School Science Lessons
Topic 13
2019-08-18
Please send comments to: J.Elfick@uq.edu.au

Topic 13 Gases, prepare gases, greenhouse gases
Table of contents

13.0.0 Gases

13.2.0 Air, atmosphere

Gas, gases

13.1.0 Gases by name, Argon to Xenon

13.1.0 Gases by name, Ammonia to Xenon
See: Gases, (Commercial)
3.48 Acid rain and nitrogen oxides, NOx
3.33 Ammonia, NH3
7.2.2.3 Argon, Ar
Bromine, Br
Butane C4H10
3.34 Carbon dioxide, CO2
13.1.5 Carbon disulfide, CS2
3.39.0 Carbon monoxide, CO
Chlorine, Cl2
13.1.9 Ethane, C2H6
13.1.10 Ethene (ethylene), C2H4
13.1.11 Ethyne (acetylene), C2H2
Fluorine, F
Helium, He
13.1.14 Heptane, C7H16
13.1.15 Hexane, C6H14
3.41.0 Hydrogen gas, H2
12.1.1 Hydrogen gas, aluminium with acids
3.41.2.0 Prepare hydrogen gas
12.1.1 Prepare hydrogen gas, aluminium with acids
3.41.2.1 Prepare hydrogen gas bubbles
12.6.3 Prepare hydrogen gas with citric acid
3.41.2.0 Prepare hydrogen gas, zinc with hydrochloric acid
13.1.17 Hydrogen bromide, HBr
3.42.0 Hydrogen chloride, HCl
Hydrogen fluoride, HF, hydrofluoric acid
3.43.0 Hydrogen sulfide, H2S
13.1.17 Krypton, Kr, Table of the elements
Methane, CH4, natural gas
Neon, Ne
13.1.25 Nitric oxide, nitrogen monoxide, NO
13.1.24 Nitrogen gas, N2
13.1.23 Nitrogen dioxide, NO2
3.48 Nitrogen oxides, Acid rain and NOx
12.19.6.13 Nitrogen triiodide, NI3
13.1.27 Nitrous oxide, N2O
16.6.8.0 Octane, C8H18
13.1.29 Oxygen gas, O2
12.17.0 Oxides, Reactions of oxides
Ozone, O3
19.4.22 Packaging gases, propellant, food additive
13.1.32 Pentane, C5H12
3.39.2 Phosgene, carbonyl chloride, CoCl2
13.1.33 Propane, C3H8
Radon, Rn
13.1.35 Sulfur dioxide, SO2
13.1.36 Sulfur trioxide, SO3
Trichloroethane (1, 1, 1-trichloroethane), methyl chloroform, CH3CCl3
Trichloromethane, CHCl3, chloroform
7.2.2.47 Xenon, Xe.

13.1.5 Carbon disulfide, CS2
Carbon disulfide, CS2, carbon bisulfide, Toxic
7.9.22 Flammable (See: 2.)
15.7.0 Flammable organic chemicals

3.41.0 Hydrogen gas, H2
See: Hydrogen Elements, Compounds, (Commercial)
Hydrogen, Table of Elements
7.2.2.19 Hydrogen, properties
3.8.5 Hydrogen gas hazards

Hydrogen experiments
12.2.5.0 Acid-base reactions (hydrogen ions, H+)
Coal gas (50% hydrogen gas)
Density of gases, Hydrogen (Table)
3.6.1 Deuterium (hydrogen isotope)
8.0.0 Direct union of elements to form compounds
12.12.10 Drain cleaners, e.g. "Drano"
7.9.22 Flammable (for hydrogen, See4.)
7.9.28 Fuel cell (hydrogen gas)
12.4.0 Hydrochloric acid (forms hydrogen gas)
3.01.3 Hydrogen bonds
3.2.1 Hydrogen bonds, Liquid viscosity
19.4.22 Packaging gases, propellant, food additive
12.3.0.4 pH (hydrogen ion concentration)

Hydrogen experiments
12.7.3 Alkalis with metals, sodium hydroxide (See 2.)
12.1.1 Aluminium with acids
12.1.2 Aluminium with sodium hydroxide
12.1.8 Aluminium chloride with water
3.41.3 Burn hydrogen gas
3.94 Catalysts and rate of reaction
17.1.1 Count bubbles, dilute hydrochloric acid with 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 sulfuric acid with aluminium
12.3.3 Dilute sulfuric acid with steel wool
3.86.1 Electrode potential order of metals
15.5.12 Electrolysis of sodium chloride solution
15.5.4 Electrolysis of water, decomposition of water, Hofmann voltameter, coulometer
3.41.5 Explode a hydrogen gas balloon
33.84.1 Galvanic cell, Voltaic cell
14.1.10 Heat of reaction, potassium with diethyl ether
4.47 "Human" water (hydrogen boy, oxygen girl )
3.41.4 Hydrogen gas generator from a boiling tube
12.3.3.3 Iron with sodium hydrogen sulfate
17.2.3 Iron with sulfuric acid (See 2.)
33.3.4 Lemon cells, lemon battery, electricity from lemons
3.74 Metals displace hydrogen from acids
12.2.4.3 Magnesium displaces ethanoic acid
12.3.3.2 Magnesium with sodium hydrogen sulfate
12.0.0 Prepare hydrogen gas
16.1.3.1.2 Prepare sodium ethoxide
12.15.1 Prepare silica and silicon, (See 4.)
12.18.5.4 Reactions of dilute sulfuric acid
12.19.9.1 Reactions of bromine, (See 1.)
12.19.9.4 Reactions of hydrogen bromide (See 7.)
12.15.3 Reactions of metals with steam
12.15.1 Reactions of metals with water
3.73 Reactions of sodium with water
16.5.1.4 Reduce copper oxide with methane
5.1.14 Relative atomic mass of magnesium
33.3.1 Simple electric cell
3.41.1.0 Tests for hydrogen gas
12.11.3.5 Tests for substances with dilute hydrochloric acid

12.0.0 Prepare hydrogen gas
12.1.1 Prepare hydrogen gas,aluminiumwith acids
3.41.2.0 Prepare hydrogen gas
3.41.2.1 Prepare hydrogen gas bubbles
3.41.2.0 Prepare hydrogen gas, zinc with hydrochloric acid
12.6.3 Prepare hydrogen with citric acid

3.41.1.0 Tests for hydrogen gas
3.41.1 Lighted splint tests for hydrogen gas
3.41.1.1 Litmus tests for hydrogen gas
3.41.1.2 Pouring test for hydrogen gas

3.42.0 Hydrogen chloride, HCl
Be careful!
Prepare hydrogen chloride gas in fume hood!

3.42.0 Prepare hydrogen chloride / hydrochloric acid
18.6.2 Air pollution from burning refuse
3.30.15 Decomposition of chlorides
14.0 Electrophiles and nucleophiles, HCl
12.19.5.0 CFCs, chlorofluorocarbons, "Freons"
Density of gases, Hydrogen chloride, (Table)
3.8.6 Hydrogen chloride, anhydrous, hazards

Experiments
31.1.14 Aluminium foil precipitator
3.33.6 Ammonium chloride smoke screen
10.1.2 Diffusion rates of ammonia and HCl
17.5.5.5 Effect of temperature on equilibrium
12.8.13 Heat hydrated iron chlorides
3.42.1.7 Hydrogen chloride fountain test
12.13.4 Phosphorus trichloride with water
3.42.0 Prepare hydrogen chloride / hydrochloric acid
12.19.8.3 Prepare iron (III) chloride, (See 2.)
12.20.2 Prepare tin (IV) chloride
12.19.8.1 Reactions of sodium chloride
12.18.5.2 Sulfuric acid displaces acids from salts
3.42.1.0 Tests for hydrogen chloride
12.11.3.6 Tests for gases with hot conc. sulfuric acid
11.3.6 Tests for gases with hot conc. sulfuric acid
12.14.2.5 Zinc with copper in sulfuric acid

3.42.1 Tests for hydrogen chloride
3.42.1.6 Ammonia test for hydrogen chloride
3.42.1.7 Fountain test for hydrogen chloride
3.42.1.4 Lighted splint test for hydrogen chloride
3.42.1.2 Litmus paper test for hydrogen chloride
3.42.1.5 Magnesium ribbon test for hydrogen chloride
3.42.1.1 Solubility test for hydrogen chloride

3.43.0 Hydrogen sulfide
Hydrogen sulfide, H2S, rotten egg gas
Be careful!
Prepare hydrogen sulfide in fume cupboard!

Coal gas
Density of gases, Hydrogen sulfide, (Table)
19.7.3 Hair products
3.8.7 Hydrogen sulfide, hazards
3.4.12.1 Hydrogen sulfide waste bottle
18.2.2.2 Iron in drinking water
4.1.0 Oxygenic phototropic bacteria
12.3.15 Acids with salts (See 5. Hydrogen sulfide)
12.11.7.4 Group IV Insoluble sulfides precipitated
17.5.7.0 Explanation of group analysis
12.12.4 Hydrogen peroxide as oxidizing agent
15.2.14 Hydrogen sulfide as reducing agent
Rotten egg gas, hydrogen sulfide, H2S
20.0.6 Standard temperature and pressure, STP
3.7.16 Sulfides, hazards
16.2.8.2 Sulfides: RSR (where R is not H)

Experiments
12.2.2.1 Heat iron with sulfur, synthesis reaction
12.19.6.1 Prepare hydrogen iodide, HI (See 5.)
3.43.0 Prepare hydrogen sulfide
12.18.2 Prepare sulfides (See 3.)
5.32 Protect mangroves (See 5. ) (Primary)
3.43.2 Reduce potassium manganate with hydrogen sulfide
3.43.3 Reduce iron (III) chloride with hydrogen sulfide
18.6.5 Smell of water, hydrogen sulfide
3.43.1 Tests for hydrogen sulfide solution
11.3.3 Tests for solubility, prepare group analysis
12.11.3.5 Tests for substances with dilute hydrochloric acid
12.11.3.6 Tests for substances with conc. sulfuric acid
9.9.18.6 Use of freshwater algae for hydroponics

13.1.25 Nitric oxide, nitrogen monoxide, NO
Density of gases, nitric oxide, NO, (Table)
Experiments
3.44.1 Catalytic conversion of nitric oxide
12.1.3 Copper cycle reactions (See1.)
12.3.11.0 Dilute nitric acid with copper
13.13.8 Dry hydrogen sulfide and dry sulfur dioxide
4.206 Float eggs in water
3.44 Prepare nitric oxide (nitrogen monoxide, NO)
12.11.1 Reactions of nitrites

13.1.24 Nitrogen gas, N2
3.30.17 Decomposition of manganates
Density of gases, Nitrogen (Table)
1.0.0 Mineral deficiency symptoms (Agriculture)
6.39 Nitrogen cycle
Nitrogen gas, N2, Table of the elements
13.9.3 Nitrogen gas generated in a motor car air bag
1.10.0 Nitrogen deficiency (Agriculture)
7.2.2.31
Nitrogen properties
13.9.1 Nitrogen reacts with metals
19.4.22 Packaging gases, propellant, food additives
Experiments
3.46 Prepare nitrogen gas
12.11.1 Reactions of nitrites (See 8. Nitrogen)

13.1.23 Nitrogen dioxide, NO2
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
3.48 Acid rain and nitrogen oxides, NOx
Density of gases, Nitrogen dioxide, (Table)
Experiments
12.1.3 Copper cycle reactions (See 1.)
17.5.6.2 Heat nitrogen tetroxide, (dinitrogen tetroxide, N2O4)
3.47.1 Pass nitrogen dioxide through water
3.47 Prepare nitrogen dioxide (nitrogen (IV) oxide, NO2)
12.11.2 Reactions of nitrates, NO3-
12.11.1 Reactions of nitrites

13.1.27 Nitrous oxide, N2O
Nitrous oxide, dinitrogen monoxide, N2O, dinitrogen oxide, laughing gas, non-toxic, causes hysteria, use fume cupboard
37.42.1 Composition of the atmosphere and greenhouse gases (See 4. Nitrous oxide)
Density of gases, Nitrous oxide (Table)
19.4.22 Packaging gases, propellants, food additives, E942
Experiments
3.45 Prepare nitrous oxide (dinitrogen oxide) N2O
3.45.1 Tests for dinitrogen oxide (nitrous oxide) N2O

13.1.35 Sulfur dioxide, SO2
Sulfur dioxide, SO2, sulfur (IV) oxide, Toxic, Corrosive, Highly irritating, fumes may affect asthmatics.
Always use a fume cupboard for sulfur dioxide.
Sulfur dioxide, SO2, Toxic, Corrosive, severe respiratory irritant, Additive E220, Sulfur dioxide from coal tar or combustion of S).
Sulfur dioxide preservative in soft drinks, dried fruit, wine.
Acid rain, SOx, from burning sulfur or sulfur compounds: 12.6.0.1
Acidity and alkalinity: 18.1.0.1
Density of gases, Sulfur dioxide, (Table)
Sulfur dioxide and sulfites, Food preservation
Sulfur dioxide to sulfuric acid 1.: 12.6.0.3.1
Sulfur dioxide to sulfuric acid 2.: 12.6.0.3.2
Sulfur dioxide with water: 13.12.0

Experiments
Bleach flowers with sulfur dioxide: 3.51.4
Dilute acids with acidic oxides: 12.3.8
Dry hydrogen sulfide and dry sulfur dioxide will not react: 13.13.8
Halide salts with hot concentrated sulfuric acid: 12.19.3.1
Heat sulfur to form sulfur dioxide: 8.2.15
Oxygen with sulfur dioxide: 12.2.6.2
Prepare sulfur dioxide by burning sulfur: 3.51.0
Properties of sulfur dioxide and sulfites: 12.18.4
Reduce iron (III) chloride with sulfur dioxide: 3.51.3
Reduce potassium manganate (VII) with sulfur dioxide: 3.51.2
Tests for sulfur dioxide: 3.51.1
Tests for substances with dilute hydrochloric acid: 11.3.5 (See 3. and 4.)

13.1.36 Sulfur trioxide, SO3
Highly irritating and toxic gas, use a fume cupboard, fumes may affect asthmatics
Sulfur trioxide with water forms sulfurous acid.
Loss of substance on heating indicates: 12.11.3.4.1 (See 6.)
Reactions of sulfamic acid: 12.18.7, (See: 4. Test for sulfur trioxide)
Sulfuric acid, ionization of hydrogen sulfate ion, sulfur trioxide: 12.6.0.0

13.2.0 Air, atmosphere
Air Experiments
1.0 Air (Primary)
3.48 Acid rain and nitrogen oxides, NOx
31.7.18 Air bubbles in oil show electric field
37.9 Air expands when heated
13.2.23 Air flow, hair dryer, vacuum cleaner, ping-pong ball, balloon
4.227 Air pressure in all directions
12.1.04 Atmospheric pressure water spray
12.3.3 Air has mass, air has weight, with a balloon
4.226 Air has mass, break ruler under newspaper
18.6.0, Air pollution
12.3.2.1 Air takes up space, transfer air under water
27.58 Air wedge interference
12.3.0 Atmospheric pressure, air pressure
13.2.0 Bernoulli force, Bunsen burner, air streams
20.0.9 Breath, Henry's law and decompression sickness, "the bends"
9.237 Breath, oxygen content of inhaled and exhaled air
9.239.1 Breath, simulated diaphragm breathing
9.241 Breath volume
6.5.1 Carbon dioxide in the air is necessary for photosynthesis
35.23.1 Coal, coal dust explosions
37.12 Cold air is heavier than warm air, inverted paper bag balance
37.42.1 Composition of the atmosphere and greenhouse gases
Density of gases, Air (Table)
11.3.20 Density of air with a balloon
37.5.4 Dust in the air
4.8.0 Expansion of air
36.26.1 Falling stars, shooting stars, meteors, air resistance
16.3.9 Fan on a sailing boat, Newton's sailboat, fan cart
3.4.6 Gas or vapour inhalation, EAR, CPR
20.1.05 Hot air balloons
37.47 Hot air rising
37.39.1 Layers of the atmosphere, lapse rate, auroras
13.2.6 Lift water by blowing, Atomizer, Venturi tube, fly spray, chimney
9.10.5 Marcotting, air layering, (Agriculture)
37.7.0 Moisture enters the air, evaporation, (Weather)
37.8.0 Moisture leaves the air, precipitation, (Weather)
4.241 Oxidation and air pressure, steel wool over water
4.223 Plastic syringes and air pressure, Boyle's Law
17.2.3 Reduce friction with air streams, balloon hovercraft
26.3.2 Resonance in air columns, musical instruments
3.34 Soil air (Primary)
36.7.0 Sunrise & sunset
12.3.2.1 Transfer air under water, air takes up space
37.11.0 Winds, convection
4.238 Volume and pressure of air, Boyle's Law

3.32.0 Prepare gases with gas generation apparatus
| 12.3.2 Saturation vapour pressure over water | See diagram 1.13a: Simple fume hood
| See diagram 3.32: Gas generation apparatus | See diagram 3.33: Collect gas with an angle tube syringe
Be careful! Prepare poisonous gases only in a fume hood or fume cupboard, e.g. preparation of chlorine, hydrogen chloride,
sulfur dioxide.

1. See diagram 3.32, No. 1. Collect more dense gas by upward displacement of air (downward delivery), if molecular mass > 28.8.
The more dense gas sinks down into, and displaces, the less dense air upwards, e.g. preparation of carbon dioxide, nitrogen dioxide.

2. See diagram 3.32, No. 2. Collect less dense gas by downward displacement of air (upward delivery), if molecular mass < 28.8.
The less dense gas rises into, and displaces, the more dense air downwards, e.g. preparation of hydrogen, ammonia.
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 mostlyair.
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. See diagram 3.32, No. 3. Collect insoluble gas, or not very soluble gas, of any density over water, i.e. by downward displacement
of water. (water displacement), e.g. preparation of oxygen.
Fill one third of the water trough.
Fill a test-tube with water by placing 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.

4. See diagram 3.32, No. 4. Collect soluble gases in water (aqueous solution), e.g. Cl2.

5. See diagram 3.32, No. 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, use an inverted glass filter funnel.

6. See diagram 3.33, No. 6. Collect gas with an angle tube syringe that can collect and store gas of any density.

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 Lime water tests for carbon dioxide
See diagram 3.34.1: Lime water tests for carbon dioxide
Lime water test A The lime water turns milky.
1. Prepare the weak alkali calcium hydroxide solution, lime water, by adding solid calcium hydroxide, slaked lime, to demineralized
water.
Shake the solution vigorously and leave to stand.
Calcium hydroxide solid is only slightly soluble in water.
When a white solid has settled as a fine white sediment, decant the clear lime water above the sediment.
To replenish the lime water, add more demineralized water to the sediment in the stock bottle, shake and allow to settle.
The settling process may take several days.

2. Prepare lime water 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.
Lime water 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 lime water.
otherwise the surface turns milky on standing because of the reaction with the carbon dioxide in the air.
Store lime water in a container with a rubber or plastic stopper.
f you use a screw top container, calcium carbonate may form in the screw of the lid so you cannot open the container.

3. Pass carbon dioxide through lime water or blow through it.
Carbon dioxide turns lime water milky.
A fine suspension of calcium carbonate causes the milky colour in the
Ca(OH)2 (s) + CO2 (g) --> CaCO3 (s)+ H2O (l)
Pass more carbon dioxide through the solution.
The solution becomes clear again because soluble calcium hydrogen carbonate forms.
CaCO3 (s) + CO2 (g) + H2O (l) <--> Ca(HCO3)2 (aq)
The reaction can be reversed to form the calcium carbonate precipitate once again.
either by boiling the solution or bubbling air through it.

4. Pass air through freshly-made lime water.
After a long time may see a faint cloudy precipitate.
The air contains about 0.4% carbon dioxide.

5. Use a drinking straw to exhale into the lime water.
A cloudy precipitate soon forms because exhaled breath contains about 4% carbon dioxide.

6. Pass carbon dioxide into water to form carbonic acid.
CO2 (g) + H2O (l) --> H2CO3 (aq)
Add lime water to neutralize the carbonic acid to form the carbonate ion
H2CO3 (aq) + 2OH- (aq) --> CO32- (aq) + 4H2O (l)
Calcium carbonate is insoluble and precipitates
Ca2+ (aq) + CO32- (aq) --> CaCO3 (s)
Pass more carbon dioxide into the solution to use up the OH- and make the solution acidic.
so the carbonate ion is converted into the soluble bicarbonate ion.
CaCO3 (s)+ H2CO3 (aq) --> Ca2+ (aq) + 2HCO3- (aq)

3.34.1.3 Burning charcoal tests for carbon dioxide
Put lime water 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 lime water.
Remove the burning charcoal.
Close the container and shake it.
The solution turns a milky colour.
The formation of this white solid in lime water 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 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, 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 mL 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
Phenol red, C19H14O5S (acid-base indicator): 28
See diagram 3.34.1: Phenol red test
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.

3.39.0 Carbon monoxide, CO, properties
Carbon monoxide, carbon oxide, carbonic oxide (carbonyl, CO), monodentate ligand
Be careful! Do NOT make carbon monoxide.
See 18.6.3: Danger of vehicle exhausts, tailpipe gases

1. Toxicity and air pollution
Carbon monoxide is very toxic.
It can cause unconsciousness due to combination of the gas with haemoglobin in the blood, and prevents the blood from acting as an
oxygen carrier.
Death can occur from carbon monoxide inhalation.
Do not prepare carbon monoxide in an open room.
Carbon monoxide is particularly dangerous because it is a colourless, odourless and tasteless gas.
It kills more people than any other gas.
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.
It is very flammable and forms explosive mixtures with air.
It forms when carbon in fuels (petrol, wood, coal, natural gas) is not burned completely.
It is soluble in some organic solvents, such as ethyl acetate, chloroform and acetic acid.
It forms toxic and flammable compounds when exposed to finely dispersed metal powders.
It may react vigorously with oxygen, acetylene, chlorine, fluorine and nitrous oxide.
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.
Metal oxides are reduced by passing carbon monoxide over the heated oxide.
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)

3.39.2 Phosgene, carbonyl chloride, CoCl2
Phosgene is used for the synthesis of isocyanate-based polymers, carbonic acid esters, acid chlorides, dyestuff manufacture and
insecticides.
It is a poisonous gas and was used as a chemical warfare agent in the First World War.
In low concentrations it smells like cut grass.
It is manufactured by direct combination of carbon monoxide with chlorine, with carbon catalyst.
At high temperate the reverse reaction occurs
CO + Cl2 --> CoCl2
Chloroform + UV --> phosgene, so chloroform kept in dark glass bottles

3.41.0 Hydrogen gas
Hydrogen gas, H2, was first discovered by Theophrastus Paracelsus, (1493-1541, Swiss- German), and Robert Boyle (1627-1691.
England), who demonstrated the existence of "inflammable air", both using sulfuric acid with iron.
In 1766, Henry Cavendish (1731-1810, England), using iron with hydrochloric acid, showed that hydrogen is an element and a
constituent of water.
Hydrogen gas, both as a compressed gas and a gas generated by experiment is highly flammable and can act as a non-toxic asphyxiant.

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.

3.41.2.0 Prepare hydrogen gas
See diagram 3.41: Collecting hydrogen gas
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.
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.

Experiments
1. Prepare hydrogen gas by reaction of an acid on an active metal, e.g. zinc with hydrochloric acid.
See diagram 3.2.33: Zinc with hydrochloric acid
Do not use a container bigger than a test-tube.
Check that the test-tubes are not chipped or cracked.
Put granulated zinc in a small borosilicate test-tube and cover it with water.
Add a crystal of copper (II) sulfate to act as a catalyst.
The copper sulfate first reacts with zinc in a displacement reaction to form metallic copper, which has a catalytic action.
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.
Use a lighted taper to "pop"' the air / hydrogen mixture.
Collect hydrogen gas by downward displacement or over water.
Let the reaction continue for some minutes to drive out all the air fromthe 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. Prepare hydrogen with iron filings and citric acid or 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 is colourless and odourless, but any impurities in the iron filings give a nasty smell.
To test 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 one test-tube full of hydrogen gas!

3. Prepare hydrogen with aluminium foil and sodium carbonate
Cut into small pieces aluminium foil or aluminium milk bottle tops and put them into a test-tube.
Add 5 mL of sodium carbonate solution (Na2CO3.10H2O, washing soda).
Heat until effervescence occurs.

4. Prepare hydrogen with calcium and 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. Prepare hydrogen with iron filings and potash 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, also shown as KAl(SO4)2.12H2O.

6. Prepare hydrogen with iron filings and 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. Prepare hydrogen with aluminium kitchen foil and caustic soda drain cleaner (sodium hydroxide)
Fill a clear glass 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 × 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

3.41.2.1 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 Burn hydrogen gas
See diagram 3.41.3: Reduction tube
Safety
For these experiments, wear safety goggles, ear protection and have afire extinguisher close by.
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.
Prepare hydrogen gas by reaction of an acid on an active metal, e.g. zinc with hydrochloric acid.
Hydrogen forms spontaneously combustible / explosive mixture in air at low temperature.
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 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.

Hydrogen in test-tubes
Generate small quantities of gas in a test-tube only.
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.
Students may be injured by "popping" (burning) hydrogen in other glass containers.
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.

3.41.4 Hydrogen gas 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 sulfuric 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.

3.41.5 Explode a hydrogen gas balloon
Never explode hydrogen in a glass or solid container.
Exploding a hydrogen balloon is safe, but loud, because there are no hard pieces of shrapnel from the explosion.
Attach a string to the balloon clip and tether it to a rail away from any combustible material.
All other observers should be at least 5 metres away from the explosion.
Attach a small birthday candle to a long stick or thin dowel to be a safe distance from the ignition.
Extend your arms and hold the lighted candle at the end of the dowel below the balloon.
Tie the string of the 30 cm hydrogen balloon to a weight on the desk, well away from any chemicals or combustible substances.
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

3.41.7 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).

3.42 Hydrogen chloride
Hydrogen chloride, HCl, Toxic, corrosive, highly irritating gas
Hydrogen chloride gas is corrosive.
Do not prepare hydrogen chloride in an open room.
Use a fume cupboard.
Be careful! Do these experiments 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.

3.42.0 Prepare hydrogen chloride / hydrochloric acid
| See diagram 3.2.36: Collecting hydrogen chloride
| See diagram 3.42: Collecting hydrogen chloride
| See diagram 13.5.2: Prepare hydrochloric acid
| See diagram 1.13a: Simple fume hood
Be careful!
Prepare hydrogen chloride gas only in a fume hood or fume cupboard!
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.

1. 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)
Repeat the experiment with concentrated hydrochloric acid and concentrated sulfuric acid.
Be careful!

2. Prepare hydrogen chloride gas by gently warming hydrochloric acid in a water bath in a flask with a gas collection tube.
Collect the gas by displacement of air.
Hydrogen chloride can be used in place ammonia in the ammonia fountain in the ammonia fountain experiment.

3. Hydrogen chloride gas fumes in air, forming droplets of hydrochloric add, so be careful not to inhale it.
Mix well together, on a creased sheet of paper, a finger width of sodium hydrogen sulfate and the same quantity of sodium chloride.
and transfer the mixture to a test-tube fitted with stopper and delivery tube.
Heat the mixture by keeping the test-tube moving in the flame to prevent the glass cracking.
The misty fumes of the heavy gas pass down wards into the second test-tube.
When it is full, as shown by the fumes coming out at the top, stopper it, and collect another as a spare.
Hold a piece of blue litmus paper in the fumes.
The blue litmus turns red.
The misty fumes are minute droplets of hydrochloric acid.
formed by the reaction of the invisible hydrogen chloride with water vapour in the air.
It is this acid that turns the blue litmus red.

4. Concentrated sulfuric acid reacts with sodium chloride to form hydrogen chloride gas, and can be reduced with copper metal to
form sulfur dioxide gas on gentle heating.
Do these experiments in a fume cupboard while wearing eye protection.

5. Prepare hydrochloric acid gas from sodium chloride.
Mix together 2 g of sodium chloride and 2 g of powdered alum (hydrated potassium sulfate), or sodium hydrogen sulfate
(sodium bisulfate).
Put the mixture into a dry test-tube and have ready a damp blue litmus paper and a bottle of strong ammonia.
Heat the mixture over a medium Bunsen burner flame, holding the test-tube in a paper holder and moving the test-tube in the flame.
Hydrochloric acid gas, or hydrogen chloride forms as steam-like fumes.
Sniff the gas cautiously and put the blue litmus paper into the fumes.
Test the gas by removing the stopper from the bottle of strong ammonia and blowing the steamy fumes across the top of the bottle.
A dense white smoke forms.
The white smoke consists of ammonium chloride.

6. Prepare hydrochloric acid
Repeat the above experiment 5.
Lead the delivery tube into a 500 mL bottle, half full of water.
Keep the end of the tube clear of the water to prevent sucking back.
As you heat the test-tube, hold the bottle in your other hand, and keep the water swirling to dissolve the hydrogen chloride gas.
Continue heating until no more gas forms.
Recharge the test-tube and repeat the procedure many times.
Label the bottle of dilute hydrochloric acid.

3.42.1.1 Solubility test for hydrogen chloride
Remove the stopper from the receiving test-tube under water.
Note the solubility of hydrogen chloride.
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.
To show the extreme solubility of hydrogen chloride, remove the stopper from the test-tube, and quickly put your thumb or finger
over the mouth of the test-tube.
Invert the test-tube of gas in a dish of water, removing your thumb only when the mouth of the test-tube is under the water.
Water rushes up into the test-tube.
Hydrogen chloride is so soluble that it dissolves almost at once in the water at the mouth of the test-tube.
Atmospheric pressure forces the water into the empty test-tube.

3.42.1.2 Litmus paper test for hydrogen chloride
1. Test the solution in the receiving test-tube with moist litmus paper.
Red litmus paper turns blue.

2. Hold a piece of blue litmus paper in the fumes.
The blue litmus turns red.
The misty fumes are minute droplets of hydrochloric acid, formed by the reaction of the invisible hydrogen chloride with water
vapour in the air.
It is this acid that turns the blue litmus red.

3.42.1.4 Lighted splint test for hydrogen chloride
Hydrogen chloride extinguishes a lighted splint.
Hydrogen chloride neither burns nor supports combustion.

3.42.1.5 Magnesium ribbon test for hydrogen chloride
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.

3.42.1.6 Ammonia test for hydrogen chloride
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.

3.42.1.7 Fountain test for hydrogen chloride
1. 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 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.

2. Fill a beaker with litmus solution.
Fit a glass jet tube into the stopper of a flask.
Remove the stopper and jet, and start filling the dry flask with hydrogen chloride.
When the flask is full of gas, replace the stopper and jet, and quickly invert the flask with the other end of the jet tube in the litmus
solution.
With the spirit burner at a safe distance, pour a finger width of methylated spirit on the flask and blow on it.
This causes the spirit to evaporate and thereby cool the flask and the gas inside it.
The contraction of the gas reduces its pressure, and atmospheric pressure forces litmus solution up the glass tube and out of the jet.
The fountain from the jet suddenly increases and the litmus changes colour.
The fountain from the jet suddenly increases for the reason given above.
The litmus changes from blue to red because the water in the litmus solution reacts with the hydrogen chloride to form hydrochloric acid.

3.43.0 Prepare hydrogen sulfide
See diagram 1.13a: Simple fume hood
Hydrogen sulfide, Extremely Toxic, Highly flammable
Hydrogen sulfide, gas, < 1% Not hazardous
Hydrogen sulfide water, solution, Toxic if ingested
Hydrogen sulfide gas is both an irritant and an asphyxiant.
Collapse, coma and death from respiratory failure may come within a few seconds after one or two inspirations.
Do not prepare hydrogen sulfide in an open room.
Hydrogen sulfide is extremely toxic so use a 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.
Hydrogen sulfide has high acute short-term toxicity to aquatic life, birds, and animals.
It is soluble in water and organic solvents and will corrode metals.
Hydrogen sulfide is used in the manufacture of pulp and paper (digesting agent), in tanneries and in sulfide ores.
It is the main offensive smell in flatus produces after a diet of certain types of beans.

Experiment
1. This experiment is regarded as the safest method to prepare hydrogen sulfide gas.
Do this experiment in a fume cupboard, fume hood.
Prepare only a very small quantity of this gas.
Have a beaker full of a weak alkali ready to stop the reaction.
Add dilute hydrochloric acid to iron (II) 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) + 3O2 (g) --> 2SO2 (g) + 2H2O (l)

2. 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.

3.43.1 Tests for hydrogen sulfide solution
See diagram 1.13a: Simple fume hood
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 pink-red.

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+ + S2-5.
Hydrogen sulfide with lead ethanoate paper forms a black precipitate of lead sulfide.
H2S (g) + Pb (C2H3O2)2 (aq) --> PbS (s) + 2HC2H3O2 (aq)
lead (II) acetate + hydrogen sulfide --> lead (II) sulfide + acetic acid

3.43.2 Reduce potassium manganate (VII) with hydrogen sulfide
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.

13.1.01 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.
They 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 secure 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. cables around a motor car engine.

13.1.02 Relative molecular mass of gases, propane
See diagram 13.1.5: Relative molecular mass of gases | 12.3.2: 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 × 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.

13.1.6 Molar volume of oxygen prepared with hydrogen peroxide
See diagram: 13.1.6: Molar volume of oxygen | 12.3.2: Saturated vapour pressure over water
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 tube1 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 × V1) / T1 = (P2 × V2) / T2
(PO2 × VO2) / Tf = (Pstp × Vstp) / Tstp
So Vstp = VO2 [(PO2 / Pstp) × (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) × (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

13.9.0 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.

13.9.1 Nitrogen reacts with metals
Lithium is the only element that reacts with nitrogen at room temperature.
6Li (s) + N2 (g) --> 2Li3N (s)
Magnesium and other alkali metals react with nitrogen at high temperature to form nitrides.
3Mg (s) + N2 (g) --> Mg3N2 (s)

13.9.3 Nitrogen gas generated in a motor car air bag
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 air bag.
2NaN3 --> 2Na + 3N2 (300oC)
10Na + 2KNO3 --> K2O + 5Na2O + N2
K2O + Na2O + SiO2 --> alkaline silicate

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)

13.13.8 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)