School Science Lessons
Topic 13 Gases, prepare gases, atmosphere and greenhouse gases
2012-02-04 SP
Please send comments to: J.Elfick@uq.edu.au
See: Interesting websites
Table of contents
13.0.0 Gases
13.1.0 Gases (different gases, ammonia to xenon)
Gases, (household gas, laboratory gas)
3.1.1 Bunsen burner safety
3.4.8 Chemical vapours and smelling chemicals, safety
13.1.6 Chlorine gas, Cl2
13.1.7 Chlorides, Cl-
3.32.1 Composition of the atmosphere and greenhouse gases
1.0.0 Density of gases, SVP, Saturation Vapour Pressure
2.4 Fume cupboards, fume hoods
2.11 Gas-pak
17.1.3 Gas burette, Dilute hydrochloric acid with marble chips
13.01 Gas bags
3.32.0 Gas generation apparatus
2.1 Gas installations and inspections
3.4.6 Gas or vapour inhalation, EAR, CPR, safety
3.25 Gases, Separate gases dissolved in a water sample
3.8.0 Hazards associated with gases
19.4.4.22 Packaging gases, propellants, food additives
3.32.0 Prepare gases with a gas generation apparatus
3.39.1 Reactions of methane with steam
3.41.4 Reduce copper oxide with natural gas, methane
13.1.5a Relative molecular mass of gases
20.0.6 Standard temperature and pressure, STP, density of gases
7.7.13.1 Volume of gas dissolved in tap water
3.2.0 Water of crystallization
13.1.0 Gases, different gases, ammonia to xenon
3.33.0 Ammonia
3.33.1 Ammonia, Tests for ammonia
3.33.2 Ammonium compounds
13.1.0 Argon, Ar
13.1.2 Bromine, Br
13.1.4 Butane, C4H10
3.34.0 Carbon dioxide
3.35.0 Carbon dioxide in the home
3.34.1.0 Carbon dioxide, Tests for carbon dioxide
13.1.5 Carbon disulfide, CS2, Toxic by all routes, Not permitted in schools
3.39.0 Carbon monoxide, CO
13.1.6 Chlorine, Cl2
13.4.3 Chlorine, Tests for chlorine
8.7.21.0 Chlorine, Tests for chlorine, swimming pool chemistry
13.1.7 Chlorides, Cl-
13.1.9 Ethane, C2H6
13.1.10 Ethene, (ethylene), C2H4
13.1.11 Ethyne, (acetylene), C2H2
13.1.12 Fluorine, F
13.1.13 Helium, He
13.1.14 Heptane, C7H16
13.1.15 Hexane, C6H14
13.1.17 Hydrogen bromide, HBr
3.42.0 Hydrogen chloride, HCl
3.42.1 Hydrogen chloride, Tests for hydrogen chloride
3.41.0 Hydrogen gas, H2
3.43.0 Hydrogen sulfide, H2S
3.43.1 Hydrogen sulfide, Tests for hydrogen sulfide
13.1.17 Krypton, Kr, Table of the elements
13.1.19 Methane, CH4
13.1.21 Neon, Ne
13.1.22 Nitrogen compounds
13.1.23 Nitrogen dioxide, NO2, nitrogen (IV) oxide, dinitrogen tetroxide
13.1.24 Nitrogen gas, N2
13.1.25 Nitrogen monoxide, (nitric oxide), NO
13.1.26 Nitrogen oxides, Acid rain and nitrogen oxides, NOx
13.1.22.1 Nitrogen triiodide
13.1.27 Nitrous oxide, (dinitrogen monoxide), N2O
13.1.28 Octane, C8H18
13.1.29 Oxygen gas, O2
12.17.0 Oxides, Reactions of oxides
13.1.31 Ozone, O3, Highly toxic gas with pungent odour:
13.1.32 Pentane, C5H12
13.1.33 Propane, C3H8
13.1.34 Radon, Rn, Table of the
13.1.35 Sulfur dioxide, SO2
13.1.37 Trichloroethane,  (1,1,1-trichloroethane, methyl chloroform), CH3CCl3
13.1.38 Trichloromethane, (chloroform), CHCl3
13.1.40 Xenon, Xe, Table of the elements
3.34.1.0 Tests for carbon  dioxide
Burning charcoal test for carbon dioxide: 3.34.1.3
Lighted splint test for carbon dioxide: 3.34.1.1
Limewater test for carbon dioxide: 3.34.1.2
Litmus test for carbon dioxide: 3.34.1.5
Pouring test for carbon dioxide: 3.34.1.4
Test carbon dioxide as a greenhouse gas: 3.38.1
Tests for carbon dioxide in the breath with limewater: 3.34.2
Tests for carbon dioxide with bromothymol blue: 6.5.13
Tests for carbon dioxide with limewater: 9.154
Tests for carbon dioxide with thymolphthalein indicator: 3.34.1.6
Tests for carbon dioxide with phenol red: 3.34.1.7
3.41.0 Hydrogen gas, H2, Table of the elements
3.99 Gases from wood
3.41a Hydrogen gas
7.2.2.19 Hydrogen gas properties
3.41.1 Hydrogen gas, Tests for hydrogen gas
3.2.0 Hydrogen bonding, in liquids
3.8.5 Hydrogen gas, hazards
33.3.8 Hydrogen / oxygen fuel cell
12.2.5 Acid-base reactions
12.7.3 Alkalis with metals, sodium hydroxide, See 2.
15.1.8 Anodize aluminium, See 3.
3.94 Catalysts and rate of reaction (hydrogen gas forms)
7.1.0 Chemical changes and physical changes, formation of water
13.4.0 Chlorine (Direct combination with hydrogen)
Coal gas
3.7.1 Concentration and rate of reaction
12.1.3 Conservation of mass in a cycle of copper reactions, the copper cycle experiment, (Step 5. hydrogen bubbles)
15.3.2 Corrosion of magnesium
17.1.1 Count bubbles, dilute hydrochloric acid with granulated zinc
12.3.2 Dilute acids with metals, hydrochloric acid
12.3.2.1 Dilute acids with metals, sulfuric acid, hydrochloric acid, ethanoic acid
12.3.3.1 Dilute acids with metals, sulfuric acid with aluminium
12.3.3 Dilute acids with metals, sulfuric acid with iron
8.0 Direct union of elements to form compounds, hydrogen gas with chlorine gas, hydrogen gas with sulfur
12.12.10 Drain cleaners
15.7.0 Electrode potential of metals, (standard hydrogen cell)
15.5.01 Electrolysis experiments
15.5.3 Electrolysis of water, conduction of water:
15.5.4 Electrolysis of water (decomposition of water) Hofmann electrolysis apparatus
15.5.5 Electrolysis of water, measure volume of hydrogen gas generated
15.5.6 Electrolysis of water, Decomposition of acidified water by electricity
33.1.0 Electrolysis, electrolytes, anode and cathode
3.41.02 Explode a hydrogen balloon
7.9.22 Flammable (hydrogen) see 4.
7.9.28 Fuel cell (hydrogen gas forms)
33.3.8 Fuel cell, hydrogen / oxygen fuel cell
12.4 Hydrochloric acid (forms hydrogen gas)
3.41.01 Hydrogen generator from a boiling tube
4.41 Ice experiments (Hydrogen boy), See 3.
12.3.0.5 Ionization reaction of carbonic acid
12.3.3.3 Iron with sodium hydrogen sulfate
33.3.4 Lemon cell, See 3.
3.41.1 Lighted splint tests for hydrogen gas
3.41.1.1 Litmus tests for hydrogen gas
3.74 Metals displace hydrogen from acids (hydrogen bubbles)
12.15.1 Metals with water, Cu, Zn, Fe, Mg, Al
12.2.4.3 Magnesium displaces hydrogen in ethanoic acid
12.3.3.2 Magnesium with sodium hydrogen sulfate
12.15.3 Metals with steam
16.1.1a Methane, CH4, prepare methane gas
12.3.11.1 Nitric acid with metals
12.3.0.4 pH
18.1.0 pH tests
19.4.4.22 Packaging gases, propellants, food additive, See E949 Hydrogen gas
3.41.1.2 Pouring test for hydrogen gas
3.41.0 Prepare hydrogen gas, H2
3.41.2 Prepare hydrogen gas bubbles
16.1.3.1.2 Prepare sodium ethoxide
12.15.1 Prepare silica and silicon, See 4.
7.2.2.1 Reactions of aluminium
12.01.1 Reactions of aluminium (1. hydrogen gas forms)
7.2.2.12 Reactions of chlorine
7.2.2.21 Reactions of iron
7.2.2.24 Reactions of lithium
12.15 Reactions of metals with water, See 6.
7.2.2.30 Reactions of nickel
7.2.2.32 Reactions of oxygen gas
12.18.5.4 Reactions of dilute sulfuric acid as an acid
12.19.9.1 Reactions of bromine, Br2, See 1.
12.19.9.4 Reactions of hydrogen bromide, HBr, See 7.
3.73 Reactions of sodium with water (hydrogen bubbles)
7.2.2.48 Reactions of zinc
3.41.3 Reduce metal oxides to metals with hydrogen gas
3.41.4 Reduce copper oxide with natural gas, methane
12.2.6 Redox reactions
5.1.5 Relative atomic mass of magnesium (hydrogen gas forms)
3.84.3 Simple electric cell, copper with zinc in dilute sulfuric acid (hydrogen bubbles)
33.3.1 Simple electric cell
12.11.3.5 Tests for substances with dilute hydrochloric acid, note gas produced
14.1.10 Test heat of reaction, potassium with diethyl ether
8.4.5 Water of crystallization, See 1
3.41.1 Tests for hydrogen gas
3.41.02 Explode a hydrogen balloon
3.41.1 Lighted splint test for hydrogen gas
3.41.1.1 Litmus test for hydrogen gas
3.41.1.2 Pouring test for hydrogen gas
15.5.12 Tests for hydrogen gas, Electrolysis of sodium chloride solution (See 1. and 2.)
13.1.5 Carbon disulfide, CS2, carbon bisulfide, Toxic by all routes, Not permitted in schools
Flammable: 7.9.22 (See: 2.)
Flammable organic chemicals: 15.7.0
Flammable liquids: 3.5.8
Prepare rayon, copper (II) sulfate with ammonia solution, "regenerated fibre", "artificial silk: 3.4.
13.1.6 Chlorine gas, Cl2
11.0  Activated carbon [commercial information]
13.4.4  Bleaching powder
13.4.8  Burn substances in chlorine
Chloric acid, HClO3
Chlorine, Cl, Table of the elements
13.4.0  Chlorine
15.2.2  Chlorine as an oxidizing agent
7.2.2.12  Chlorine, descriptions of common elements
3.8.4  Chlorine, hazards
Chlorine monoxide, dichlorine oxide, Cl2O
Chlorine dioxide, ClO2
Chlorine reagent test kit
18a2 Chlorine used in swimming pools
12.19.5.0  Chlorofluorocarbons, CFCs
8.0 Direct union of elements to form compounds, chlorine with sodium, chlorine with hydrogen gas
18.7.6  Dissolve chlorine in pool water by electrolysis
3.4.4  Electric writing
13.4.6.1 Halogenation of benzene
7.1.3  Heat metals in chlorine
18.7.4  Measure the free chlorine in water
16.1.1a  Methane with chlorine
13.4.12  Pass chlorine through iodine solution
13.4.13  Pass chlorine through iron (II) chloride solution
13.4.9  Pass chlorine through water
13.4.1  Prepare chlorine
13.4.2  Prepare chlorine water
13.4.15  Reactions of chlorine with alkalis, bleaching powder
13.4.6  Reaction of chlorine with benzene, dichlorbenzene, C4H4Cl2
13.4.14  Reactions of chlorine with heated copper and steel tool
13.4.7  Reactions of chlorine with sodium
13.4.3  Tests for chlorine
18.7.21.0  Tests for chlorine, swimming pool chemistry
12.11.5.8  Tests for chlorides  12.11.5.8
15.5.12 Tests for chlorine, Electrolysis of sodium chloride solution (See 4.)
13.1.7 Chlorides, Cl-
3.7.3 Chlorides, Cl-, hazards
18.2.2.1 Chlorides in groundwater
Chlorinated lime, calcium hypochlorite
12.19.6.4 Compare silver chloride, silver bromide and silver iodide
12.19.8.4 Concentrated sulfuric acid with potassium chlorate, KClO3, (forms chlorine dioxide)
Hydrochloric acid
3.42.0  Hydrogen chloride, HCl
3.42.1  Hydrogen chloride, Tests for hydrogen chloride
3.7.7 Hypochlorites, hazards
12.19.6.4 Compare silver chloride, silver bromide and silver iodide
1.9 List of chlorides
12.2.1 Precipitation reactions, double decomposition
12.19.8.3 Prepare iron (IlI) chloride, FeCl3
12.19.8.5 Prepare potassium perchlorate by fractional crystallization
12.19.8.1 Reactions of chlorides, Cl-
12.19.8.6 Recover silver from silver chloride, AgCl2
3.43.3 Reduce iron (III) chloride with hydrogen sulfide
3.51.3 Reduce iron (III) chloride with sulfur dioxide
12.11.5.8 Tests for chlorides
13.1.29 Oxygen gas, O2, Table of the elements
Oxygen gas, compressed, cylinders for oxy-acetylene welding
Oxygen liquid, refrigerated oxygen, with flammable or combustibles may explode, causes frostbite (cold burns), use ventilated gloves
15.2.1 Oxygen as oxidizing agent
12.12 Oxygen experiments
33.3.8 Oxygen fuel cell, hydrogen / oxygen fuel cell
7.2.2.32 Oxygen, properties
3.48 Acid rain and nitrogen oxides, NOx
3.1.1 Bunsen burner
3.1.1.1 Bunsen burner gases
3.1.4 Bunsen burner flame and candle flame
13.3.4 Burn magnesium ribbon in oxygen
13.3.3 Burn steel wool in oxygen, burn iron filings
13.3.2 Burn sulfur in oxygen
3.36 Carbon dioxide and photosynthesis
3.39.0 Carbon monoxide, properties
3.94 Catalysts and rate of reaction
7.1.0 Chemical changes and physical changes
7.1.1 Chemical changes, burn magnesium
3.44.1 Catalytic conversion of nitric oxide (nitrogen monoxide)
3.28.4 Collect and weigh the gaseous products of a burning candle
12.17.2.2 Competition for oxygen, Heat metals with oxides of another metal
3.32.1 Composition of the atmosphere and greenhouse gases
8.6.0 Conditions for combustion and ignition temperature
12.2.3.1 Decomposition reactions
12.3.11.0 Dilute nitric acid with copper
18.3.2 Dissolved oxygen in water, Winkler method
3.98 Elements in foods
7.2.2 Elements and chemical reactions
8.2.0 Elements combine with oxygen gas when heated in air
7.2.2.21 Fe, Iron
8.3.3 Heat nitrates
8.3.6 Heat manganates
8.1.0 Heat sources
12.11.3.4 Heat substances in a dry test-tube, See 5.
3.49a Hydrogen peroxide concentration
3.87 Lead accumulator cell, lead-acid rechargeable battery, See 3. 4
7.2.2.24 Li, Lithium
3.52.1 Mass of iron and its temperature increases during rusting
13.1.6a Molar volume of oxygen prepared with hydrogen peroxide
15.3.4 Need for oxygen for rusting
15.3.5 Need for oxygen for corrosion of magnesium
16.4.6 Oxyacetylene welding, Tests for gases from burning hydrocarbons, oxyacetylene welding
3.52.2 Oxygen gas combines with iron during rusting
9.237 Oxygen content of inhaled and exhaled air
19.4.4.22 Packaging gases, propellants, food additive
3.49.0 Prepare oxygen gas with hydrogen peroxide, O2
12.12.7 Prepare oxygen with hydrogen peroxide using catalysts
13.3.1 Prepare oxygen foam with hydrogen peroxide
13.3.5 Prepare oxygen absorbent
7.2.0 Pure substances and impure substances
23.4.2 Reactions in liquid oxygen
3.73 Reactions of sodium with water
12.2.6 Redox reactions, (oxidation-reduction reactions, electron transfer reactions)
18.3.0 Tests for air and dissolved oxygen in water, dissolved oxygen, DO
3.49.1 Tests for oxygen gas, glowing splint test
15.5.12 Tests for oxygen gas, Electrolysis of sodium chloride solution (See 1. and 3.)
13.1.24 Nitrogen gas, N2, asphyxiant at high concentrations, liquid nitrogen, LN2, burns skin
Use eye and face protection when using compressed gas.
Do not use in small enclosed spaces with limited ventilation.
Liquefied nitrogen gas can cause frostbite or cold "burns" so wear gloves and eye / face protection
Nitrogen gas, N2, Table of the elements
Ammonia and the ammonium ion, NH3,NH4+: 12.11.3
Chemical fertilizers: 6.9.17.0
Nitrogen properties: 7.2.2.31
Nitrogen experiments: 12.11
Nitrogen: 13.9.0
Nitrogen gas generated in a motor car airbag: 13.9.3
Prepare nitrogen using ammonium nitrite: 3.46
Properties of nitrogen: 7.2.2.31
Reactions of the nitrites, NO2-: 12.11.1
Reactions of the nitrates, NO3-: 12.11.2
13.1.22 Nitrogen compounds
Nitrogen compounds
Acid rain and nitrogen oxides, NOx: 3.48
Nitrate, nitrite
Nitrates, List of nitrates: 1.19
Nitric acid
Nitric oxide (nitrogen monoxide), NO: 3.44.0
Nitriles, -CN, RC = -N, (cyano group, cyanides, cyanide ion, CN-), acid nitriles, alkyl cyanides: 16.2.4.2
Nitrites, (NO2-, dioxonitrate ion, salts or esters of nitrous acid, O=NOH), Nitrites group: (-C=N), suffix: -nitrite: 16.2.4.5
Nitro, -NO2
Nitroalkanes, (nitroparaffins), (CnH2n+1NO2): 16.2.4.4
Nitrobenzene, oil of mirhan, aromatic nitro compounds: 16.3.4.0.2
Nitrogen compounds, one atom of nitrogen: 16.2.4
Nitrogen compounds, two or more nitrogen atoms: 16.2.5
Nitrogen gas generated in a motor car airbag, sodium azide, NaN3: 13.9.3
Nitrogen oxides, Acid rain and nitrogen oxides: NOx
Nitrogen compounds, one atom of nitrogen, 16.2.4
Nitrogen compounds, two or more nitrogen atoms: 16.2.5
13.1.22.1 Nitrogen triiodide, triiodoamine, NI3, touch powder, ammonium triiodide, contact explosive, Explosive if dry, Not permitted in schools
Nitroglycerine (UK), nitroglycerin (USA), Explosive, Not permitted in schools: 16.1.3.0.3
Packaging gases, propellants, food additive: 19.4.4.22
13.1.23 Nitrogen dioxide, NO2, nitrogen (IV) oxide, dinitrogen tetroxide, Highly toxic gas that forms acids deep in the lungs
Nitrogen dioxide, < 0.1% Not hazardous but use cross-ventilation
Acid rain and nitrogen oxides, NOx: 3.48
Conservation of mass in a cycle of copper reactions, the copper cycle experiment, Step 1: 12.1.3
Heat nitrogen tetroxide, (dinitrogen tetroxide, N2O4): 17.5.6.2
Pass nitrogen dioxide through water: 3.47.1
Prepare nitrogen dioxide, NO2, nitrogen (IV) oxide, copper with nitric acid: 3.47
Prepare nitrogen dioxide, NO2, concentrated nitric acid with copper: 13.10.2
13.1.25 Nitrogen monoxide, (nitric oxide), NO
3.44.1 Catalytic conversion of nitric oxide (nitrogen monoxide)
3.44 Prepare nitrogen monoxide (nitric oxide) NO
13.1.27 Nitrous oxide, (dinitrogen monoxide), N2O, dinitrogen oxide, laughing gas, non-toxic but may cause hysteria, use fume cupboard
3.32.1 Composition of the atmosphere and greenhouse gases, (See 3)
19.4.4.22 Packaging gases, propellants, food additives
3.45 Prepare dinitrogen oxide (nitrous oxide) N2O
3.45.1 Tests for dinitrogen oxide (nitrous oxide)
3.32.0 Prepare gases with a gas generation apparatus
See diagram 3.32: Gas generation apparatus | See: SVP, Saturation vapour pressure over water
Diagram 1. Collect more dense gas by upward displacement of air, if molecular mass > 28.8.
Diagram 2. Collect less dense gas by downward displacement of air, if molecular mass < 28.8.
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.
Diagram 3. Collect insoluble gas or not very soluble gas of any density over water, i.e. by downward displacement of water.
Fill one third of the water trough. Fill a test-tube with water by placing it it on its side in the water trough. Put your thumb over the end of the test-tube and invert it. Fix the end of the gas delivery tube inside or under the test-tube. If the gas is slightly soluble in water its solubility will be less in warmer water, Some gases dissolve in water to produce heat and form an acid solution, e.g. HCl, SO2, NO, NO2. Some gases dissolve in water to form a basic solution, e.g. NH3.
Diagram 4. Collect soluble gases in water (aqueous solution), e.g. Cl2
Diagram 5. When the gas preparation equipment uses downward displacement of water, (Diagram 3.), or collection in water, (Diagram 4.), water may be forced back into the equipment by atmospheric pressure, i.e. "sucked back", and break hot glassware or dilute reactants. To prevent sucking back used an inverted glass filter funnel.
3.32.1 Composition of the atmosphere and greenhouse gases
Gas and percentage volume in dry air: N2 78.08%, 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 molecular mass air molecules is 28.8 (80% of 28 + 20% of 32). The apparent molar mass is 28.96 g / mol.
The main greenhouse gases are as follows:
1. Water vapour from evaporation of water or sublimation of ice
2. Carbon dioxide, an acidic oxide, from burning of fossil fuels, wood and chemical reactions. However, plants remove carbon dioxide from the atmosphere, sequester, during photosynthesis so concentration drops during the Northern hemisphere growing season. Carbon dioxide transmits visible light but absorbs infrared radiation.
3. Methane, CH4, from volcanoes, coal, natural gas, oil, digestion by herbivores and anaerobic decay of plants in rice paddy and solid waste landfills.
4. Nitrous oxide, N2O from combustion of fossil fuels and solid wastes and from chemical reactions and agricultural activities including emission by tropical soils.
5. Ozone, O3, concentrated in the ozone layer of the atmosphere, shields the earth from excess high frequency ultraviolet light. However, it harms the respiratory system.
6. Chlorofluorocarbons, CFCs, contain C, Cl and C, and hydrofluorocarbons, HCFCs, that also contain H, 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 (CCl2F2, R-12, "Freon-12") HCFCs, e.g. difluoromonochloromethane (CHClF2, "Freon 22") and halons, e.g. bromochlorodifluoromethane (CF2ClBr, "Halon 1211").
3.34.1.1 Lighted splint tests for carbon dioxide
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.
3.34.1.2 Limewater tests for carbon dioxide
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)
Prepare 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.
3.34.1.3 Burning charcoal tests for carbon dioxide
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.
3.34.1.4 Pouring tests for carbon dioxide
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.
3.34.1.5 Litmus tests for carbon dioxide
(See 12.3.0): Properties of acids
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.
H2O (l) <--> H+ (aq) + OH- (aq)
2H+ (aq) + CO32- (aq) <--> H2CO3 (aq) carbonic acid
CO2 + H2O <--> H3O+ + HCO3-
HCO3- + H2O <--> H3O+ + CO32-
3.34.1.6 Tests for carbon dioxide with thymolphthalein indicator
See diagram 3.34.1: Thymolphthalein test
Thymolphthalein, C28H30O4, acid-base indicator, pH 9.4 colourless, pH 10.6 blue, Quantity of indicator per 10 mL: 3.1
Put 125 cc of ethanol in a beaker and add 5 drops of thymolphthalein indicator. Add drops of dilute sodium hydroxide solution until the solution turns blue. Blow through a tube into the solution until it becomes colourless.
CO2 (g) + H2O (l) --> H2CO3 (aq) <--> H+ (aq) + HCO3- (aq)
CO2 (g) + 2NaOH (aq) + CO2 (g) --> Na2CO3 (aq) + H2O (l)
The sodium hydroxide is added to make the solution slightly alkaline at the beginning of the experiment and to absorb any initial carbon dioxide or any other acid. Na2CO3 is less basic than NaOH.
3.34.1.7 Tests for carbon dioxide with phenol red:
See diagram 3.34.1: Phenol red test
See: Phenol red indicator
Put 125 cc of ethanol in a beaker and add 2 drops of phenol red  indicator. Add drops of dilute sodium hydroxide solution until the solution turns red. Blow through a tube into the solution until it becomes yellow.
CO2 (g) + H2O (l) --> H2CO3 (aq) <--> H+ (aq) + HCO3- (aq)
CO2 (g) + 2NaOH (aq) + CO2 (g) --> Na2CO3 (aq) + H2O (l)
The sodium hydroxide is added to make the solution slightly alkaline at the beginning of the experiment and to absorb any initial carbon dioxide or any other acid. Na2CO3 is less basic than NaOH.
3.35.0 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.
Baking powder: 19.1.6.1
3.39.0 Carbon monoxide, properties
Be careful! Do NOT make carbon monoxide.
See 18.6.3: Danger of vehicle exhausts, tailpipe gases
1. Air pollution
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 + oxygen gas --> carbon monoxide
2. 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)
3. 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)
4. Use fume cupboard to reduce metallic oxides to the metal by passing carbon monoxide over the heated oxide.
3.39.1 Reactions of methane with steam
At 700oC and nickel catalyst forms hydrogen and carbon monoxide.
CH4 (g) + H2O (g) --> 3H2 (g) + CO (g)
3.41a Hydrogen gas, H2
Hydrogen gas, compressed gas and gas generated by experiment, Highly flammable gas, can act as non-toxic asphyxiant
Hydrogen forms spontaneously combustible / explosive mixture in air at low temperature
Students may be injured by "popping" (burning) hydrogen in glass containers. Generate small quantities of gas in a test-tube only
Hydrogen gas forms explosive mixtures with air when mixed in almost any proportions. Pure hydrogen burns in air with a hot colourless flame without exploding. However, mixtures of hydrogen and oxygen or air burn with explosions. In experiments to demonstrate ignition of hydrogen, air / hydrogen mixtures are often ignited instead. For example, when igniting a jet of hydrogen gas from a flask containing metal and acid, sufficient air may be left in the flask to form an explosive mixture so the whole flask explodes throwing shards of glass in all directions. Exploding about 1 mL of hydrogen / air or hydrogen / oxygen mixture is safe if there is no possibility of broken materials being propelled outwards by the explosion. Exploding soap bubbles containing hydrogen / oxygen mixture on the surface of a beaker of water is safe, if the delivery tube is at least 10 cm below the surface of the water. There may be some risk associated with popping air / hydrogen mixtures in test-tubes because the test-tube may shatter. Check that the test-tubes are not flawed. Ensure that all ignition of hydrogen occurs at least 2 metres away from any other source of hydrogen and away from glass light fittings or glass-fronted cabinets. Wear eye protection when doing experiments with hydrogen. Prepare hydrogen gas by reaction of an acid on an active metal, e.g. zinc with hydrochloric acid.
3.41.0 Prepare 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. Prepare hydrogen gas by reaction of an acid on an active metal, e.g. zinc with hydrochloric acid.
1. 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.
2. Aluminium foil and sodium carbonate, washing soda
Cut into small pieces some aluminium foil or aluminium milk bottle top and put into a test-tube. Add 5 mL sodium carbonate solution (Na2CO3.10H2O, washing soda). Heat until effervescence.
3. Aluminium kitchen foil with caustic soda drain cleaner (sodium hydroxide)
Fill a clear glass bottle, e.g. a wine bottle, one third full with tap water. Do not use a plastic bottle. Cut a plastic bottle with diameter greater than the glass bottle at half its length to make a beaker. Half fill the plastic beaker and stand the glass bottle in it. With safety glasses and gloves, use a funnel to add three heaped teaspoons of  caustic soda to the water in the glass bottle. Keep the bottle open and swirl the bottle gently without spilling any reagents then return the bottle to the plastic beaker. The  dissolving is exothermic so the contents of the bottle will get hot. Roll three 20 X 30 cm sheets of kitchen aluminium foil into cylinders and then drop them into the bottle. Inflate a 30 cm (helium quality) balloon to stretch the rubber. Deflate the balloon and attach it to the mouth of the bottle, with another person holding the bottle.. Swirl the bottle contents again and return the bottle to the plastic beaker. Let the balloon swell to 30 cm diameter then remove from the bottle, with another person holding the bottle,  and attach a balloon clip with attached string. If the balloon does not rise up when you hold the attached string either the balloon is not yet at full capacity or the reaction has occurred too quickly and condensed water vapour inside the balloon is weighing it down. Wash the remaining reagents down the laboratory sink.
2Al (s) + 2 NaOH (aq) + 2H2O --> 2NaAlO2 + 3H2 + energy
aluminium + sodium hydroxide + water --> sodium aluminate + hydrogen gas + energy
4. Calcium with hydrochloric acid
Use forceps to transfer about 0.1 g of calcium metal turnings to dilute hydrochloric acid in a test-tube.
Ca (s) + 2HCl (aq) --> CaCl2 (aq) + H2 (g)
5. 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.]
6. 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.
7. Iron filings with alum
Put 5 g of iron filings in 1 cm depth of alum solution [Al2(SO4)3.K2(SO4).24H2O, potash alum, alum], [also shown as KAl(SO4)2.12H2O] in a test-tube. Heat the solution until effervescence occurs.
8. Zinc with hydrochloric acid
See diagram 3.2.33: 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, or through a syringe. 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)
9. Zinc with hydrochloric acid
Prepare hydrogen gas in a small borosilicate test-tube by reaction of zinc with dilute hydrochloric acid and 'pop.' the air / hydrogen mixture with a lighted taper. Check that the test-tubes are not chipped or cracked.
10. 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.
11. Iron filings with citric acid, sulfuric 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 gas is colourless and odourless but any impurities in the iron filings give a nasty smell. To tests for hydrogen gas, remove from heating, place your thumb over the end of the test-tube, count to five, apply a lighted paper to the end of the test-tube, the hydrogen gas explodes with a loud pop sound. Never test more than a test-tube full of hydrogen gas!
3.41.01 Hydrogen generator from a boiling tube
See diagram 3.2.34: Small hydrogen generator
Hydrogen generators are not allowed in many school systems.
Make holes in the bottom of a boiling tube. Heat the bottom of a boiling tube and a glass rod to red heat in a Bunsen burner flame. Fuse the glass rod on to the bottom of the boiling tube then pull it away to form a shred of glass pulled out from the boiling tube. Break off the shred of glass and form smooth rounded edges to the hole in a hot flame. Make three holes. Put granulated zinc into the boiling tube and fix a one-hole stopper with a delivery tube, rubber tube extension and screw clip. If the zinc is in very small pieces put glass wool in the bottom of the boiling tube before adding the zinc. Put the apparatus in a jar containing M sulphuric acid and drops of copper sulfate solution. Open the clip to allow acid to enter the boiling tube and react with the zinc to form hydrogen gas. Close the clip to allow pressure inside the boiling tube to prevent acid reacting with the zinc.
3.41.02 Explode a hydrogen balloon
Explode a hydrogen balloon
Never explode hydrogen in a glass or solid container. Tie the string of the 30 cm hydrogen balloon to a weight on the desk, well away from any chemicals or combustible substances. Attach a small birthday candle to a long stick or thin dowel. All observers must stand at least 5 metres away and must protect their ears. Keep a soda acid fire extinguisher nearby. With your safety goggles and ear protection on, light the candle and hold it under the floating balloon at arm's length.
2H2 + O2 --> 2H2O + energy
3.41.1 Lighted splint test for hydrogen gas
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.
1. 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)
hydrogen gas + oxygen gas --> water
3.41.1.1 Litmus test for hydrogen gas
Hydrogen does not change the colour of moist litmus.
3.41.1.2 Pouring test for hydrogen gas
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 gas 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 gas
See diagram 3.41.3: Hydrogen over heated copper oxide
Be careful! Use a safety screen and wear eye protection
1. Pass hydrogen gas over copper (II) oxide, or lead (II) oxide (lithage) or iron (III) oxide. Hydrogen gas reduces metal oxides to metals. The products are the metal and water.
CuO (s) + H2 (g) --> Cu (s) + H2O (l)
2. 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.
3. 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: 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)
13.01 Gas bags
See diagram 13.01: Gas bag, cable tie
1. Party balloons can be inflated with gas only from a high pressure source, e.g. a gas cylinder.
2. Snap lock resealable polythene bags. can be resealed with a finger press sealing strip to give a gas-tight seal. The closure system which reseals an opened bag includes a pressure sensitive adhesive on the front side and a defined release surface on the back of the bag. The top portion of the bag is folded so that the defined release surface comes into contact with the adhesive to reseal the opened bag.
3. The plastic bag used in a 2 litre wine cask can be washed out and used to store gas. Use a cork borer to insert a glass tube through a one hole rubber stopper. Be Careful! Leave 1 cm of glass tube to protrude from the top of the stopper. Pull tight around the neck of the plastic bag around the rubber stopper and securely it tightly with a cable tie. To check for leaks, close the end of the glass tube with a rubber cap, immerse the bag in water and squeeze the bag. To fill the bag, squeeze it flat then fill it from a gas cylinder or chemical generator. Refill the bag and squeeze out the gas more than once to ensure that any air is flushed out. When the bag is finally filled, close the glass tube with a rubber cap. To get a gas sample, inject the hypodermic needle of a syringe through the rubber cap and suck gas into the syringe.
A cable tie usually consists of a Nylon tape with a gear rack and a ratchet within a small open case. When the pointed tip of the cable tie has been pulled through the case and past the ratchet, it cannot be pulled back, but the loop formed may be pulled tighter. Cable ties are used to bind several cables together, e.g. around a motor car engine.
13.1.5a Relative molecular mass of gases
See diagram 13.1.5: Relative molecular mass of gases | See SVP: Saturated vapour pressure over water
The relative molecular mass, M, of a compound is the ratio of the average mass of molecules of the substance to 1 / 12 of the mass of one atom of C-12. Number of moles = volume in litres / 22.4 litres / mol. At STP, 1 mol of most gases occupies 22.4 L at STP. At 25oC, 1 mol of most gases occupies 24.45 L
Weigh a gas container. Collect 1 litre of gas in an inverted measuring cylinder over water. The levels of water inside and outside the measuring cylinder must be the same. Weigh the gas container again. Calculate the loss in weight (about 2 g). Note the temperature and atmospheric pressure. For propane, if loss in weight of gas container = 1.8 g, 1.8 X 24.45 = 44 = relative molecular mass of propane. Use other sources of gas, e.g. a cigarette lighter. Hold it under water below the measuring cylinder with the valve kept open with a rubber band.
13.1.6a Molar volume of oxygen prepared with hydrogen peroxide
See diagram: 13.1.6: Molar volume of oxygen | See SVP: Table of saturated vapour pressure over water, Psvp
Put 15 mL of 3% w/w (3 g H2O2 /100 g solution) hydrogen peroxide solution into flask A. Put 0.05 g of yeast in a small test-tube then lower the test-tube into flask A. Weigh flask A and its contents, W1. Attach Plastic tube 1 and Plastic tube 2 to flask B only. Put water into flask B leaving a space in the neck of the flask. Add water to a beaker until it is one third full. Siphon water into the beaker by blowing into the open end of Rubber tube 1 or by using a pipette bulb. Raise and lower the beaker to remove any air bubbles from Plastic tube 2. Adjust the height of the beaker so that the levels of the water in the beaker and in flask B are the same. Connect Plastic tube one to flask A. Raise the beaker to check for leaks in the apparatus. Again, adjust the height of the beaker so that the levels of the water in the beaker and in flask B are the same. Close the pinch clamp. Replace the glass tube in the beaker and open the pinch clamp to allow some water to flow into the beaker. With the pinch clamp still open, tip flask A so that the yeast falls into the hydrogen peroxide solution. Swill flask A until the reaction is completed when the water level in the beaker does not change. Again, adjust the height of the beaker so that the levels of the water in the beaker and in flask B are the same. Close the pinch clamp. Remove the stopper in flask A, insert a thermometer and note the temperature of the gas inside, T1. Repeat this measurement with flask B, T2. Disconnect Plastic tube 1 from flask A and again weigh flask A and its contents, W2. Measure the oxygen produced, by measuring the volume or weight water in the beaker, V. Find the vapour pressure of water at that temperature from the Table of saturated vapour pressure over water, Psvp. Note the room temperature. Note the barometric pressure from a barometer or ask the weather bureau or local airport.
Calculate the volume at STP. of 32 g, one mole, of oxygen gas.
(W2 - W)1 = Loss in weight
VO2 = volume of water in the beaker = volume of oxygen collected
Tf = average temperature in the flasks = (T1 + T2) / 2
Patm = atmospheric pressure = pressure of oxygen in the flask, PO2 + saturation vapour pressure of water at that temperature, Psvp.
so PO2 = (Patm - Psvp)
VO2 = volume of oxygen in the flask
Tstp = temperature at STP. (Standard Temperature and Pressure) = 0oC, 273 K
Pstp = pressure at STP. = 760 mm Hg = 101325 Pa
Vstp = volume at STP.
P1V1 / T1 = P2V2 / T2 (Boyle's law and Charles's law)
(P1 X V1) / T1 = (P2 X V2) / T2
(PO2 X VO2) / Tf = (Pstp X Vstp) / Tstp
So Vstp = VO2 [(PO2 / Pstp) X (Tstp / Tf)]
Relative molecular mass of oxygen = 32 g
So number of moles of oxygen = (W2-W1) / 32
So the molar volume of oxygen at stp = Vstp / number of moles of oxygen = litres / mole
A mole of an ideal gas occupies 22.4 litres at STP.
If barometric pressure = 1016 kPa, average temperature = 20oC, loss in weight of flask = 0.2 g, volume of oxygen collected at average temperature = 140 mL
Pressure of oxygen in apparatus = (barometric pressure - SVP at 20oC) = (1016 - 2.3) = 1013.7 kPa
Vstp = 140 [(1013.66 / 101.325) X (293 / 273)] = 503.17 mL
Number of moles = 0.2 / 32 = 0.00625 moles
Molar volume = 140 / 0.00625 = 22400 = 22.4 litres = 22.4 L / mole
13.3.2 Burn sulfur in oxygen
Dip a wire loop into sulfur powder. Ignite the sulfur in a burner flame and then put it into a test-tube of oxygen. The sulfur burns with a bright blue flame.
13.3.3 Burn steel wool in oxygen, burn iron filings
Collect oxygen in test-tubes with stoppers. Store test-tubes in a test-tube rack and remove the stoppers just before inserting the burning element. Fasten steel wool to wire. Heat the steel wool in a burner flame. Put it into a test-tube of oxygen. The steel wool burns with bright sparkles to form black-grey iron oxide, Fe3O4(FeO.Fe2O3) magnetite, lodestone, iron ore.
Repeat the experiment by sprinkling iron filings into a Bunsen burner flame. A shower of sparks occurs as in some fireworks.
13.3.4 Burn magnesium ribbon in oxygen
Wrap a 3 cm piece of magnesium ribbon around the loop at the end of a wire. Ignite it in a burner and put it quickly in the oxygen. Magnesium burns with a very bright flame.
BE CAREFUL! Do not look directly at the flame because its brightness can cause injury to eyes. The white smoke is magnesium oxide. Its toxicity is low, but inhalation should be avoided.
Put the ash on a watch glass and add 3 mL of deionized water to wet the ash thoroughly and leave it lying in a small pool of water. Add one small drop of phenolphthalein solution and leave to stand for two minutes. Magnesium oxide has a low solubility in water, so you will not see any visible evidence that any of the solid has dissolved. Add one drop of dilute hydrochloric acid solution and leave to stand until the solution around the solid ash will turn pink, showing that the solution has become alkaline. This is evidence that some magnesium oxide has dissolved. Oxide ions in the solid react with water to form aqueous hydroxide ions. When no further change occurs, add a second drop of dilute hydrochloric acid. The pink colour disappears almost instantly, showing that the hydroxide ions have been neutralized very quickly, and replaced by an excess of hydrogen ions. During the next 2 to 15 minutes, depending on the size and concentration of the drop of acid added, the mixture changes slowly back to pink as the excess acid being neutralized slowly by solid magnesium oxide, followed by slow dissolving of remaining magnesium oxide to make the solution. When no more changes occur, add a second small drop of dilute hydrochloric acid. The same cycle of discharge and reappearance of pink colour can be repeated for as long as any solid magnesium oxide remains. Magnesium oxide is a metallic oxide, and is therefore basic. Magnesium oxide has a low solubility in water. Dilute hydrochloric acid reacts rapidly with aqueous magnesium hydroxide, but slowly with solid magnesium oxide. Magnesium oxide dissolves slowly in water. Phenolphthalein is an indicator that shows changes in alkalinity of the solution. An equilibrium is established between solid magnesium oxide and dissolved magnesium ions. The addition of acid disrupts the equilibrium by removing hydroxide ions from the solution. An equilibrium is established between solid magnesium oxide and dissolved magnesium ions. The addition of acid disrupts the equilibrium by removing hydroxide ions from the solution. Equilibrium is restored by slow dissolving of more magnesium oxide. Addition of larger drops or higher concentration of acid causes a larger initial excess of acid in the solution. Because the reaction of acid with the solid magnesium oxide is slow, it will take a much longer time for the pink colour to return to the mixture. The magnesium oxide formed from combustion of magnesium ribbon forms a hard mass with a small surface area for reaction. The rate of reaction with acid, and the rate of solution of the solid to form an alkaline solution, would be increased by crumbling the ash.
13.3.5 Prepare oxygen absorbent
Dissolve 300 g of ammonium chloride in 1 litre of water and add 1 litre of concentrated ammonia solution. Shake the solution. Pass the gas through the solution after adding half the volume of copper turnings.
13.4.0 Chlorine, Cl2, Highly toxic, granular, liquid, powder, tablets, Highly irritant to lungs, chloride Cl-, chloro -Cl,
Solution < 3%, Highly toxic by all routes
Chlorine gas, < 3%, Not Hazardous if small volume in cross-ventilation
1. Chlorine gas is a powerful lung irritant, causing coughing and diminishing lung efficiency. It was used as a poison gas in the First World War. Carry out all reactions that may result in evolution of chlorine in a fume cupboard. Chlorine gas is very toxic. Can react to cause fires or explosions upon contact with turpentine, ether, ammonia gas, illuminating gas, hydrocarbon, hydrogen gas and powdered metals. It dissolves readily in water forming highly corrosive solution. Do not prepare chlorine in an open room but use a fume cupboard. Direct combination of chlorine and hydrogen gas occurs in bright light or ignition of the mixture by lighted taper or electric spark. It reacts with metals, solid non-metals, hydrocarbon. Use small quantities only. Chlorine is a green-yellow, dense gas that causes rapid corrosion of metals and destruction of plastics. It is also a dangerous gas because it attacks the mucous membrane linings of the eyes, nose, throat and lungs, causes the lungs to fill with fluid and the victim drowns. During the First World War it was used as a chemical weapon. Chlorine is prepared commercially from electrolysis of concentrated sodium chlorine (brine) solution. Chlorine is a very reactive non-metal and free chlorine never occurs naturally. Do all chlorine experiments in a fume cupboard. Chlorine kills most living things and is used to sterilize drinking water and disinfect swimming pools. Chlorine is used to manufacture the plastic PVC, to bleach wood pulp and to prepare organic compounds such as solvent tetrachloroethene CCl2.CCl2, solvent tetrachloromethane (carbon tetrachloride) CCl4, safer solvent 1,1,1-trichloroethane CH3CCH3 and the insecticide DDT (C6H4Cl)2CH-CCl3 [Former name: dichlorodiphenyltrichloroethane, New IUPAC name: 1,1,1-trichloro-2,2-bis (4-chlorophenyl)ethane] However, many of these substances cannot be broken down in the environment (biodegraded) so you should avoid using them. Chlorine is a powerful oxidizing agent.
13.4.1 Prepare chlorine
See diagram 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 and can be poured from one container to another. 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. Solutions of bleach, (sodium hypochlorite). and solid bleaching powder produce small amounts of chlorine gas when exposed to the air, with its characteristic smell. Adding acid causes a vigorous production of chlorine gas. The most convenient ways to prepare chlorine gas are to use the reaction of dilute acid with bleaching powder or the reaction of potassium permanganate with concentrated hydrochloric acid. Collect the chlorine gas by displacement of air.
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.
2. 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.
CaOCl2 + H2SO4 (aq) --> CaSO4 (s) + H2O (l) + Cl2 (g)
3. Repeat the experiment with dilute hydrochloric acid
CaOCl2 + 2HCl (aq) --> CaCl2 (s) + H2O (l) + Cl2 (g)
4. Add 5 M hydrochloric acid to sodium hypochlorite. Keep a 5M alkali solution nearby to stop the reaction.
5. Add a dilute acid to bleach solution to form chlorine gas.
NaOCl (aq) + HCl (aq) --> NaCl (aq) + H2O (l) + Cl2 (g)
6. Prepare chlorine with concentrated hydrochloric acid and manganese (IV) oxide
Put some manganese (IV) oxide in a boiling tube and add drops of concentrated hydrochloric acid from the reservoir. Heat the test-tube gently. Observe the slight green colour in the tube. The wider the tube, the easier this is to see.
Use potassium manganate (VII) instead of manganese (IV) oxide to prepare chlorine because the reaction does not require heating avoiding hot concentrated hydrochloric acid.
Be careful! Prepare chlorine only in a fume cupboard
4HCl (aq) + MnO2 (s) --> MnCl2 (aq) + 2H2O (l) + Cl2 (g)
7. Prepare chlorine, concentrated hydrochloric acid with potassium manganate (VIII)
This experiment may not be allowed in some school systems. Use a fume cupboard.
See diagram 12.19.8.2: Prepare chlorine
8. Put 5 cc of potassium permanganate into a flask. Fill the dropping funnel with concentrated hydrochloric acid and allow the acid to run on to the permanganate to produce chlorine.
2MnO4- + 16H + + 10Cl- --> 2Mn2+ + 8H2O + 5Cl2 (g)
Pass the gas through water to remove hydrogen chloride and pass the gas through concentrated sulfuric acid to dry it. Collect the gas by downward displacement of air
9. Connect a conical flask by means of a delivery tube to a collection vessel. Put about 5 g of solid potassium manganate (VII) (potassium permanganate) in a conical flask and add concentrated hydrochloric acid drop by drop.
16HCl (aq) + 2KMnO4 (s) --> 2KCl (aq) + 2MnCl2 (aq) + 8H2O (l) + 5Cl2 (g).
13.4.2 Prepare chlorine water
Chlorine water, chlorine solution, Toxic by all routes, forms lung irritant chlorine gas, (Use small quantities, < 10 mL or 10 g in well-ventilated area)
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)
13.4.3 Tests for chlorine
1. Lighted splint test for chlorine
Chlorine extinguishes a lighted splint, but hot steel wool burns in it
2. Bleaching test for chlorine
Chlorine bleaches moist red or blue litmus paper, flowers and some dyes in cloth. Use chlorine gas to bleach flower petals, leaves and hair suspended in the gas.
13.4.4 Bleaching powder
Domestic bleach is manufactured by mixing chlorine solution with sodium hydroxide solution. A bleach is a chemical used to make white or remove the colouring from fibres, textiles and hair ("peroxide blondes"), e.g. hydrogen peroxide, sodium hypochlorite. Commercial bleaching powder is probably a mixture of calcium chlorate (I), calcium chloride and calcium hydroxide. The commercial value of bleaching powder depends on its available chlorine. Bleaching may occur by oxidation using free chlorine or combined chlorine or by exposure to sunlight, e.g. "fading" of the curtains. Bleaching is important in the textile and paper industries. Domestic bleach is manufactured by mixing a solution of chlorine with sodium hydroxide solution. Hypochlorous acid, HClO, "bleach", is used as a bleach and a disinfectant.
Cl2 (g) + 2OH- (aq) --> Cl- (aq) + ClO- (aq) + H2O
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.
13.4.6 Reaction of chlorine with benzene, dichlorbenzene, C4H4Cl2
Chlorine reacts with benzene, C6H6, to form three liquid aromatic isomers by chlorinating benzene with an iron filings catalyst, then separate isomers by fractional distillation. The 3 separate compounds, isomers. called dichlorobenzene are  1,2-dichlorobenzene (ortho-), 1,3-dichlorobenzene (meta-), and 1.4-dichlorobenzene (para-).
1,2-dichlorobenzene, (ortho-dichlorobenzene), b.p. 179oC, used in insecticides and to make dyes, C6H4Cl2, chloroben (used for termite treatment), o-dichlorobenzene, Toxic, avoid inhalation of vapour
Solution < 5% Not hazardous
1.4-dichlorobenzene, (para-dichlorobenzene), b.p. 174oC, colourless solid, strong odour, used as a deodorant and an insecticide substitute for naphthalene in the common household "mothballs" or "moth crystals". C6H4Cl2, p-dichlorobenzene, paradichlorobenzene, PDB,  p-DCB, 1,4-dichlorobenzol, paramot, "dichlorobenzene", Toxic if ingested, may be carcinogenic
Solution < 25%, Not hazardous
May react with sulfide groups to form the polymer poly(p-phenylene sulfide) and sodium chloride
C6H4Cl2 + Na2S --> 1/n (C6H4S)n + 2Nacl
The two isomers above can react to form (meta-dichlorobenzene) 1,3-dichlorobenzene
C6H6 (l) + Cl2 (g) --> C2H4Cl2 (l)
benzene + chlorine --> dichlorbenzene
13.4.6.1 Halogenation of benzene
Halogenation is the substitution of -H by -Cl or -Br
Benzene reacts with bromine or chlorine in the presence of aluminium bromide / aluminium chloride or iron, as electrophilic substitution reactions.
Iron takes part in the reaction so it is actually the bromide or chloride formed that acts as the catalyst, in the same way as does aluminium chloride.
2Fe + 3Br2 --> 2FeBr2
2Fe + 3Cl2 --> 2FeCl3
Chlorinating benzene with aluminium chloride catalyst, AlCl3. Bubble chlorine into a mixture of benzene and anhydrous aluminium chloride.
C6H6 (l) + Cl2 (g) --> C6H5Cl (l) + HCl (g)
With iron catalyst
C6H6 + Br2 --> C6H5Br + HBr
13.4.7 Reactions of chorine with sodium
See diagram 13.4.7: Reactions of chlorine with sodium
BE CAREFUL! THE REACTION IS VERY VIGOROUS! Do this experiment in a fume cupboard.
Dry a small piece of sodium with absorbent paper. Grip a piece of sodium with a pair of tongs. File the sodium and let the obtained sodium filings fall into chlorine gas collected in a test-tube. The sodium filings react violently with the chlorine, sparks flying off, to form many smoke particles of sodium chloride and a crust of sodium chloride on what is left of the piece of sodium. When the reaction has stopped, wash the residue in methylated spirit to remove unreacted chlorine. Let chlorine leave the test-tube by diffusion. Crystals of sodium chloride remain in the test-tube.
2Na (s) + Cl2 (g) --> NaCl (s)
2. Put a pin head volume of sodium in the bowl of a deflagrating spoon. In a fume cupboard, put the spoon into a test-tube of chlorine and leave to stand. When the reaction stops, remove the spoon, allow it to cool and place it in a small amount of alcohol. Let excess chlorine diffuse away in the fume cupboard. Let the mixture of alcohol and solid stand until no further reaction takes place. Wash the crystals with alcohol and let them cool and dry.
sodium (s) + chlorine (g) --> sodium chloride (s)
13.4.8 Burn substances in chlorine
1. Burn steel wool in chlorine
Ignite steel wool held by tongs in a burner flame, then put into chlorine gas.
BE CAREFUL! The reaction occurs with strong combustion to form a brown-red cloud that condenses to black flakes of anhydrous iron (III) chloride.
2Fe (s) + 3Cl2 (g) --> 2FeCl3 (s)
2. Burn copper in chlorine
Heat copper foil in a burner flame and put into chlorine gas. The reaction forms a layer of brown copper (II) chloride that turns green in the presence of moisture.
Cu (s) + Cl2 (g) --> CuCl2 (s)
3. Burn magnesium in chlorine
A burning magnesium ribbon burns violently in chlorine gas to form magnesium chloride. .
4. Burn a wax taper in chlorine, reaction of chlorine with non-metals
Chlorine has such a strong attraction for hydrogen that it removes all the hydrogen from the hydrocarbon paraffin leaving behind the carbon as a residue. A mixture of chlorine and hydrogen gas does not react in the dark but if heated or exposed to strong sunlight the mixture reacts explosively to form hydrogen chloride.
BE CAREFUL! Do not mix chlorine and hydrogen gas!
H2 (g) + Cl2 (g) --> 2HCl (g) + energy
5. Burn a paraffin wax taper or a small birthday candle in chlorine
The taper keeps burning with a dull red flame and forms black carbon particles (soot) and hydrogen chloride gas.
nCl2 + (CH3-CH2-CH2-CH2-) --> nHCl + nC
6. Burn a taper in chlorine
A burning wood taper continues to burn in chlorine gas, forming hydrogen chloride and organic chloro-compounds.
13.4.9 Pass chlorine through water
Chlorine is available commercially for school laboratory use as chlorine water. Hypochlorous acid HClO, a bleach and a disinfectant, is an aqueous 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)
13.4.12 Pass chlorine through iodine solution
The more reactive chlorine displaces iodine from its salt. The colourless potassium iodide solution turns red then black as iodine is displaced from the solution. Tests for iodine with starch solution.
Cl2 (g) + 2KI (aq) --> I2 (aq) + 2KCl (aq)
13.4.13 Pass chlorine through iron (II) chloride solution
Pass chlorine through iron (II) chloride solution. The chlorine oxidizes iron (II) chloride to iron (III) chloride. The solution changes from green to brown.
2Fe2+ (aq) + Cl2 (g) --> 2Fe3+ (aq) + 2Cl- (aq)
FeCl2 (aq) + Cl2 (g) --> 2FeCl3 (aq)
13.4.14 Reactions of chlorine with heated copper and steel wool
BE CAREFUL! Do not breath this poisonous gas.
1. In a fume cupboard, put a heated spiral of copper wire into a small test-tube of chlorine. The heated copper is immediately covered with brown copper (II) chloride that turns green in the presence of water.
Cu (s) + Cl2 (g) --> CuCl2 (s)
2. Heat a lump of steel wool and plunge it into the chlorine gas. Brown fumes form that condense to black flakes of anhydrous iron (III) chloride.
2Fe (s) + 3Cl2 (g) --> 2FeCl3 (s)
13.4.15 Reactions of chlorine with alkalis, bleaching powder
1. Pass chlorine slowly through test-tubes containing dilute sodium hydroxide solution, dilute potassium hydroxide solution and solid calcium hydroxide. Add dilute sulfuric acid to the products of any reactions. Chlorine reacts with cold alkali solutions to form chloride ions, Cl- and hypochlorite ions, ClO-, powerful bleaching agents.
Cl2 (g) + 2OH- (aq) --> Cl- (aq) + ClO- (aq) + H2O (l)
2. Pass excess chlorine into hot alkali solutions to form chloride and chlorate ions. If passed into potassium hydroxide, the less soluble potassium chlorate can be separated from the less soluble potassium chloride by fractional distillation.
3Cl2 (g) + 6OH- (aq) --> Cl- (aq) + ClO3- (aq) + 3H2O (l)
3. Chlorine reacts with strongly basic hydroxides, e.g. calcium hydroxide, and strongly basic oxides, e.g. calcium oxide, in the solid state. The products have a variable composition. Reactions of chlorine with calcium hydroxide produces bleaching powder, a convenient source of chlorine and a powerful bleaching agent in dilute acid solutions. Bleaching powder reacts with sulfuric acid to give off chlorine.
Bleaching powder (s) + sulfuric acid (aq) --> calcium sulfate (s) + chlorine (g) + water (l)
13.9.0 Nitrogen
Nitrogen gas N2 is colourless, odourless, tasteless, neutral and unreactive. Nitrogen does not support combustion. Magnesium and calcium will continue to burn in nitrogen to form nitrides. Nitrogen is manufactured by fractional distillation of air. Air contains about 78% of nitrogen.
13.9.3 Nitrogen gas generated in a motor car airbag
A gas generator containing a mixture of sodium azide, NaN3, potassium nitrate, KNO3 and silica, SiO2 is ignited electrically to allow a slow detonation so that nitrogen fills the airbag.
2NaN3 --> 2Na + 3N2 (300oC)
10Na + 2KNO3 --> K2O + 5Na2O + N2
K2O + Na2O + SiO2 --> alkaline silicate
13.10.2 Prepare nitrogen dioxide, NO2, heat lead (II) nitrate crystals
Heat lead (II) nitrate crystals. The decomposition may be noisy. Nitrogen dioxide and oxygen form, leaving yellow lead oxide.
Pb(NO3)2 (s) --> 4NO2 (g) + 2PbO (s) + O2 (g)
lead (II) nitrate --> nitrogen dioxide + lead oxide + oxygen
13.12.0 Sulfur dioxide with water
Sulfur dioxide, SO2, is a colourless gas that irritates the lungs. Sulfur dioxide dissolves in water to form, mainly, sulfurous acid (H2SO3). Sulfur dioxide is one component of acid rain.
SO2 (g) + H2O (l) --> H2SO3 (l)
13.13.3 Prepare sulfur dioxide with sulfuric acid and sodium sulfite
See diagram 13.13.3: Prepare sulfur dioxide
Na2SO3 (s) + H2SO4 (l) --> Na2SO4 (aq) + H2O (l) + SO2 (g)
13.13.4 Prepare sulfur dioxide with sulfuric acid and copper
Add hot concentrated sulfuric acid to copper to form copper (II) sulfate, water and sulfur dioxide. BE CAREFUL!
Cu (s) + 2H2SO4 (l) --> CuSO4 (aq) + 2H2O (l) + SO2 (g)
13.13.8 Dry hydrogen sulfide and dry sulfur dioxide will not react
Collect sulfur dioxide in a dry test-tube after passing the gas slowly through concentrated sulfuric acid to dry it. Collect a test-tube of hydrogen sulfide, after passing it over calcium chloride tube to dry it. Invert the test-tube containing sulfur dioxide over the test-tube containing the hydrogen sulfide. No reaction occurs. Leave to stand then pour drops of water into the lower test-tube and quickly replace the upper test-tube. sulfur immediately precipitates in the test-tubes.
2H2S + SO2 --> 2H2O + 3S (s)