H₂O)
8.2.10 Heat ammonium chloride

8.2.0 Elements combine with oxygen gas 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
8.2.11 Heat aluminium foil
8.2.12 Heat copper foil
8.2.13 Heat iron nail
8.2.14 Heat magnesium ribbon
8.2.15 Heat sulfur

8.3.0 Substances that may lose mass when heated
3.30 Substances that decompose and lose mass when heated
8.3.1 Heat salts
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 water
8.4.5 Water of crystallization

8.6.0 Conditions for combustion and ignition temperature
3.52 Conditions necessary for rusting
3.52.1 The mass of iron and its temperature increases during rusting
3.52.2 Oxygen gas combines with iron during rusting
3.52.3 Metals can prevent rusting
8.6.1 Burn (Monopoly, fake money) bank notes, ethanol
8.6.2 Oxygen gas 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 Heat sources
See 16.1.1cc LPG (liquefied petroleum gas, LP gas)
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 a not luminous and occurs because of complete burning of hydrocarbons with plenty of oxygen gas. The 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 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 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.
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 35^oC, a combustible substance in oxygen gas 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 gas 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 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 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 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

8.1.1.1 Parts of a candle flame
See diagram 3.1.4: Bunsen burner flame and candle flame
Hold a piece of white cardboard behind the flame so that you can see each part of the clearly.
The candle flame has three parts, regions. each region has the shape of the flame around the wick.
1. The innermost region closest to the wick consists of vapours from the molten wax and is dark in colour because air cannot reach that region, so the gases are not burning.
2. The second region is bright yellow orange to blue near the bottom 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 800^oC. The outer orange and yellow areas are hotter than this region. Some combustion occurs in this region.
3. 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. 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 so intermediate substances form. Tiny black specks in this region are particle of carbon (soot) that glow on ignition and emit most of the light from the candle.
C + O₂ --> CO₂ + light energy

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 glass 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 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.

8.1.1.5 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.

8.1.1.6 Floating tea candle
A tea candle is 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 the maintain the burning of the vapour.

8.1.2.1 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 about 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.

8.1.2.2 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. Add ink or cochineal to colour the water. 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 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.

8.1.3.1 Spirit burner, alcohol lamp, methylated spirit
See diagram 8.1.3.1: Spirit burners
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. The wick should protrude about 3 cm from the cap and fit tightly into the wick holder. 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 sully container of methylated spirit store in another room.
4. 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.
5. 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.
6. 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.
7. 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; a wick made up of wasted cotton or cotton bath towel of length more than two times of 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 some hard round object (for example round file) to burnish the hole. Roll the small metallic sheet into a cylinder. The outer diameter of the cylinder is equal to 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 solder the cylinder on the cap; even the cracks between the cylinder and the cap also are soldered tightly. Insert the wick into the cylinder on the cap and leave a part of fit length outside of the cap and trim the part well. Fill fuel into the bottle but not too full. Screw the cap on the bottle tightly to prevent evaporation.

8.2.01 Heat copper sulfate crystals
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. Move the test-tube in the flame or move the flame up and down under the test-tube so that overheating does not occur. 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:
* to one part in a test tube, add one drop of methylated spirit,
* to one part in a test tube add white spirit (dry-cleaning fluid, C7 to C12 hydrocarbons),
* to the third part remaining in the evaporating basin, hold it 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.

8.2.02 Heat cobalt chloride crystals
Repeat the experiment by heating pink cobalt chloride crystals. The cobalt chloride turns blue and water condenses on the cooler part of the test-tube.
Store and label the dry cobalt chloride crystals.
Add water to the cobalt chloride crystals.
Grasp the cool test-tube upright in your hand and add water, drop by drop. Describe what you see. The blue residue turns pink, and the test-tube becomes hot. When water is added to these substances, a chemical reaction occurs and heat is given out. The substances combine with the water and become as they were before being heated. The effect of adding water is used as tests for the presence of water.

8.2.03 Cobalt chloride invisible ink (CoCl₂.6H₂O)
Make a weak solution of cobalt chloride by adding one or two crystals to half a test-tube of water and shaking it. The solution should be very pale pink, almost colourless. Using a pen with a clean nib and containing no ink, or an old fashioned dip pen, write a message with the invisible ink you have made, and allow the writing to dry. If the solution was weak enough, your writing will be invisible. Heat the paper but not over a flame. Note whether the writing now shows. Breathe on the visible writing. Describe what you see. Heat the paper to make the writing blue. Breathing on the blue writing makes it invisible again.

8.2.04 Lemon juice invisible ink (citric acid, COOHCH₂C(OH)COOHCH₂COOH.H₂O, C₆H₈O₇.H₂O)
Write with the lemon juice and heat the paper as before. Note the colour of the writing. It is brown.

8.2.05 Heat copper carbonate CuCO₃.Cu(OH)₂.H₂O, basic copper carbonate
Use the copper carbonate from 10.01.2
Put a finger width of copper carbonate in a test-tube and heat it until the colour changes. Leave to cool. The colour changes to black. The copper carbonate decomposes into black copper oxide and the invisible gas, carbon dioxide. Most carbonates similarly split on heating.

8.2.06 Heat crystals to find water of crystallization
Heat alum ( AlK(SO₄)₂.12H₂O, aluminium potassium sulfate crystals), Epsom salts (MgSO₄.7H₂O, magnesium sulfate crystals) , and household salt (Nacl,
sodium chloride crystals). Decide which substances contain water of crystallization. Alum and magnesium sulfate contain water of crystallization. Although most salts crystallize from their solutions as hydrates with water of crystallization, household salt (sodium chloride crystals), does not form water of crystallization.

8.2.07 Heat iron sulfate crystals (FeSO₄.7H₂O)
Heat a few of the pale green crystals in a test-tube until they turn white. Note any other changes. Water condenses on the cooler. The iron sulfate has lost its water of crystallization and become white, anhydrous iron sulfate.

8.2.08 Heat ammonium chloride
Place ammonium chloride in the test-tube and heat, only at the bottom of the test-tube at first. Heat more strongly by holding the bottom of the test-tube quite still in the flame. Describe what you see. The ammonium chloride partly vaporizes and this vapour turns back to solid chloride again, forming a white deposit higher up the test-tube. The solid. ammonium chloride, on heating, turns directly into a gas without first melting into a liquid. When the gas is cooled, it turns directly back into a solid again. A substance that behaves in this way is said to sublime.

8.2.13 Heat cane sugar (sucrose, C₁₂H₂₂O₁₁)
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.

8.2.14 Heat tartaric acid [(CHOHCOOH)₂]
Repeat the previous experiment with tartaric acid. Tartaric acid contains the elements carbon, hydrogen, and oxygen, so a residue of carbon is left on heating. and steam forms.

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.

8.2.11 Heat aluminium foil
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 seeany changes because aluminium does not change color when heated. The melting point is 655oC to 660^oC. 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
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)

8.2.13 Heat 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.
Repeat the experiment by heating fine iron wire
Heat a small piece of fine 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.
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 oxide

8.2.14 Heat magnesium ribbon
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.
NEVER 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 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.

8.2.15 Heat sulfur
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)

8.3.1 Heat salts
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 and lead (II) nitrate. The brown gas that forms is nitrogen dioxide. Sodium nitrate and potassium nitrate do not give off a brown gas on heating because they only breakdown to the nitrite and oxygen gas.
2KNO3 (s) --> 2KNO2 (s) + O2 (g)
potassium nitrate --> potassium nitrite + 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 gas.

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 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 gas. Filter the solution while hot and leave 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. Observe 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, 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 and ignition temperature
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 400^oC before it will burn
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. 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 230^oC (Fahrenheit 451^o!).
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. 78.5^oC.
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. 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 two-holes 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.