UNChem1

Chemistry
Updated: 2008-08-19 R
Please send comments to: J.Elfick@uq.edu.au

Table of contents
3.1.0 Bunsen burner
3.2.0 Identify pure substances
3.9.0 Solubility and solutions
3.18.0 Separate substances from mixtures
3.28 Substances that gain mass when heated, copper foil
3.28.1 Substances that gain mass when heated, magnesium ribbon
3.28.2 Substances that gain mass when heated, preparation of lithium oxide
3.28.3 Substances that gain mass when heated, preparation of calcium oxide
3.29 Collect and weigh the gaseous products of a burning candle
3.30.0 Substances may decompose and lose mass when heated, thermal decomposition
3.31.0 Hygroscopic, deliquescent and efflorescent chemicals
3.32.0 Prepare, collect and test gases
3.52.0 Rusting
3.53.0 Indicators of acids and bases
3.54.0 Crystal growth
3.55.0 Matter as particles
3.59.0 Electrical conductivity of substances
3.61.0 Construction materials
3.70.0 Chemical reactions
3.80.0 Energy from chemical reactions
3.84.0 Electrical energy from chemical reactions
3.91.0 Rate of reaction
3.95.0 Break down large molecules to small molecules
3.100.0 Building up molecules

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

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

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

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

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

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

3.32.0 Prepare, collect and test gases
3.32 Prepare gases with a gas generation apparatus
3.32.1 Composition of the atmosphere and greenhouse gases
3.33 Prepare ammonia, NH₃
3.33.1 Tests for ammonia, ammonia fountain experiment, ionization reaction of ammonia
3.34 Prepare carbon dioxide, CO₂
3.34.1 Tests for carbon dioxide
3.34.2 Test the breath for carbon dioxide
3.34.3 Solubility of carbon dioxide in water, acidity of soda water
3.34.4 Reduce carbon dioxide with burning magnesium
3.34.5 Frozen carbon dioxide ("dry ice", "hot ice")
3.34.5.1 Dry ice in water
3.34.6 Soda-acid fire extinguisher
3.35 Carbon dioxide in the home
3.35.1 Washing soda
3.35.2 Baking soda
3.35.3 Baking powder
3.35.4 Yeast cells
3.36 Carbon dioxide and photosynthesis
3.37 Carbon dioxide and respiration
3.38 Carbon dioxide and fermentation for brewing
3.39 Carbon monoxide, CO
3.40 Prepare chlorine, Cl₂
3.40.1 Tests for chlorine
3.40.2 Pass chlorine through water
3.41 Prepare hydrogen, H₂
3.41.1 Tests for hydrogen
3.41.2 Prepare hydrogen bubbles
3.41.3 Reduce metal oxides to metals with hydrogen
3.41.4 Reduce copper oxide with natural gas, methane
3.42 Prepare hydrogen chloride, HCl
3.42.1 Tests for hydrogen chloride
3.43.0 Prepare hydrogen sulfide, H₂S
3.43.1 Tests for hydrogen sulfide solution, ionization of hydrogen sulfide
3.43.2 Reduce potassium manganate (VII) with hydrogen sulfide
3.43.3 Reduce iron (III) chloride with hydrogen sulfide
3.44 Prepare nitrogen monoxide (nitric oxide) NO
3.44.1 Catalytic conversion of nitrogen monoxide (nitric oxide)
3.45 Prepare dinitrogen oxide (nitrous oxide) N₂O
3.45.1 Tests for dinitrogen oxide (nitrous oxide)
3.46 Prepare nitrogen, N₂
3.47 Prepare nitrogen dioxide, NO₂
3.47.1 Pass nitrogen dioxide through water
3.48 Acid rain and nitrogen oxides, NO_x
3.49 Prepare oxygen, O₂
3.49.1 Tests for oxygen
3.50 Ozone, O₃
3.51 Prepare sulfur dioxide, SO₂
3.51.1 Tests for sulfur dioxide
3.51.2 Reduce potassium manganate (VII) with sulfur dioxide
3.51.3 Reduce iron (III) chloride with sulfur dioxide
3.51.4 Bleach flowers with sulfur dioxide

3.1.1 Bunsen burner
See diagram 3.1: Bunsen burner flame | See diagram 3.1.1: Bunsen burners
Be careful! Do not turn the gas on without lighting the Bunsen burner. Gas forms an explosive mixture in air.
1. Combustion is the burning in oxygen of a substance to produce heat and sometimes light energy. A flame appears during combustion when a gas has such a high temperature that it emits heat and light. A flame appears only where the burning gas and oxygen are in contact.
2. The Bunsen burner consists of 2.1. a barrel, shaft, 2.2 an air regulator, i.e. a sleeve with a hole in it, 2.3 a jet, air mixture valve, needle valve, 2.4 a base, 2.5. a gas inlet opening.
3. Adjust the flame by opening or closing the gas tap. When the air regulator is open, the gas burns with a noisy blue flame that may be nearly invisible in strong light. If the flame rises up from the burner, turn down the gas supply.
4. When not using the Bunsen burner, either turn off the gas or close the air regulator to give a safety flame. The flame is yellow because of the incandescence of carbon particles. It is not as hot as the blue flame and leaves black soot deposits on glassware.
5. Regularly inspect gas fittings on the benches and hoses connecting Bunsen burners to gas turrets to make sure that connections are free of leaks.
Tests for leaks by dipping the part in soapy water. Be careful! Do not use a lighted match.
6. Heat flammable liquids in water baths using electrical hot plates, not Bunsen burners. Turn the gas off first at the gas tap, then at the cylinder or main supply tap.
7. Use the Bunsen burner only in a draught free area. Allow the Bunsen burner to cool before you move or store it.
8. Do not heat low melting point objects, e.g. plastics, solder, lead, over the barrel of the burner. Melted pieces may fall inside the barrel. Hold the burner at an angle. If a match is blown out, turn gas off, then light the Bunsen burner again.
9. The gas is usually natural gas, i.e. mostly methane CH₄, or LPG, bottled gas, mostly propane, C₃H₈. In a laboratory, the pilot light should burn with a 90% blue flame. If the flame is yellow, the gas may be contaminated with condensates. Do not use such a gas but immediately inform the local gas authority. Previously, laboratories used town gas, based on coal gas, containing equal volumes of methane and hydrogen, some carbon monoxide and hydrogen sulfide as a safety smell. Heating values of fuels: town gas 88 MJ / kg, natural gas 55.6 MJ / kg, LPG gas 49 MJ / kg, diesel fuel oil 38 MJ / L, kerosene 36.7 MJ / L, coke or coal 27 MJ / kg, dry split wood 12.5 MJ / kg.
CH₄(g) + 2O₂(g) --> CO₂(g) + 2H₂O(g) + heat

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

3.28.1 Substances that gain mass when heated, preparation of magnesium oxides
Use magnesium ribbon because magnesium powder is too reactive. Be careful! Do not heat magnesium powder.
Magnesium has density 1.74 g / cm³and melting point 650^oC, but magnesium oxide has density 3.58 g /cm³ and melting point 2800^oC because the Mg²⁺--O^2- chemical bond is stronger than the Mg--Mg bond.
1. Hold a 10 cm strip of magnesium ribbon in a pair of tongs. Place the ribbon in a Bunsen burner flame until it starts to burn. Be careful! Magnesium burns with a very bright white light. Magnesium ribbon corrodes slightly in air and burns with an intense white flame to form a white ash of magnesium oxide.
Mg + 1/2O₂ --> MgO
2. Clean 25 cm of magnesium ribbon and cut into pieces 1 cm long. Put the pieces into a crucible with a lid. Weigh the crucible + lid + contents = W1. Put the crucible on a pipe clay triangle on a tripod stand. Heat gently then strongly. Use tongs to raise the lid. The magnesium darkens before it melts. When the magnesium starts to burn, put the lid back on the crucible and remove the burner. Every few seconds raise the lid slightly to let more air enter. Do not let white magnesium oxide smoke escape. When the magnesium does not burn after you raise the lid, remove the lid and heat the crucible strongly. Hold the lid ready in case the magnesium starts to burn again. Let the crucible cool. Again weigh the crucible + lid + contents = W2. Note W2 > W1. The formation of magnesium oxide causes the increase in weight.

3.28.2 Substances that gain mass when heated, preparation of 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.

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

3.29 Collect and weigh the gaseous products of a burning candle
See diagram 3.29: Gaseous products of burning candle
Candle wax is a mixture of different alkanes (paraffins) saturated hydrocarbons with general formula C_nH_2n+2that are solid at room temperature. Soda lime is a grey-white mixture of sodium hydroxide and calcium hydroxide as granules or powder that absorbs the products of combustion, carbon dioxide and water. Use soda lime instead of sodium hydroxide because soda lime is not deliquescent. Weigh a candle, C1. Weigh a U-tube containing granules of soda lime, U1. Put a candle under an inverted glass filter funnel connected to one arm of the U-tube. Attach a filter pump to the other arm to draw air through the U-tube. Light the candle and turn on the filter pump to draw air over the candle. Let the candle burn for five minutes. Extinguish the candle and disconnect the filter pump. Weigh the candle again, C2. The candle has lost weight, C1-C2. When the U-tube is cool, weigh it again, U2. The U-tube containing the soda lime has gained weight, U2-U1. The U-tube gains more weight than the candle loses weight (U1-U2) > (C2-C1) for two reasons: 1. The candle wax combines with oxygen in the air to form carbon dioxide gas and water. 2. The air sucked in by the filter pump contains some water vapour absorbed by the soda lime. To measure the weight of water absorbed from the air, in a control experiment, repeat the experiment for the same period, but without the candle.

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

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

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

3.30.3 Hydroxides decompose to form water and the metallic oxide

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

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

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

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

3.30.8 Salts with water of crystallization, hydrous salts, lose their water of crystallisation, and become anhydrous powders and then become stable. They also lose their crystalline shape and sometimes their colour.
Prepare test-tubes containing 1 cm of 1. iodine crystals 2. sodium hydrogen carbonate granules or crystals 3. silica sand 4. zinc oxide. Fix a cotton wool plug in the mouth of each test-tube to prevent loss of solid during heating, then weigh each test-tube. Heat each test-tube and cotton wool plug thoroughly and weigh it again. Note any change in weight because of the loss of water of crystallization.
(blue) copper (II) sulfate crystals + heat < = > (white) anhydrous copper (II) sulfate + water.
CuSO₄.5H₂O (s) --> CuSO₄ (s) + 5H₂O (vapour)
blue --> grey white
Na₂CO₃.10H₂0 --> Na₂CO₃ (s) + 10H₂O (vapour)
colourless --> white

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

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

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

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

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

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

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

3.31 Hygroscopic, deliquescent and efflorescent chemicals
1. Hygroscopic substances absorb water from the air.
2. Deliquescent substances are hygroscopic substances that absorb water to such an extent that they form a concentrated solution of the substance. Deliquescent chemicals absorb water from the air and dissolve in it to form a concentrated solution, e.g. citric acid (slight) cobalt (II) nitrate Co(NO₃)₂.6H₂O, magnesium chloride MgCl₂, potassium hydroxide KOH, potassium iodate KIO₃ (slight) potassium iodide KI (slight) sodium nitrate NaNO3 (in moist air) sodium thiosulfate Na₂S₂O₃.5H₂O (in moist air). Store deliquescent chemicals in an airtight container or in a desiccator. When exposed to the air, sodium chloride neither gains nor loses water. Pure NaCl is not hygroscopic. However, sodium chloride as table salt in a salt shaker may become sticky and hard to shake out because it contains deliquescent magnesium chloride as an impurity. Add calcium carbonate or rice grains to table salt to stop it deliquescing.
3. Hygroscopic and deliquescent substances may absorb moisture from tissue and so should be treated as potentially highly corrosive. Hygroscopic chemicals include the following: (anhydrous) calcium chloride CaCl₂, glycerol CH₂OH.CHOH.CH₂OH, iron (II) sulfate-7-water FeSO₄.7H₂O (concentrated) nitric acid HNO₃, potassium carbonate K₂CO₃, potassium chloride KCl (slight) potassium iodide KI (slight) silica gel desiccant SiO₂ (anhydrous) sodium carbonate Na₂CO₃ (concentrated) sodium hydroxide NaOH, sodium nitrite NaNO₂ (anhydrous) sodium sulfate Na₂SO₄ (concentrated) sulfuric acid H₂SO₄. Dry silica gel is hygroscopic. It absorbs water from the air, but does not dissolve in the water.
4. Efflorescent chemicals lose water of crystallization on exposure to the air. Efflorescent chemicals include the following: copper (II) sulfate CuSO₄.5H₂O (in dry air) iron ammonium sulfate Fe(NH₄)₂(SO₄).6H₂O, lead acetate (CH₃COO)₂Pb.3H₂O (slow) magnesium sulfate MgSO₄.7H₂O, di-sodium orthophosphate Na₂HPO₄.12H₂O, sodium sulfate decahydrate Na₂SO₄.10H₂O, sodium tetraborate decahydrate (borax) Na₂B₄O₇.10H₂O (in dry air) zinc sulfate ZnSO₄.7H₂O.

3.31.1 Expose different salts to the air
Put equal volumes of different salts on separate watch glasses, e.g. dry calcium chloride, copper (II) sulfate-5-water, iron (II) sulfate-7-water, potassium carbonate, dry silica gel, pure sodium chloride, and sodium sulfate-10-water. Examine the salts after one hour and after one day.

3.31.2 Expose sodium carbonate decahydrate, washing soda, to the air
The relative molecular mass of sodium carbonate-10-water = 286.15.
The relative molecular mass of anhydrous sodium carbonate = 105.99.
The relative molecular mass of sodium carbonate in sodium carbonate crystals = (105.99/286.15) X 100 = 37%
Open an unopened packet of washing soda. Put 10 g of washing soda in an evaporating basin. Record the time for the crystals to change into a white powder. Leave the basin for two days. Record the time for all the crystals to change into a white powder. Weigh the powder. Calculate the weight of the powder expressed as the percentage of the original weight of the crystals.

3.31.3 Tests for water with cobalt (II) chloride
Tests 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.]

3.32 Prepare gases with a gas generation apparatus
See diagram 3.32: Gas generation apparatus | See also: Saturation vapour pressure over water
Collect less dense gas by downward displacement of air, see diagram 1.
Collect more dense gas by upward displacement of air, see diagram 2.
Collect insoluble gas over water, see diagram 3.
Use a borosilicate test-tube that is not cracked. Clamp the test-tube to a stand. Put the solid reagent in the sidearm test-tube and the liquid reagent in the reservoir. Add the liquid reagent very slowly drop by drop. Keep the reservoir tap closed and the reservoir full to prevent gases blowing back. Grease the stopper and insert it so that if an accidental sudden increase in pressure occurs, the stopper blows out of the test-tube. Use rubber tubing to collect the sidearm to a delivery tube that leads into the receiving test-tube. Discard the first gas coming out of the delivery tube because it is mostly air. Never allow a flame near the gas as it comes out of the delivery tube. Some air probably remains in the receiving test-tube. Use the gas bubbler to collect over water insoluble gases with similar density to air. Some water vapour remains in the receiving test-tube. Gases can also be collected in balloons, inflatable footballs, and plastic bags.

3.32.1 Composition of the atmosphere and greenhouse gases
Gas and percentage volume in dry air: N₂78.08%, O₂ 20.95%, Ar 0.93%, CO₂ 0.03%, Ne 0.0018%, He 0.00052%, Kr 0.00011%, Xe 0.000009%, Rn 6 X 10^-18%. The average formula weight of air is 28.8. The apparent molar mass is 28.96 g / mol. The main greenhouse gases caused by human activities are as follows:
1. Carbon dioxide from burning of fossil fuels, wood and chemical reactions. However, plants remove carbon dioxide from the atmosphere, sequester, during photosynthesis.
2. Methane, CH₄, from coal, natural gas, oil, digestion by herbivores and anaerobic decay of plants in rice paddy and solid waste landfills.
3. Nitrous oxide, N₂O from combustion of fossil fuels and solid wastes and from chemical reactions and agricultural activities.
4. Fluorinated gases, i.e. hydrofluorocarbons, e.g. tetrafluoroethane (CH₂FCF₃, R-134a) perfluorocarbons, e.g. tetrafluoromethane (CF₄, carbon tetrafluoride, R14) and sulfur hexafluoride (SF₆) from chemical reactions. They have high global warming potential but they are not ozone-depleting as are CFCs, e.g. dichlorodifluoromethane (CCl₂F_2,R-12, "Freon-12") HCFCs, e.g. difluoromonochloromethane (CHClF_2, "Freon 22") and halons, e.g. bromochlorodifluoromethane (CF₂ClBr, "Halon 1211").

3.33 Prepare ammonia
See diagram 3.33.1: Preparing ammonia | See diagram 3.33.2: A fountain experiment
Ammonia is an extremely irritating gas and is flammable in the presence of sufficient oxygen. Do not prepare ammonia in an open room. Use a fume cupboard.
Ammonia is less dense than air. Ammonia solution is a weak electrolyte so the properties of the molecules and the ions in the solution affect its properties. Ammonia (NH₃) is produced industrially by the Haber process with a catalyst, with direct synthesis at high pressure and temperature 45^oC. Cloudy ammonia is clear ammonia solution with soap added in memory of the days before the Haber Process when ammonia was made from coal tar and had cloudy impurities.
N₂(g) + 3H₂(g) < = > 2NH₃(g) + energy released
1. Put a mixture of calcium hydroxide and ammonium chloride into a test-tube to a depth of 4 cm. Fill a U-tube with lumps of calcium oxide mixed with cotton wool. The cotton wool is to prevent blocking of the U-tube. Gently heat the test-tube. The calcium oxide is to dry the ammonia gas. Collect the gas by downward displacement of air. Test whether the receiver test-tube is full by holding a piece of moist red litmus paper at the opening. Ammonia gas turns red litmus blue. Collect test-tubes of ammonia gas and apply stoppers.
2NH₄Cl(aq) + Ca(OH)₂(s) --> 2NH₄OH(s) + CaCl₂(aq)
then NH₄OH(s) --> NH₃(g) + H₂O(l)
2. Prepare ammonia with ammonium chloride and sodium carbonate. Put 5 g of ammonium chloride (sal ammoniac) in 2 cm depth sodium carbonate (washing soda) solution. Heat the test-tube. Tests for ammonia gas and with wet red litmus paper.
3. Prepare ammonia with ammonia solution and sodium hydroxide. Add 15 g of granular sodium hydroxide to 30 mL of concentrated ammonia solution contained in a conical flask. Immediately fix in the flask a stopper with a straight delivery tube inserted in it. A large quantity of ammonia forms quickly. Simultaneously, the temperature of the reaction increases and froth seethes inside the flask.

3.33.1 Tests for ammonia, ammonia fountain experiment, ionization reaction of ammonia
1. Concentrated hydrochloric acid test (hydrogen chloride test)
Dip one end of a glass rod into concentrated ammonia solution and one end of another glass rod into concentrated hydrochloric acid. Bring the two ends close to each other but do not let them touch. Blue-white smoke of ammonium chloride forms.
NH₃(g) + HCl(g) -> NH₄Cl(s)
2. Odour test
Tests for ammonia by very cautious smelling. Use very small amounts of reacting chemicals. Do not inhale directly from a test-tube but fan the air above the test-tube towards you.
3. Moist litmus paper test
Ammonia dissolves in water to form a weak base that turns moist red litmus paper blue.
4. Solubility test
Ionization reaction, Kb = 1.8 X 10^-5
NH₃+ H₂O <--> NH₄⁺ + OH^-
Dip the open end of a test-tube containing ammonia under water. The test-tube fills with water.
Ammonia is the most soluble of all gases. Ammonia dissolves in water to form ammonia solution, NH₃(aq). Do not call it "ammonium hydroxide" because while "NH₄⁺" ions and "OH^-" ions can be detected, "NH₄OH" cannot be detected.
5. Ammonia fountain test
Heat the end of a delivery tube and draw it out to form a fine jet. Fill a flask with ammonia and close the flask with a one-hole stopper with a delivery tube. Add litmus to acidified water in a beaker. Warm the flask gently to expand the gas and then hold the flask upside down with the lower end of the delivery tube in the acidified water. Water soon sprays into the flask through the fine jet as the ammonia dissolves in the water and the pressure of ammonia in the flask decreases. The litmus in the water changes from red to blue.
NH₃(g) + H₂O(l) < = > NH₃(aq) + H⁺ + OH^-(aq)
or
NH₃(g) + H₂O(l) < = > NH₄⁺(aq) + OH^-(aq)

3.34 Prepare carbon dioxide
See diagram 3.34: Collecting carbon dioxide, testing when the receiving jar is full
Carbon dioxide gas does not support life so it is a simple asphyxiant. Carbon dioxide and other gases that could accumulate in coal mines to cause choking and suffocation were called choke-damp, after-damp, foul-damp, black damp. Miners used to keep a caged canary bird with them that would die before a concentration of carbon dioxide fatal to humans occurred.
Carbon dioxide is used in photosynthesis. Excess carbon dioxide in the atmosphere from excess burning of fossil fuels causes a greenhouse effect so the temperature of the atmosphere rises, called global warming. An increase of the concentration of carbon dioxide in the atmosphere may increase the rate of photosynthesis.
1. Add dilute hydrochloric acid to 1.1 calcium carbonate (marble chips) 1.2 sodium carbonate (washing soda) 1.3 sodium hydrogen carbonate (baking soda) or basic copper (II) carbonate, CuCO₃.Cu(OH)₂.H₂O. Carbon dioxide is slightly soluble in water so it can be collected over water or by upward displacement of air in dry containers. apply stoppers on the receiving test-tubes to prevent diffusion of the gas into the room.
CaCO₃(s) + 2HCl(aq) --> CaCl₂(aq) + H₂O(l) + CO₂(g)
carbonate + hydrochloric acid --> salt + water + carbon dioxide
2. Test gases from the reaction with a lighted splint.
3. Heat zinc carbonate or basic copper (II) carbonate
CuCO₃.Cu(OH)₂.H₂O --> 2CuO(s) + 2H₂O(l) + CO₂(g)
ZnCO₃(s) --> ZnO(s) + CO₂(g)
4. Add water to sodium carbonate or sodium hydrogen carbonate. No carbon dioxide forms. Add vinegar (acetic acid) or lemon juice (citric acid). The reaction with these acids forms carbon dioxide.
5. Mix vinegar (acetic acid) with sodium hydrogen carbonate in a big container. Drop naphthalene mothballs into the solution. The carbon dioxide formed by the reaction of the vinegar with the sodium hydrogen carbonate forms bubbles of carbon dioxide on the mothballs at the bottom of the container. The mothballs rise to the surface, lose the bubbles and sink again.
2NaHCO₃(s) --> Na₂CO₃(s) + CO₂(g) + H₂O(l)
NaHCO₃(s) + HC₂H₃O₂(aq) --> NaC₂H₃O₂(aq) + CO₂(g) +H₂O(l)
6. Add alum solution (Al₂(SO₄)₃.K₂(SO₄).24H₂O, potash alum) to baking soda or washing soda. The reaction forms carbon dioxide.

3.34.1 Tests for carbon dioxide
1. Lighted splint test
Carbon dioxide extinguishes a lighted splint. Carbon dioxide does not support combustion. Lower a lighted splint into a dry container of carbon dioxide. The level where the flames are extinguished shows the level of carbon dioxide in the container.
2. Limewater test
See diagram 3.34.1: Limewater test
Carbon dioxide turns limewater milky. A fine suspension of calcium carbonate causes the milky colour
Ca(OH)₂(s) + CO₂(g) --> CaCO₃(s) + H₂O(l)
Make limewater by adding calcium oxide (quicklime) to water to form calcium hydroxide.
CaO(s) + H₂O(l) --> Ca(OH)₂(s)
calcium oxide + water --> calcium hydroxide
Then the calcium hydroxide dissolves in water to form a weak alkaline solution. Limewater is a saturated solution of calcium hydroxide.
Ca(OH)₂(aq) < = > Ca²⁺(aq) + 2OH^-(aq)
When testing for the presence of carbon dioxide, make a fresh solution of limewater, otherwise the surface turns milky on standing because of the reaction with the carbon dioxide in the air. Store limewater in a container with a rubber or plastic stopper. If you use a screw-top container, calcium carbonate may form in the screw of the lid so you cannot open the container.
3. Burning charcoal test
Put limewater into a container with a lid. Attach some charcoal to the end of a wire. Ignite some charcoal with a Bunsen burner. Hold the burning charcoal in the container above the surface of the limewater. Remove the burning charcoal. Close the container and shake it. The solution turns a milky colour. The formation of this white solid in limewater is a test for carbon dioxide. No other gas does this.
4. Pouring test
1. Test whether carbon dioxide gas is heavier than air by "pouring" the gas into a test-tube held either above the first test-tube or below it. Use a lighted taper to investigate where the carbon dioxide has gone. 2. Test the density of the carbon dioxide by "pouring" the gas into a container containing a short lighted candle, e.g. a happy birthday candle. The carbon dioxide extinguishes the lighted candle.
5. Litmus test
See also 12.3.0: Properties of acids, ionization of carbonic acid
Carbon dioxide does not change the colour of moist litmus paper. Carbon dioxide dissolves in water to form weak carbonic acid that does not affect moist litmus paper.

3.34.2 Test the breath for carbon dioxide
See diagram 3.34.1: Limewater test
Breathe out through a drinking straw into limewater. The limewater turns milky.

3.34.3 Solubility of carbon dioxide in water, acidity of soda water
See also 35.22.7.1 Calcium carbonate dissolves in rain water
Carbon dioxide is an acidic oxide that dissolves in water to form the weak acid carbonic acid (H₂CO₃) pH about 4, and the carbonate ion. Do not store carbonic acid because it easily decomposes to carbon dioxide and water. Soda water is carbon dioxide dissolved in water under pressure that makes the gas more soluble. Carbonic acid is the basis for all aerated waters, e.g. fizzy lemonade or cola, gaseous natural spring waters and sparkling wines. Carbonic acid soon decomposes, but it can form stable sodium carbonate, potassium carbonate and hydrogen carbonate salts.
1. Open a bottle of soda water or fizzy lemonade. Bubbles of carbon dioxide appear as the gas leaves the solution under the lower atmospheric pressure. Carbon dioxide leaves the solution. Tests for carbon dioxide by putting a lighted splint in the bottle above the lemonade. Test the pH of soda water at room tremperature with drops of methyl red (red below pH 4.2 , yellow above pH 6.3). Boil the soda water and test the pH. Reducing the pressure cause carbon dioxide to come out of solution, equilibrium 1 moves to the left, then equilibrium 3 moves to the left removing hydrogen ions from the solution making the solution less acidic.
Equilibria reactions
CO₂(g) <--> CO₂ (aq) (equilibrium 1)
CO₂(aq) + H₂O(l) <--> H₂CO₃(aq) carbonic acid (equilibrium 2)
H₂CO₃(aq) + OH^-(aq) <--> H₂O(l) + HCO₃(aq)^- (hydrogen carbonate ion, hydrogencarbonate ion) (equilibrium 3)
or
H₂CO₃(aq) <--> H⁺(aq) + HCO₃^-(aq) (hydrogen carbonate ion, hydrogencarbonate ion) (equilibrium 3)
HCO₃^-(aq) + OH^-(aq) <--> H₂O(l) + CO₃^2-(aq) (carbonate ion) (equilibrium 4)
or
HCO₃^-(aq) <--> H+(aq) + CO₃^2-(aq) (carbonate ion) (equilibrium 4)
or
CO₂ + H₂O <--> H₃O⁺ + HCO₃^-
HCO₃^-+ H₂O <--> H₃O⁺ + CO₃^2-

3.34.4 Reduce carbon dioxide with burning magnesium
Attach a small piece of magnesium ribbon to the end of a wire. Light the magnesium ribbon and put it quickly into a test-tube of carbon dioxide. The magnesium continues to burn with a spluttering reaction. White magnesium oxide and specks of black carbon form. The magnesium reduces the carbon dioxide to carbon. If you see no carbon specks, add sulfuric acid to remove the magnesium oxide and unburned magnesium so that the carbon becomes more visible.
2Mg(s) + CO₂(g) --> 2MgO(s) + C(s)

3.34.5 Frozen carbon dioxide ("dry ice", "hot ice")
Be careful! When handling dry ice wear eye protection and wear gloves to avoid burns and frost bite. Store dry ice in an expanded polystyrene box.
If dry ice is touched, the moisture on the skin freezes and the dry ice sticks to the skin. Never lick dry ice because your tongue will stick to it.
When carbon dioxide is cooled under pressure, it becomes a solid called "dry ice" or "hot ice". Dry ice is used for a refrigerant by mobile ice cream sellers and in fire extinguishers. At atmospheric pressure dry ice sublimes at -79^oC. It changes directly from solid to gas. Hold a piece of dry ice in a gloved hand. Watch it disappear as the carbon dioxide sublimes.

3.34.5.1 Dry ice in water
Fill a 10 cc measuring cylinder water and add universal indicator. Add drops of sodium hydroxide solution. Add a lump dry ice. Note how it sinks to the bottom and gives off bubbles of carbon dioxide to make a fog at the mouth of the measuring cylinder. The universal indicator slowly changes colour from blue, pH 9, to orange, pH 5, as the pH reaches about 4.5.
OH^-(aq) + CO₂(g) –> HCO₃^-(aq)
Repeat the experiment with ammonia solution. The colour change of the universal indicator is more gradual because of the reaction of weak acids with weak bases.
H₂O(l) + NH₃(aq) +CO₂(g) –> NH₄⁺(aq) + HCO₃^-(aq)

3.34.6 Soda-acid fire extinguisher
Use a plastic drink-bottle with a one-hole rubber stopper fitted with a plastic tube. Connect rubber tubing with a nozzle to the tube. Use a test-tube that can fit inside the bottle. Partly fill the bottle with sodium hydrogen carbonate solution. Fill the test-tube with dilute sulfuric acid solution and lower it gently into the bottle so that it rests upright. Fit the stopper and plastic tube. Add a detergent to the acid to produce the blanketing effect of foam. Aim the bottle at the fire and invert the bottle rapidly. A strong reaction forms carbon dioxide. The pressure of the gas pushes the liquid out through the jet to extinguish the fire.
2NaHCO₃(aq) + H₂SO₄(l) --> Na₂SO₄(s) + H₂O(l) + CO₂(g)
To make a foam similar to the foam blanket produced by fire extinguishers, add sodium hydrogen carbonate to warm soapy water in a beaker. Add concentrated aluminium sulfate solution and note the mass of white bubbles that looks like ice-cream soda.

3.35 Carbon dioxide in the home
3.35.1 Washing soda
Washing soda (sodium carbonate decahydrate, Na₂CO₃.10H₂O) allows sodium ions to displace calcium ions in clay particles so that clay particles in mud can be dispersed and held in suspension in the washing water.

3.35.2 Baking soda
Baking soda (sodium hydrogen carbonate, bicarbonate of soda, baking powder) is used in cooking to form bubbles of carbon dioxide to expand bread dough, cake mix and pastry dough, to make them light and pleasant to eat. Commercial baking powders often contain a solid acid that reacts with the sodium hydrogen carbonate only when moist.

3.35.3 Baking powder
Baking powder contains sodium hydrogen carbonate (sodium bicarbonate) that reacts with an acid, e.g. 2-hydroxypropanoic acid (lactic acid) from sour milk, to form carbon dioxide. The heat from the oven helps the decomposition of sodium hydrogen carbonate to form carbon dioxide.
baking powder, or sodium bicarbonate, NHCO₃, reacts with an acid such as lactic acid from sour milk to produce carbon dioxide. Commercial "baking powder" often contains a solid acid that reacts with the sodium bicarbonate only when moist, e.g. tartaric acid or hydrogen carbonates.
1. Put baking powder into water and note whether carbon dioxide gas forms. Put sodium bicarbonate into water and note whether carbon dioxide forms. Put baking powder in a test-tube containing vinegar (acetic acid, ethanoic acid) or lemon juice (citric acid) and note whether carbon dioxide forms.

3.35.4 Yeast
Yeast cells convert sugar to carbon dioxide gas and alcohol to make bread rise.
See diagram 3.35: Yeast reacting with sugar solution
1. Make a sugar solution and half fill a container with this solution. Add a spoonful of dry yeast and leave to stand for two days. Construct a bubbler to fit on the top of the container. Note whether the yeast forms a gas. Note whether carbon dioxide gas collects in the upper part of the container. Yeast breaks down sugar into ethanol using enzymes that act as catalysts in the conversion:
C₆H₁₂O₆ --> 2C₂H₅OH + 2CO₂(g)
glucose --> ethanol + carbon dioxide

3.36 Carbon dioxide and photosynthesis
nCO₂ + nH₂A --> (CH₂O)n + nO₂
carbon dioxide + hydrogen donor --> carbohydrate + oxygen
Water is the most common hydrogen donor.
nCO₂ + nH₂O + --> (CH₂O)n + nO₂
carbon dioxide + water (+ light energy) --> carbohydrate + oxygen (dioxygen)
The chlorophyll molecules in green plants absorb mainly red and blue light from the visible range of the electromagnetic radiation from the sun to form higher energy electrons. These excited electrons pass to an electron acceptor to cause a series of reactions resulting in the formation of carbohydrates, e.g. glucose. The electrons removed from the chlorophyll molecules are replaced from the reaction of splitting the water molecule. The protons (H⁺) combine with carbon in the photosynthesis reaction.
2H₂O < = > 2H⁺+ 2OH^- --> 4H⁺+ O₂ + 4e^-Summary equations
6CO₂(g) + 12H₂O(l) + light energy --> C₆H₁₂O₆(aq) + 6O₂(g) + 6H₂O
carbon dioxide + water + light energy --> glucose + oxygen + water (This equation shows water on both sides of the equation.)
6CO₂(g) + 6H₂O(l) + light energy --> C₆H₁₂O₆(aq) + 6O₂(g) (This equation may be preferred because it shows water only on one side of the equation.)

3.37 Carbon dioxide and respiration
Carbon burns to form carbon dioxide. Carbon dioxide is a colourless, odourless gas with a slight smell of soda water, and is about 0.03% of the air. Carbon dioxide is denser than air. Carbon dioxide is slightly soluble in water and the solubility increases with pressure. Carbon dioxide extinguishes a lighted splint.
Fermentation or anaerobic respiration
C₆H₁₂O₆ --> 2C₃H₄O₃+ 4H (combined with other groups)
glucose --> pyruvic acid
Aerobic Respiration
(CH₂O)n + nO₂--> nCO₂+ nH₂O
carbohydrate + oxygen ---> carbon dioxide + water
C₆H₁₂O₆ + 6O₂--> 6CO₂+ 6H₂O
glucose + oxygen ---> carbon dioxide + water + energy

3.38 Carbon dioxide and fermentation for brewing
Carbon dioxide is made in large quantities by the brewing industry. The yeast fungus, Saccharomyces sp. forms enzymes that act as catalysts. Carbon dioxide forms in bread dough, but the fermentation is slower.
Add 5 g of powdered brewer's yeast to 50 mL of 10% sucrose (cane sugar) solution or molasses or treacle. Collect the carbon dioxide over water. After leaving the fermentation for 2 days in a warm place the smell of alcohol is obvious.
invertase enzyme C₁₂H₂₂O₁₁+ H₂O ---> C₆H₁₂O₆+ C₆H₁₂O₆
sucrose + water ---> (+)glucose + fructose
zymase enzyme C₆H₁₂O₆---> 2C₂H₅OH + 2CO₂
(+)glucose ---> ethyl alcohol + carbon dioxide

3.39 Carbon monoxide, CO
Be careful! Do NOT make carbon monoxide.
See also 18.6.3: Danger of vehicle exhausts, tailpipe gases
Carbon monoxide is very toxic. It can cause unconsciousness because of combination of the gas with haemoglobin in the blood. Death can occur from carbon monoxide inhalation. Do not prepare carbon monoxide. Metal oxides are reduced by passing carbon monoxide over the heated oxide. Carbon monoxide is very poisonous and particularly dangerous because it is colourless and has no smell. It kills more people than any other gas. Carbon monoxide is poisonous because it reacts with the haemoglobin in blood and prevents the blood from acting as an oxygen carrier. The gas can form accidentally by leaving a car engine running in a closed garage or by burning a gas fire with restricted ventilation. When carbon or carbon compounds burn in a limited supply of air, the reaction forms carbon monoxide.
2C(s) + O₂(g) --> 2CO(g)
Carbon monoxide is insoluble in water, but it is absorbed by potassium hydroxide solution. Carbon monoxide burns with a pale blue flame forming carbon dioxide.
2CO(g) + O₂(g) --> 2CO₂(g)
Carbon monoxide can act as a reducing agent and is the main reducing agent in a blast furnace. At high temperatures, carbon monoxide reduces the oxides of copper, lead and iron to the metal.
CuO(s) + CO(g) --> Cu(s) + CO₂(g)
Fe₂O₃(s) + 3CO(g) --> 2Fe(s) + 3CO₂(g)

3.40 Prepare chlorine, Cl₂See also: 1.13a Simple fume hood
Chlorine gas is very toxic. Can react to cause fires or explosions upon contact with turpentine, ether, ammonia gas, illuminating gas, hydrocarbon, hydrogen and powdered metals. Dissolves readily in water forming highly corrosive solution. Do not prepare chlorine in open room. Use fume cupboard. Direct combination of chlorine and hydrogen in bright light or ignition of the mixture by lighted taper or electric spark. Reactions of chlorine with metals, solid non-metals, hydrocarbon. Use small quantities only.
Fume cupboards, fume chambers, fume hoods,
Chlorine is a greenish yellow gas with an irritating and choking odour. Be careful! Chlorine gas is poisonous and damages the respiratory organs. Do not inhale gases directly from the test-tube. Fan the gas towards the nose with the hand and sniff cautiously. If no odour is detected, move closer and try again. Prepare chlorine with bleaching powder, bleach solution. Bleaching powder is a mixture of calcium chloride, calcium hydroxide and calcium chlorate (I). Bleaching powder is made commercially by the reaction of chlorine with solid calcium hydroxide. Do the following experiments in a fume cupboard, fume hood or near an open window. Before doing these experiments, make available sodium thiosulfate or calcium hydroxide solution to be used for a chlorine trap to absorb excess chlorine gas. Also prepare ammonia solution because the effect of inhaling chlorine gas may be counteracted by inhaling ammonia vapour. The best treatment for inhaling chlorine gas is plenty of fresh air.
1. With great care, warm bleaching powder and smell it until you notice a choking smell because of chlorine gas being produced by the action of carbon dioxide in the air. Test with wet red or blue litmus paper that becomes colourless because of the bleaching action of chlorine.
2. Put 5 g of bleaching powder (calcium hypochlorite) into a test-tube. Add drops of a weak acid, e.g. citric acid or vinegar. Test with wet red or blue litmus paper. Hold a piece of white paper behind the apparatus to note the green chlorine gas.
3. Add dilute sulfuric acid to bleaching powder. After collecting a small amount of chlorine gas put a stopper in the receiving test-tube and put the end of the delivery tube into sodium thiosulfate solution to absorb excess chlorine.
Bleaching powder + H₂SO₄(aq) --> CaSO₄(s) + H₂O(l) + Cl₂(g)
4. Domestic bleach is manufactured by mixing a solution of chlorine with sodium hydroxide solution
Cl₂(g) + 2OH^-(aq) --> Cl^-(aq) + ClO^-(aq) + H₂O
Add a dilute acid to bleach solution to form chlorine gas.
NaOCl(aq) + HCl(aq) --> NaCl(aq) + H₂O(l) + Cl₂(g)

3.40.1 Tests for chlorine
1. Bleaching test
Chlorine bleaches moist red or blue litmus paper, flowers and some dyes in cloth.
2. Lighted splint test
Chlorine extinguishes a lighted splint, but hot steel wool burns in it.

3.40.2 Pass chlorine through water
Chlorine is available commercially for school laboratory use as chlorine water. Hypochlorous acid HClO, a bleach and a disinfectant, is a solution of chlorine (I) oxide that forms salts called hypochlorites. Hypochlorous acid is a weak acid that easily decomposes back to chlorine gas and water. When chlorine passes through water, a mixture of HCl and HClO forms. The chlorine is oxidized and reduced.
Cl₂(g) + H₂O(l) < = > HCl(aq) + HClO(aq)

3.41 Prepare hydrogen
See diagram 3.41: Collecting hydrogen gas
Do not allow direct combination of hydrogen and chlorine in bright light or ignition of the mixture by lighted taper or electric spark. You can ignite a jet of hydrogen issuing from a delivery tube. Hydrogen reduces metal oxides.
Hydrogen, H₂, is a colourless odourless diatomic gas with the lowest density of any element. Hydrogen does not change the colour of moist litmus. The hydrogen ion, H⁺,is a proton.
1. Zinc with hydrochloric acid
Do not use a container bigger than a test-tube. Put granulated zinc in a test-tube and cover it with water. Add a crystal of copper (II) sulfate to act as a catalyst. Slowly add dilute hydrochloric acid through a funnel, as in diagram 2.41.1 or through a syringe, as in 2.41.2. Bubbles of hydrogen appear on the surface of the zinc. The test-tube feels hot because the reaction is exothermic. Collect hydrogen gas by downward displacement or over water. Let the reaction continue for some minutes to drive out all the air from the test-tube. Discard the first two test-tubes of hydrogen because they will contain displaced air. Collect test-tubes of the gas and apply stoppers.
Zn(s) + 2HCl(aq) --> ZnCl₂(s) + H₂(g)
2. Iron with sulfuric acid or citric acid, or sodium hydrogen sulfate
Put 1 cm depth of iron filings in a test-tube. Just cover the iron filings with a dilute acid solution. Warm the test-tube until frothing starts. Hydrogen is colourless and odourless but any impurities in the iron filings give a nasty smell.
3. Aluminium with sodium carbonate
Cut into small pieces some aluminium foil or aluminium milk bottle tops and put into a test-tube. Add 5 mL of sodium carbonate solution (Na₂CO₃.10H₂O, washing soda). Heat until effervescence occurs.
4. Iron with alum
Put 5 g of iron filings in a 1 cm depth of alum solution in a test-tube. Heat the solution until effervescence occurs. [Potash alum, "alum" has the formula Al₂(SO₄)₃.K₂(SO₄).24H₂O and is also shown as KAl(SO₄)₂.12H₂O.]
5. Iron with ammonium chloride
Put an equal volumes mixture of iron filings and ammonium chloride in a dry test-tube and heat. Hydrogen gas and ammonia are given off.
6. Calcium with hydrochloric acid
Use forceps to transfer about 0.1 g of calcium metal turnings to dilute hydrochoric acid in a test-tube.
Ca(s) + 2HCl(aq) --> CaCl₂(aq) + H₂(g)

3.41.1 Tests for hydrogen
1. Be careful! A dangerous explosion may occur if you use anything bigger than a small test-tube when igniting the gas, particularly if the gas is mixed with air. Never test more than a test-tube full of hydrogen gas. Never dry hydrogen gas with concentrated sulfuric acid.
Hold a lighted splint or burning taper to the mouth of a test-tube. The gas explodes with a squeaky pop sound. The splint is extinguished. The squeaky pop shows rapid combustion of hydrogen to form water vapour. Look for vapour on the sides of the test-tube. However, as 2 litres of gas forms only about 1 mL of liquid, the liquid on the sides of the test-tube may just show that test-tube was already wet before the experiment.
2H₂(g) + O₂(g) --> 2H₂O(l)
2. Hydrogen does not change the colour of moist litmus.
3. Pouring test
Test whether hydrogen is lighter than air by "pouring" the gas into a test-tube held either above the first test-tube or below it. Use a lighted taper to investigate where the hydrogen has gone.

3.41.2 Prepare hydrogen bubbles
Hydrogen is much lighter than air and was formerly used in airships, dirigible balloons. It has now been replaced by helium because hydrogen ignites easily.
Pass hydrogen through soapy water to form soap bubbles full of hydrogen. Shake the bubbles gently to make them float up. The hydrogen bubbles rise into the air, showing the low relative density of hydrogen gas. Try to ignite the bubbles with a lighted splint.

3.41.3 Reduce metal oxides to metals with hydrogen
See diagram 3.41.3: Hydrogen over heated copper oxide
Be careful! Use a safety screen and wear eye protection
1. Pass hydrogen over 5 g of copper (II) oxide (CuO, black copper oxide) or lead (II) oxide (lead monoxide, PbO, lithage) or iron (III) oxide (haematite, Fe₂O₃). Hydrogen reduces metal oxides to metals. The products are the metal and water.
Weigh a reduction tube empty then with copper oxide. Pass hydrogen over the copper oxide and light the gas as it comes out of the hole in the end of the combustion tube. Heat the copper oxide with a Bunsen burner flame until it glows then turns pink. The glow shows that reduction occurs. Remove the Bunsen burner. Let the combustion tube cool then discontinue the supply of hydrogen. When the flame has gone out remove the stopper and weigh the reduction tube and contents again.
CuO(s) + H₂(g) --> Cu(s) + H₂O(l)
In the industrial process, blistered copper is heated in a furnace and natural gas is passed through the molten copper oxide until the flame burns green to indicate that almost pure copper remains.
2. Repeat the experiment with 5 g of copper (I) oxide (red copper oxide, Cu₂O)

3.41.4 Reduce copper oxide with natural gas, methane
1. Pass natural gas, about 95% methane, over heated copper (II) oxide powder. The reduction reaction is slow and may need twenty minutes of strong heating. The copper does not glow with heating so it is not clear when all the copper oxide has been reduced.
4CuO(s) + CH₄(g) --> 4Cu(s) + 2H₂O(l) + CO₂(g)
2. See also: Metaldehyde
Repeat the experiment with a 1 cm cubic piece of metaldehyde in the reduction tube. The reduction is quicker.
3. Repeat the experiment with natural gas that has bubbled through ethanol. The reduction is quicker and a slight glow is seen as the copper oxide is reduced.
6CuO(s) + C₂H₅OH(l) --> 6Cu(s) + 3H₂O(l) + 2CO₂(g)

3.42 Prepare hydrogen chloride
See diagram 3.42: Collecting hydrogen chloride | See also: 1.13a Simple fume hood
Hydrogen chloride gas is corrosive. Do not prepare hydrogen chloride in an open room. Use fume cupboard.
Be careful! Do this experiment in a fume cupboard, fume hood. Hydrogen chloride gas has a choking odour because it combines with the water vapour in the air to form hydrochloric acid. Concentrated sulfuric acid reacts with metal chlorides to form hydrogen chloride that dissolves in water to form hydrochloric acid.
Put sodium chloride crystals in a 100 mL filter flask or sidearm test-tube. Coarse rock salt causes less frothing than the fine salt. Carefully add concentrated sulfuric acid down a funnel to just cover the sodium chloride crystals. Heat the mixture if necessary. Collect the hydrogen chloride gas in test-tubes by upward displacement of air then put a stopper in the receiving test-tube and put the end of the delivery tube into water to absorb excess hydrogen chloride.
NaCl(s) + H₂SO₄(aq) --> HCl(g) + NaHSO₄(aq)

3.42.1 Tests for hydrogen chloride
1. Solubility test
1. Remove the stopper from one receiving test-tube under water. Note the solubility of hydrogen chloride. 2. Invert a receiving test-tube over water. The gas dissolves immediately to form hydrochloric acid. The water rises almost to the top because collection by upward displacement of air results in some residual air remaining in the test-tube.
2. Moist litmus paper test
Test the solution in the receiving test-tube with moist litmus paper. Red litmus paper turns blue.
3. Ammonium chloride test
Hold a piece of cotton wool soaked in ammonia solution at the mouth of a bottle of hydrochloric acid. Note the white cloud of ammonium chloride.
4. Lighted splint test
Hydrogen chloride extinguishes a lighted splint. Hydrogen chloride neither burns nor supports combustion.
5. Magnesium ribbon test
Shake a receiving test-tube with water to form a solution of hydrogen chloride, hydrochloric acid. Put a piece of magnesium ribbon in the solution. Collect any gas formed and test for hydrogen with the glowing splint test.
6. Ammonia solution test
Hold a piece of cotton wool soaked in ammonia solution, NH₃(aq) ("ammonium hydroxide") at the mouth of a receiving test-tube and note the white cloud of ammonium chloride above the hydrochloric acid.
7. Hydrogen chloride fountain test
This test is similar to the ammonia fountain test. Heat the end of a delivery tube and draw it out to form a fine jet. Fill a flask with hydrogen chloride and close the flask with a one-hole stopper with a delivery tube. Add litmus to alkaline water in a beaker. Warm the flask gently to expand the gas and then hold the flask upside down with the lower end of the delivery tube in the alkaline water. Water soon sprays into the flask through the fine jet as the hydrogen chloride gas dissolves in the water and the pressure of hydrogen chloride in the flask decreases. The litmus in the water changes from blue to red.

3.43.0 Prepare hydrogen sulfide
See also: 1.13a Simple fume hood
Hydrogen sulfide gas is both an irritant and an asphyxiant. Do not prepare hydrogen sulfide in an open room. Use fume cupboard. You can ignite a jet of hydrogen sulfide issuing from a delivery tube.
Be careful! Hydrogen sulfide is an extremely poisonous colourless flammable gas with an unpleasant smell of rotten eggs. At less than 1% concentration the smell disappears. So a student may be breathing in this poisonous gas without being aware of it. Do NOT use a Kipp's apparatus for generating hydrogen sulfide.
1. Do this experiment in a fume cupboard, fume hood. Put 5 sodium thiosulfate crystals in a metal screw cap. Heat the metal screw cap gently by holding it with pincers in a Bunsen burner flame until the crystals have melted and solidified again, with steam given off.
Be careful! Do NOT inhale gas directly from the metal screw cap.
With more careful heating, note the "rotten egg" smell of hydrogen sulfide. Allow the metal screw cap to cool. Moisten the white residue with a weak acid, e.g. vinegar. The smell of hydrogen sulfide gas becomes stronger. Dip a strip of clean newspaper in the copper (II) sulfate solution and hold it over the meal screw cap. The paper turns black.
2. Do this experiment in a fume cupboard, fume hood. Add dilute hydrochloric acid to iron sulfide. Collect the gas over warm water by downward displacement.
FeS(s) + 2HCl(aq) --> FeCl₂(aq) + H₂S(g)
Ignite the gas as it leaves the delivery tube.
2H₂S_(g) + 3O_2(g)--> 2SO₂(g) + 2H₂O(l)

3.43.1 Tests for hydrogen sulfide solution, ionization of hydrogen sulfide
Be careful! The gas is soluble in water, so use a solution of hydrogen sulfide in water instead of the gas.
1. Odour test
Hydrogen sulfide has the odour of rotten eggs.
Be careful! Do NOT inhale gases directly from the test-tube. Fan the gas towards the nose with the hand and sniff cautiously. If you detect no odour, move closer and try again.
2. Lead (II) nitrate test
Hydrogen sulfide solution turns lead (II) nitrate solution test paper black.
3. Litmus test
Hydrogen sulfide solution turns blue litmus slightly red-pink.
4. Copper (II) sulfate test
Hydrogen sulfide solution turns copper (II) sulfate solution black.
Ionization of hydrogen sulfide
H₂S + H₂O --> H₃O⁺ + HS^-
HS^- + H₂O --> H₃O⁺ + S^2-

3.43.2 Reduce potassium manganate (VII) with hydrogen sulfide
See also: 1.13a Simple fume hood
Do this experiment in a fume cupboard, fume hood. Pass hydrogen sulfide through a dilute acidified potassium manganate (VII) solution. The colour of the manganate ion is lost and a milky precipitate of sulfur forms.
2MnO₄^-(aq) + 6H⁺(aq) + 5H₂S(g) --> 2Mn²⁺(aq) + 8H₂O(l) + 5S(s)

3.43.3 Reduce iron (III) chloride with hydrogen sulfide
Hydrogen sulfide reduces yellow acidified iron (III) chloride to green Fe²⁺ with precipitation of sulfur.
Add sodium hydroxide to the filtered precipitate to form a green-brown precipitate of iron (II) hydroxide.