UNChem1a

School Science Lessons
Chemistry
2012-05-20 SPwp
Please send comments to: J.Elfick@uq.edu.au

Table of contents
3.33.0 Ammonia
3.33.1 Tests for ammonia
3.33.2 Ammonium compounds
3.34.0 Carbon dioxide
3.34.1 Prepare carbon dioxide
3.34.1.0 Tests for carbon dioxide
3.42.0 Hydrogen chloride
3.42.1 Tests for hydrogen chloride
3.43.0 Hydrogen sulfide
3.43.1 Tests for hydrogen sulfide
3.32.0 Prepare gases with a gas generation apparatus

3.33.0 Ammonia gas, NH₃
Ammonium compounds: 3.33.2
Ammonium ion (NH₄⁺), solution, Nessler's reagent, Tests for ammonium radical, NH₄⁺: 12.11.3.10.2, (See 5.)
Ammonium solution, ammonium hydroxide, concentrated, aqueous ammonia, spirits of hartshorn, Toxic by all routes, Lung and eye irritant
Ammonia solution (10% ammonia solution): 5.3.4
Ammonium hydroxide (>25% in water), solution of ammonia in water, ammonia hydrate
Ammonium hydroxide, prepare by mixing household ammonia cleaner with strong ammonia with water.
Ammonia (25-28%) Specific gravity (relative density 0.91)
Ammonia absorbs impurities on charcoal
Ammonia anhydrous, ammonia liquefied gas, ammonia gas, Ammonia, anhydrous, hazards: 3.8.1
Ammonia cell test (0.2-10 mg / l)
Ammonia cell test kit (1-28 mg / l)
Ammonia, NH_3, and the ammonium ion, NH₄⁺: 12.11.3
Ammonia solution (10% ammonia solution): 5.3.4
Ammonia solution < 2 M, ammonium hydroxide, < 5%, Not hazardous, > 18 M, 35-50%, > 50%, 2-5 M, 5-10%, ammonium hydroxide
Ammonia solution 3 M: dilute 200 mL (28%) 14.8 M concentrated solution to 1 litre of water
Ammonia solution 4 M: dilute 220 mL (28%) 18 M concentrated solution to 1 litre of water ("ammonium hydroxide")
Ammonia solution 5-18 M, 10-35%, ammonium hydroxide, cloudy ammonia
Ammonia solution, concentrated, solution 35% w / w, SG 880, 880 ammonia, conc., liquor ammonia fort
Ammonia solution, For 2 M soln., dilute 330 mL of 10% soln in 1 L water
Ammonia solution in a neutralization reaction: 3.33.3
Ammonia with sulfuric acid, acid-base neutralization: 12.8.1
Ammonia with copper sulfate solution: 3.33.4
Ammonia with cobalt chloride solution: 3.33.5
Catalytic oxidation of ammonia forms nitrogen monoxide, with red-hot platinum wire: 13.6.6.1
Catalytic oxidation of ammonia, with chromium (III) oxide catalyst: 13.6.6.2
Catalytic oxidation of methyl alcohol, and ammonia solution: 17.3.12
Group 6 Tests for K⁺, Mg²⁺, Na⁺, NH₄⁺: 12.11.4.6
Iron (II) sulfate with ammonia: 12.1.2
Magnesium sulfate with ammonia: 12.1.10
Prepare ammonia: 3.33.0
Prepare ammonia, Haber process: 3.33.01
Prepare chrome alum: 12.14.4
Prepare hydroxides with ammonia solution: 3.33.2
Prepare iron (III) ammonium alum: 12.14.3
Prepare rayon, copper (II) sulfate with ammonia solution: 3.4.8.0
Prepare rayon, basic copper carbonate with ammonia solution: 3.4.8.1
Reactions of ammonium salts and potassium salts with water: 14.2.1
Reduce copper (II) oxide to copper with ammonia: 13.6.7
Store alkalis: 2.0. (Ammonia)
Tests for ammonia, ammonium ions: 12.11.3.10.2

3.33.1 Tests for ammonia
3.33.1.6 Ammonia fountain test
3.33.1.1 Concentrated hydrochloric acid test (hydrogen chloride test)
3.33.1.3 Litmus tests for ammonia
3.33.1.2 Odour test for ammonia
3.33.1.5 Sodium bicarbonate test for ammonia
3.33.1.4 Solubility tests for ammonia
13.6.5 Tests for ammonia, and hydroxyl ions (hydroxide ions)
12.11.3.10.2 Tests for ammonium radical, NH⁴⁺

10.1.1.1 Tests for ammonia
3.33.1.2 Odour test
Put a drop of aqueous ammonia solution in a test-tube with a stopper. Shake the test-tube and warm it with the hands. Note the pungent odour of ammonia.
See diagram 1.13: Smelling test
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. Add powdered wood charcoal. Shake the test-tube and smell the contents again. Filter the mixture. Heat the charcoal and note the absence of the pungent odour of ammonia.

3.33.2 Ammonium compounds
Ammonium ion NH₄⁺
Ammoniacal silver nitrate solution, Tollen's reagent
Ammonium acetate, ammonium ethanoate
Ammonium alum, aluminium ammonium sulfate (VI)-12-water, AlNH₄(SO₄)₂.12H₂O, Tschermigite mineral
Ammonium alum, Alums: 7.2.2.1a
Ammonium benzoate
Ammonium bicarbonate, ammonium hydrogen carbonate, NH₄HCO₃, white crystals, crystal ammonia, used in some baking powders.
Ammonium bromide, used in smelling salts
Ammonium carbonate, smelling salts: 12.16.3.5
Ammonium chlorate, Toxic, Explosive
Ammonium chloride, smoke screen: 3.33.6
Ammonium chloride, NH₄Cl, sal ammoniac, Harmful if ingested: 3.33.2.2
Ammonium chloride, Common ion effect in ammonium chloride solution: 17.5.5.6
Ammonium chromate, NH₄)₂CrO₄, yellow monoclinic crystals, decomposes at 185^oC, loses ammonia gas, alkaline solution, needs cool storage.
Ammonium copper (II) chloride, tetrammine copper (II) chloride
Ammonium cupric sulfate, diammonium copper (II) sulfate (VI)-6-water, Cu(NH₃)₄SO₄.H₂O, in fire extinguishers, Environment danger
Ammonium dichromate
Ammonium dihydrogen phosphate, ammonia dihydrogen orthophosphate (V), ammonium biphosphate, Harmful if ingested
Ammonium dihydrogen phosphate, monobasic ammonium phosphate, (piezoelectric crystal in microphones and transducers)
Ammonium ethanoate, CH₃CO₂NH₄, ammonium acetate, ammonium oxalate
Ammonium ferric citrate (brown)
Ammonium ferric citrate (green)
Ammonium ferric sulfate, iron (III) ammonium sulfate
Ammonium ferrous (II) sulfate, iron (II) ammonium sulfate
Ammonium ferrous sulfate (COD determination)
Ammonium fluoride
Ammonium formate
Ammonium hexafluorovanadate, ammonium hexafluoridovanadate(V), (NH₄)₃VF₆
Ammonium hydrogen carbonate, NH₄HCO₃, ammonium bicarbonate, Smelling salts (ammonium carbonate): 12.16.3.5
Ammonium hydrogen carbonate, Baking powder: 19.1.6.1
Ammonium hydrogen difluoride
Ammonium hydrogen phosphate, diammonium hydrogen phosphate, (NH₄)₂HPO₄, in fertilizers, Toxic if ingested
Ammonium hydrogen sulfate
Ammonium iodide, Harmful if ingested
Ammonium iron (II) sulfate
Ammonium magnesium chloride
Ammonium magnesium phosphate, struvite mineral, ((NH₄)MgPO₄.6H₂O), (used in phosphate recycling from sewage)
Ammonium metavanadate, ammonium vanadate, NH₄VO₃
Ammonium molybdate, (NH₄) ₆Mo₇O₂₄.4H₂O, H₂₄Mo₇N₆O₂₄, Irritant, Harmful if ingested: 5.4.1
Ammonium molybdate (VII) 4-water, powder, ammonium paramolybdate
Ammonium molybdate catalyst, Hydrogen peroxide with sodium thiosulfate: 17.3.9
Ammonium molybdate, Tests for phosphates: 12.11.5.15
Ammonium nickel sulfate, ammonium nickel (II) sulfate 6H₂O, Harmful if ingested
Ammonium nitrate
Ammonium oxalate, C₂O₄(NH₄)₂.2H₂O, ammonium ethanedioate-1-water, For 0.1 M soln., 16 g in 1 L water, Harmful if ingested
Ammonium perchlorate, Toxic if ingested, Explosive, Not permitted in schools
Ammonium persulfate
Ammonium phosphate, triammonium phosphate(V)-3-water, ammonium phosphate tribasic, (NH₄)₃PO₄
Ammonium sodium hydrogen phosphate (V)-4-water, Na(NH₄)HPO₄.4H₂O, ammonium sodium hydrogen orthophosphate, microsmic salt (from urine)
Ammonium purpurate, (Murexide)
Ammonium salts, Reactions of ammonium salts and potassium salts with water: 14.2.1
Ammonium sodium hydrogen phosphate (V)-4-water, Na(NH₄)HPO₄.4H₂O, ammonium sodium hydrogen orthophosphate, microsmic salt (from urine)
Ammonium sulfate
Ammonium sulfamate
Ammonium sulfide solution, (>20% in water), (10% solution), (20% solution), Highly toxic fumes
Ammonium tartrate
Ammonium thiocyanate, NH₄NCS, Highly toxic fumes
Ammonium thiocyanate, Tests for iron: 12.11.3.10.12
Ammonium thiosulfate, diammonium thiosulfate, (NH₄)₂S₂O₃, photography rapid fixer, fertilizer, reducing agent, Environment danger
Ammonium triiodide, nitrogen triiodide
Ammonium vanadate, ammonium metavanadate, Harmful if ingested
Diammonium copper (II) sulfate (VI)-6-water, ammonium cupric sulfate
Diammonium hydrogen phosphate, ammonium phosphate dibasic
Monammonium phosphate
Tests for ammonium compounds with sodium hydroxide solution: 12.11.3.7

3.33.2.1 Ammonium carbonate, (NH₄)₂CO₃ or (NH₄)₂CO₃.3H₂O, crystal ammonia, white crystalline solid, Harmful if inhaled
Harmful if swallowed or inhaled, irritates skin, eyes, respiratory tract, 3.33.2.1
Acids with ammonium carbonate: 12.12.3
Alkalis with ammonium carbonate: 12.12.2
Ammonium carbonate, For 2 M soln., 300 g in 450 mL 10% NH₃, then dilute in 1 L water
Ammonium carbonate with ethanoic acid: 14.2.3
Ammonium carbonate solution precipitates metal carbonates: 12.16.3.4
Ammonium carbonate with acids: 12.16.3.3
Ammonium carbonate with alkalis: 12.16.3.2
Decompose ammonium carbonate: 12.12.1
Heat ammonium carbonate, (smelling salts): 12.16.3.1
Precipitates with ammonium carbonate: 12.12.4
Smelling salts: 12.16.3.5

3.33.2.2 Ammonium chloride, NH₄Cl, sal ammoniac, torch battery, dry cell, soldering flux, Harmful
Ammonium chloride, NH₄Cl, For 5 M soln., 270 g in water
Ammonium chloride, Heat sensitive paper: 5.3.10
Ammonium chloride, Secret writing inks: 3.2.5.5
Ammonium chloride smoke screen: 3.33.6 Ammonium chloride, Solution < 20%, Not hazardous
Decomposition of ammonium chloride: 3.30.15
Dry cell using ammonium chloride, Leclanche cell: 3.88
Fireproof cloth: 3.5.11
Heat ammonium chloride crystals: 8.4.1
Heat sensitive paper, cobalt (II) chloride, ammonium chloride (sal ammoniac): 5.3.10
Hydrolysis of ammonium chloride: 12.10.3.2
Hydrolysis of sodium carbonate: 12.10.3
Thermal dissociation of ammonium chloride, effect of temperature: 17.5.5.5

3.33.2.2a Ammonium dichromate, (NH₄)₂Cr₂O₇, ammonium dichromate (VI), (NOT "ammonium bichromate"), Thermal decomposition: 3.30.14
Ammonium dichromate, powder, solid, Highly toxic if ingested, Corrosive, Oxidizing (OXD 1439)
Ammonium dichromate, solid, Toxic, Corrosive, Oxidizing (OXD 1439)
Ammonium dichromate, Solution < 0.5%, Not hazardous

3.33.2.2b Ammonium iron (II) sulfate, ammonium iron (II) sulfate 6-hydrate, (NH₄)₂SO₄FeSO₄.6H₂O, Mohr's salt, ferrous ammonium sulfate
Ammonium iron (II) sulfate, iron (II) ammonium sulfate, For 0.1 M soln., 39.2 g in water, add 5 mL conc. H₂SO₄ in 1 L water
Ammonium iron (II) sulfate, Secret writing ink using ammonium iron (II) sulfate
Ammonium iron (III) sulfate, iron (III) ammonium sulfate, (std soln)
Ammonium iron (III) sulfate, NH₄Fe(SO₄)₂.12H₂O, iron (III) ammonium sulfate, ferric ammonium sulfate (FAS), iron alum: 12.1.6
Ammonium iron (III) sulfate, diamond-shaped crystals, similar to potash alum and chrome alum, can be grown on a piece of cotton
Ammonium iron (III) sulfate, Prepare ferric tannate with tea leaves: 16.8.2 (See 4.)
Ammonium iron (II) sulfate, Secret writing inks: 3.2.5

3.33.2.3 Ammonium nitrate, NH₄NO_3,(OXD 1942), Toxic if ingested
Ammonium nitrate, colourless crystals, soluble in water and ethanol, nitrate of ammonia garden fertilizer, in explosives
Ammonium nitrate must not be mixed with diesel oil to make an explosive mixture.
Ammonium nitrate cold pack: 14.2.4
Ammonium nitrate, Movement of copper ions in ammonium nitrate solution: 3.4.3
Ammonium nitrite, Toxic if ingested, Prepare nitrogen gas with ammonium nitrite: 3.46

3.33.2.3a Ammonium persulfate, (NH₄)₂S₂O_8,ammonium peroxydisulfate, Toxic if ingested
Ammonium persulfate, Solution < 25%, Toxic if ingested
Ammonium persulfate, Etchants: 7.9.20.1
Ammonium persulfate, Persulfate-iodide clock reaction: 17.1.10

3.33.2.4 Ammonium sulfate, (NH₄)₂SO₄, sulfate of ammonia, analytical reagent, molecular weight 132.13, Harmful if ingested
Ammonium sulfate, (agriculture) Straight fertilizers and mixed fertilizers: 6.9.17.1
Ammonium sulfate, Ammonia with sulfuric acid, acid-base neutralization: 12.8.1
Ammonium sulfate, For 0.1 M soln., 13.2 g in 1 L water
Ammonium sulfate, Substances that decompose and lose mass when heated, thermal decomposition: 3.30.0
Prepare ammonium sulfate by neutralization: 12.14.1

3.34.0 Carbon dioxide
Carbon dioxide, CO_2:3.34
Alkalinity, total alkalinity and buffer capacity: 18.7.9
Alkalis with acidic oxides, carbon dioxide: 12.7.6
Baking powder: 19.1.6.1
Balloon as a sound lens, acoustic lens: 26.7.3
Caffeine, extraction with supercritical carbon dioxide: 19.4.2.3.0
Carbon dioxide, acidic oxides, (non-metal oxides): 12.17.3
Carbon dioxide affects acid-base titration: 12.8.5
Carbon dioxide and photosynthesis: 3.36
Carbon dioxide and respiration: 3.37
Carbon dioxide and fermentation for brewing: 3.38
Carbon dioxide cartridge rocket, Dangerous experiment: 16.6.4.4
Carbon dioxide has mass: 12.3.3.1
Carbon dioxide, hazards: 3.8.2
Carbon dioxide in drinks chemistry: Unit 6.
Carbon dioxide in the air is necessary for photosynthesis, Nasturtium: 6.5.1
Carbon dioxide in the home: 3.35.0
Carbon dioxide is a product of combustion: 8.6.3
Carbon dioxide gas is produced during respiration: 5.05
Carbon dioxide, solid, Dry ice, observe frozen carbon dioxide, (hot ice): 3.34.5
Carbon dioxide soda syphon bulbs: 3.3.5
Carbon dioxide, Tests for carbon dioxide: 3.34.1.0
Carbon dioxide with barium hydroxide solution, ionization of barium hydroxide: 12.17.3.2
Carbon dioxide with calcium carbonate suspension: 12.16.1
Carbon dioxide with calcium hydroxide solution (limewater), tests for carbon dioxide: 12.16.1.1
Carbon dioxide with sodium hydroxide solution: 12.17.3.1
Carbon dioxide with water, thymolphthalein indicator: 3.34.1.6
Carbon dioxide with water, phenol red indicator: 3.34.1.7
Carbonic acid, H₂CO₃, ionization reaction of carbonic acid: 12.3.0.5
Compare the weight of carbon dioxide and air: 3.34.7
Composition of the atmosphere and greenhouse gases: 3.32.1, (See 2.)
Diffusion of heavier than air gas, carbon dioxide: 3.55.1
Dilute acids with acidic oxides, carbon dioxide, sulfur dioxide: 12.3.8
"Dry ice", "hot ice", Frozen carbon dioxide : 3.34.5
"Dry ice" in water: 3.34.5.1
Elimination of carbon dioxide during human respiration: 9.5.7
Greenhouse gases
Oxides and the periodic table: 12.17.1.1
Packaging gases, propellants, food additives: 19.4.4.22
Prepare carbon dioxide: 3.34.1
Production of carbon dioxide during plant respiration: 9.157
Reduce carbon dioxide by burning magnesium: 3.34.4
Respiration, aerobic respiration: 9.4.0
Respiration is a form of combustion: 8.6.5
Simulated boiling: 13.7.13
Solubility of acidic oxide carbon dioxide in water, acidity of soda water: 3.34.3
Soda-acid fire extinguisher: 3.34.6
Syphon bulbs, safety: 3.3.5
Tests for carbon dioxide: 3.34.1.0
Yeast cells convert glucose to carbon dioxide gas and alcohol: 3.35.4

3.34.1 Prepare carbon dioxide
Carbon dioxide and fermentation for brewing: 3.38
Prepare carbon dioxide, acids with carbonates, acids with bicarbonates: 3.34.0
Prepare carbon dioxide, alum with baking soda: 13.7.9
Prepare carbon dioxide, heat carbonates: 13.7.6
Prepare carbon dioxide, heat hydrogen carbonates: 13.7.7
Prepare carbon dioxide, sodium hydrogen carbonate with sour milk, vinegar, fruit juice: 19.1.7
Prepare carbon dioxide, spearmint sweet, "Mentos", with cola drink: 13.7.8
Prepare self-leavened flour, "self-raising flour": 19.1.8

3.42.0 Hydrogen chloride
3.42 Hydrogen chloride, HCl, Toxic, corrosive, highly irritating gas
18.6.2 Air pollution from burning refuse
31.1.4.2 Aluminium foil precipitator
3.30.15 Decomposition of ammonium chloride, NH₄Cl
3.33.6 Ammonium chloride smoke screen
12.19.5.0 CFCs, chlorofluorocarbons (reaction of hydrogen chloride with methane)
10.1.2 Diffusion rates of ammonia and hydrogen chloride gases
17.5.5.5 Effect of temperature on chemical equilibrium, thermal dissociation of ammonium chloride
12.8.13 Heat hydrated iron chlorides
Hydrochloric acid (solution of hydrogen chloride gas)
3.8.6 Hydrogen chloride, anhydrous, hazards
3.42.1.7 Hydrogen chloride fountain test
3.42.0 Prepare hydrogen chloride, HCl
3.42.01 Prepare hydrochloric acid
12.19.8.3 Prepare iron (III) chloride, FeCl₃, See 2.
12.20.2 Prepare tin (IV) chloride, SnCl₄
12.19.8.1 Reactions of chlorides, Cl^-
12.18.5.2 Sulfuric acid acts as a displacer of acids from their salts, sodium chloride, sodium acetate, sodium formate
3.33.1.1 Tests for ammonia, Concentrated hydrochloric acid test (hydrogen chloride test)
3.42.1.0 Tests for hydrogen chloride
12.11.3.6 Tests for substances with hot concentrated sulfuric acid, note gas produced, See 1.
12.13.4 Water with chlorides of phosphorus, PCl₃, PCl₅
12.14.2.5 Zinc with copper in sulfuric acid

3.42.1 Tests for hydrogen chloride
3.42.1.6 Ammonia solution test for hydrogen chloride
3.42.1.3 Ammonium chloride test for hydrogen chloride
3.42.1.4 Lighted splint test for hydrogen chloride
3.42.1.5 Magnesium ribbon test for hydrogen chloride
3.42.1.2 Moist litmus paper test for hydrogen chloride
3.42.1.1 Solubility tests for hydrogen chloride

3.43.0 Hydrogen sulfide
3.43 Hydrogen sulfide, H₂S, sewer gas, stink gas, rotten egg gas, sulfuretted hydrogen
3.1 Hydrogen sulfide, warning when treating residues
12.3.15 Acids and with salts, See 5.
Coal gas
13.13.8 Dry hydrogen sulfide and dry sulfur dioxide will not react
4.206 Float eggs in water
12.11.7.4 Group IV Insoluble sulfides precipitated by hydrogen sulfide, ZnS, MnS, CoS, NiS
17.5.7.0 Group, Explanation of group analysis, See Group II, Group IV
19.7.3 Hair products
12.2.2.1 Heat iron with sulfur
12.12.4 Hydrogen peroxide reacts as an oxidizing agent
15.2.14 Hydrogen sulfide as a reducing agent
3.8.7 Hydrogen sulfide hazards
9.9.18 Hydroponics, soil-less culture solutions
18.2.2.2 Iron in drinking water
3.43.0 Prepare hydrogen sulfide, H₂S
12.18.2 Prepare sulfides, (See 3.)
12.19.6.1 Prepare hydrogen iodide, HI, (See 5.)
3.43.3 Reduce iron (III) chloride with hydrogen sulfide
3.44.1.0 Oxygenic phototropic bacteria
3.44.1.4 Purple nonsulfur bacteria
5.32 Protect mangroves, (See 5.) (Primary)
3.43.2 Reduce potassium manganate (VII) with hydrogen sulfide
3.43.3 Reduce iron (III) chloride with hydrogen sulfide
18.6.5 Smell of water, hydrogen sulfide
12.11.3.3 Solubility tests, prepare a solution for group analysis, (See 4.)
20.0.6 Standard temperature and pressure, STP, density of gases
3.7.16 Sulfides, hazards
16.2.8.2 Sulfides: RSR (where R is not H), old name: thioethers
3.43.1 Tests for hydrogen sulfide solution, ionization of hydrogen sulfide
12.11.3.5 Tests for substances with dilute hydrochloric acid, note gas produced, (See 5.)
12.11.3.6 Tests for substances with hot concentrated sulfuric acid, note gas produced, (See 7.)

3.33.0 Prepare ammonia
See diagram 3.2.37.1: Prepare ammonia | See diagram 3.2.37.2: Fountain experiment | See 3.32.0: Prepare gases with a gas generation apparatus
1. Put a mixture of calcium hydroxide and ammonium chloride into a test-tube to a depth of 4 cm. Fill a U-tube with lumps of calcium oxide mixed with cotton wool. The cotton wool is to prevent blocking of the U-tube. Gently heat the test-tube. The calcium oxide is to dry the ammonia gas. Collect the less dense gas (0.769 kg / m³) by downward displacement of air. Test whether the receiver test-tube is full by holding a piece of moist red litmus paper at the opening. Ammonia gas turns red litmus blue. Collect test-tubes of ammonia gas and apply stoppers.
2NH₄Cl (aq) + Ca (OH)₂ (s) --> 2NH₄OH (s) + CaCl₂ (aq)
then NH₄OH (s) --> NH₃ (g) + H₂O (l)
2. Repeat the experiment using sodium hydroxide
NH₄Cl (aq) + NaOH (s) --> NaCl (s) + H₂O (l) + NH₃ (g)
3. Heat a finger width of ammonium chloride mixed with an equal amount of calcium hydroxide in a large test-tube fitted with stopper and delivery tube. The test-tube should be inclined slightly, otherwise the water formed in the reaction can flow back into the hot end of the tube. Collect the gas by passing it upwards into another test-tube, since ammonia is less dense than air. A piece of moist red litmus paper, held at the mouth of the container, will show when each is full. Stopper and store the test-tubes of ammonia.
4. Mix a finger width of calcium hydroxide with an equal quantity of ammonium chloride on a sheet of paper. Rub the mixture with the glass rod, and notice the smell of ammonia gas evolved. Hold a piece of moistened red litmus paper over the mixture. The litmus turns blue. Ammonia is a gas which, when moist, has alkaline properties.
5. Prepare ammonia with ammonium chloride and sodium carbonate. Put 5 g of ammonium chloride (sal ammoniac) in 2 cm depth sodium carbonate (washing soda) solution. Heat the test-tube. Test for ammonia gas and with wet red litmus paper.
6. Prepare ammonia with ammonia solution and sodium hydroxide. Add 15 g of granular sodium hydroxide to 30 mL of concentrated ammonia solution contained in a conical flask. Immediately fix in the flask a stopper with a straight delivery tube inserted in it. A large quantity of ammonia forms quickly. Simultaneously, the temperature of the reaction increases and froth seethes inside the flask.
7. Prepare ammonia with sodium hydroxide and ammonium salts. Moisten a salt spoon of ammonium chloride, carbonate, or sulfate with drops of caustic soda solution. Ammonia gas forms. It can be tested by smell and by the turning blue of damp red litmus paper.

3.33.01 Prepare ammonia, Haber process
Ammonia is produced industrially by the Haber process with a catalyst, with direct synthesis at high pressure and temperature 45^oC. Cloudy ammonia is clear ammonia solution with soap added in memory of the days before the Haber Process when ammonia was made from coal tar and had cloudy impurities.
N₂ (g) + 3H₂ (g) < = > 2NH₃ (g) + energy released

3.33.1.1 Concentrated hydrochloric acid test (hydrogen chloride test)
Dip one end of a glass rod into concentrated ammonia solution and one end of another glass rod into concentrated hydrochloric acid. Bring the two ends close to each other but do not let them touch. A blue-white smoke of ammonium chloride forms.
NH₃ (g) + HCl (g) --> NH₄Cl (s)

3.33.1.2 Odour test for ammonia
1. Test for ammonia by very cautious smelling. Use very small amounts of reacting chemicals. Do not inhale directly from a test-tube but fan the air above the test-tube towards you.
2. Put a drop of aqueous ammonia solution in a test-tube with a stopper. Shake the test-tube and warm it with the hands. Note the pungent odour of ammonia.
See diagram 1.13: Smelling test
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. Add powdered wood charcoal. Shake the test-tube and smell the contents again. Filter the mixture. Heat the charcoal and note the absence of the pungent odour of ammonia.

3.33.1.3 Litmus tests for ammonia
Ammonia dissolves in water to form a weak base that turns moist red litmus paper blue.

3.33.1.4 Solubility tests for ammonia
1. Ionization reaction, Kb = 1.8 × 10^-5
NH₃ + H₂O <--> NH₄⁺ + OH^-
Dip the open end of a test-tube containing ammonia under water. The test-tube fills with water.
Ammonia is the most soluble of all gases. Ammonia dissolves in water to form ammonia solution, NH₃ (aq). Do not call it "ammonium hydroxide" because while "NH₄⁺" ions and "OH^-" ions can be detected, "NH₄OH" cannot be detected.
2. Show the extreme solubility of ammonia
Remove the stopper from a test-tube of ammonia 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. Describe what you see. The solution made in this rushes up into the test-tube. Ammonia is so soluble that it dissolves almost at once in the water at the mouth of the test-tube. Atmospheric pressure therefore forces the water into the empty test-tube. Ammonia has a greater solubility than hydrogen chloride.

3.33.1.5 Sodium bicarbonate test for ammonia
Add solid sodium bicarbonate and heat until you can smell the ammonia given off.

3.33.1.6 Ammonia fountain test
1. Heat the end of a delivery tube and draw it out to form a fine jet. Fill a flask with ammonia and close the flask with a one-hole stopper with a delivery tube. Add litmus to acidified water in a beaker. Warm the flask gently to expand the gas and then hold the flask upside down with the lower end of the delivery tube in the acidified water. Water soon sprays into the flask through the fine jet as the ammonia dissolves in the water and the pressure of ammonia in the flask decreases. The litmus in the water changes from red to blue.
NH₃ (g) + H₂O (l) < = > NH₃ (aq) + H⁺ + OH^- (aq)
or
NH₃ (g) + H₂O (l) < = > NH₄⁺ (aq) + OH^- (aq)

2. Fill a beaker with litmus solution. Add a few drops of acid solution to the litmus in the cup until the colour just changes to red. Fit a glass jet tube into the stopper of a flask. Remove the stopper and jet, and start filling the dry flask with ammonia. 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 red litmus solution turns blue, because the water in the litmus solution turns part of the ammonia into the alkali ammonium hydroxide.

3. Gently heat in a bath of warm water 100 mL of concentrated ammonia solution in a 250 mL conical flask fitted with a rubber stopper and a short glass tube. Collect the ammonia gas produced in an inverted round bottom flask, by displacement of air. Ammonia is less dense than air. When the round bottom flask is full of ammonia, insert a rubber stopper fitted with a short glass tube tapered inside to a narrow jet. Immerse the external end of the glass tube in water. Ammonia will dissolve in the water in the tube, reducing the pressure inside the flask. Atmospheric pressure will then force water into the tube, creating a jet of water, called the ammonia fountain. Add phenolphthalein to the water to make the reaction more spectacular. Demonstrators and observers must wear safety glasses.

3.33.2 Prepare hydroxides with ammonia solution, double decomposition
Dissolve ammonia gas in water to form ammonia solution, ammonium hydroxide. Prepare dilute solutions of alum (mainly aluminium sulfate), magnesium sulfate, and manganese sulfate. Add ammonia solution to each prepared solution. Note the colours of the insoluble hydroxides formed
Aluminium sulfate + ammonium hydroxide –> aluminium hydroxide (faintly white) + ammonium sulfate
Magnesium sulfate + ammonium hydroxide –> magnesium hydroxide (white) + ammonium sulfate.
Manganese sulfate + ammonium hydroxide –> manganese hydroxide (white brown) + ammonium sulfate.
These double decomposition reactions occur because one of the products is insoluble.

3.33.3 Ammonia solution in a neutralization reaction
Add a finger width of dilute sodium hydrogen sulfate solution in a test-tube, add a few drops of litmus solution. gradually add, with shaking, ammonia solution to the test-tube, using the dropping pipette, until one drop just changes the colour of the mixture to purple blue.
The ammonium hydroxide in the ammonia solution reacts with the sulfuric acid in the sodium hydrogen sulfate solution. As ammonia solution is added, the more acid is destroyed, until a point is reached when there is no more acid and no extra ammonia has been added. The alkali has exactly neutralized the acid (when the drop just changed the colour of the litmus), forming a salt (ammonium sulfate) and water.
Ammonium hydroxide + sulfuric acid --> ammonium sulfate + water.

3.33.4 Ammonia with copper sulfate solution
Add a finger width of ammonia solution to half a test-tube of copper sulfate solution. A double decomposition reaction occurs as you would expect. Note the pale blue solid formed? add more ammonia solution, and shake, until the solid disappears. Try the experiment again, making the copper sulfate solution so dilute that the blue colour is scarcely visible, and adding all the ammonia at once. The pale blue solid is copper hydroxide. When more ammonia solution is added, it reacts with the copper hydroxide, forming a complex copper-ammonia compound which has a deep blue colour. This blue colour appears even with very dilute solutions of copper compounds, and so is a useful test for them.

3.33.5 Ammonia with cobalt chloride solution
To a dilute solution of cobalt chloride add a finger width of ammonia solution. Describe what happens. A blue-green precipitate a complex cobalt-ammonia compound forms.

3.33.6 Ammonium chloride smoke screen
Before starting the experiment, make sure the laboratory is well-ventilated. In one test-tube place a finger width of a mixture of sodium chloride (salt) and sodium hydrogen sulfate, and in another a finger width of a mixture of ammonium chloride and calcium hydroxide. Heat both tubes at the same time, as in the diagram. Hold the mouths of the test-tubes together so that the two colourless gases can combine. Be careful not to inhale the gases and fumes. The gases are hydrogen chloride and ammonia. The gases, hydrogen chloride and ammonia combine, forming the solid salt, ammonium chloride, tiny particles of which form the white smoke.

13.6.5 Tests for ammonia and hydroxyl ions (hydroxide ions)
Ammonia solution is a weak electrolyte. When a strong electrolyte dissolves in water, it almost completely dissociates into ions. Weak electrolytes do not dissociate so much. Water is a very weak electrolyte. The properties of weak electrolytes are affected both by the properties of the molecules in the solution and the properties of the ions in the solution.
1. Note the odour of dilute aqueous ammonia solution.
BE CAREFUL! The odour of ammonia indicates the presence of ammonia molecules in the solution.
2. Test for the presence of hydroxyl ions. Add drops of iron (III) chloride to aqueous ammonia solution. The reaction forms a brown precipitate that indicates the presence of hydroxyl ions in the solution.

13.6.6.1 Catalytic oxidation of ammonia forms nitrogen monoxide, with red-hot platinum wire
See diagram 13.6.6.1: Oxidation of ammonia with platinum wire
BE CAREFUL! Do this experiment in a fume cupboard.
Use concentrated aqueous ammonia solution in a test-tube. Heat a spiral of platinum wire until it becomes red-hot. Insert the wire in the test-tube above the solution. The wire stays red hot and the reaction forms nitrogen monoxide that reacts with oxygen in the air to form nitrogen dioxide.
4NH₃ (g) + 5O₂ (g) --> 4NO (g) + 6H₂O (g)
2NO (g) + O₂ (g) --> 2NO₂ (g)

13.6.6.2 Catalytic oxidation of ammonia, with chromium (III) oxide catalyst
See diagram 13.6.6.2: Oxidation of ammonia with chromium (III) oxide catalyst
BE CAREFUL! Do this experiment in a fume cupboard. Chromium (III) oxide may be carcinogenic
Use chromium (III) oxide as catalyst. Put 0.5 g of ammonium dichromate (VI) in an evaporating dish. Heat with an alcohol lamp until the dichromate starts to decompose. Move the lamp away and the dichromate keeps on decomposing. Wait until the decomposition is completed. Heat the obtained chromium (III) oxide again to dry it thoroughly. To make a catalyst tube, put the freshly prepared chromium (III) oxide in a dry glass tube and squeeze a little glass wool on both sides.
Be careful! Do not touch glass wool with the fingers! Do not breathe in the fibres! Heat the catalyst tube for about 3 minutes to raise the temperature of the catalyst to above 500^oC. By using an air pump, send slowly a stream of air through the concentrated aqueous ammonia solution contained in a conical flask, and then to pass the air ammonia mixed gas over the heated catalyst. When the catalyst becomes red hot, stop heating and continue sending the mixed gas. Prepare the gas coming from the catalyst tube pass through a gas washing bottle of concentrated sulfuric acid to remove the excess ammonia and the water produced in the reaction. A brown-red gas appears in the collecting conical flask. Into this flask pour a little deionized water, shake, then add a few drops of litmus makes the solution turn red to prove that nitric acid forms in this flask.
4NH₃ (g) + 5O₂ (g) --> 4NO (g) + 6H₂O (g)
2NO (g) + O₂ (g) --> 2NO₂ (g)

13.6.7 Reduce copper (II) oxide to copper with ammonia
See diagram 13.6.7: Reduce copper (II) oxide
Pass dry ammonia over copper (II) oxide in a heated hard glass tube. The ammonia reduces the black copper (II) oxide to brown copper and is oxidized to nitrogen gas.
2NH₃ (g) + 3Cu (s) --> 3Cu (s) + 3H₂O (l) + N₂ (g)

3.34 Carbon dioxide
Carbon dioxide gas does not support life so it is a simple asphyxiant. Carbon dioxide and other gases that could accumulate in coal mines to cause choking and suffocation were called choke-damp, after-damp, foul-damp, black damp. Miners used to keep a caged canary bird with them that would die before a concentration of carbon dioxide fatal to humans occurred. Carbon dioxide is used in photosynthesis. Excess carbon dioxide in the atmosphere from excess burning of fossil fuels causes a greenhouse effect so the temperature of the atmosphere rises, called global warming. An increase of the concentration of carbon dioxide in the atmosphere may increase the rate of photosynthesis.

3.34.0 Prepare carbon dioxide with acids and carbonates or bicarbonates, e.g. sodium hydrogen carbonate
See diagram 3.2.38: Collecting carbon dioxide, testing when the receiving jar is full
1. Add dilute hydrochloric acid to carbonates, e.g. calcium carbonate (marble chips) sodium carbonate (washing soda) sodium hydrogen carbonate (baking soda) basic copper (II) carbonate, CuCO₃.Cu(OH)₂.H₂O.
Carbon dioxide is slightly soluble in water so it can be collected over water or by upward displacement of air in dry containers. apply stoppers on the receiving test-tubes to prevent diffusion of the gas into the room.
CaCO₃ (s) + 2HCl (aq) --> CaCl₂ (aq) + H₂O (l) + CO₂ (g)
carbonate + hydrochloric acid --> salt + water + carbon dioxide
2. Add vinegar (acetic acid) or lemon juice (citric acid) to sodium hydrogen carbonate (bicarbonate of soda. The neutralization reaction with these acids forms carbon dioxide.
HC₂H₃O₂ (s) + NaHCO₃ (s) --> NaC₂H₃O₂ (aq) + H₂CO₃ (s)
acetic acid + sodium bicarbonate --> sodium acetate + carbonic acid H₂CO₃ (s) --> H₂O (l) + CO₂ (g)
carbonic acid --> water + carbon dioxide
3. Attach a drawing pin, sharp side up, to the corner of a flat table. Attach a small plastic bag to each end of a wooden ruler. Suspend the centre of the ruler with attached plastic bags over the point of the drawing pin so that the ruler balances horizontally. Add vinegar to powdered sodium hydrogen carbonate in a small beaker. Pout the gas above the mixture into on of the plastic bags. This bag sinks because of the weight of the transferred carbon dioxide gas.
4. Mix vinegar (acetic acid) with sodium hydrogen carbonate in a big container. Drop naphthalene mothballs into the solution. The carbon dioxide formed by the reaction of the vinegar with the sodium hydrogen carbonate forms bubbles of carbon dioxide on the mothballs at the bottom of the container. The mothballs rise to the surface, lose the bubbles and sink again.
2NaHCO₃ (s) --> Na₂CO₃ (s) + CO₂ (g) + H₂O (l)
NaHCO₃ (s) + HC₂H₃O₂ (aq) --> NaC₂H₃O₂ (aq) + CO₂ (g) + H₂O (l)

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

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

or}CO₂ + H₂O --> H₂CO₃

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

3.34.5 Frozen carbon dioxide ("dry ice", "hot ice")
Be careful! When handling dry ice wear eye protection and wear gloves to avoid burns and frost bite. Store dry ice in an expanded polystyrene box. If dry ice is touched, the moisture on the skin freezes and the dry ice sticks to the skin. Never lick dry ice because your tongue will stick to it.
When carbon dioxide is cooled under pressure, it becomes a solid called "dry ice" or "hot ice". Dry ice is used for a refrigerant by mobile ice cream sellers, in fire extinguishers, and for stage effects to produce artificial smoke or mist. At atmospheric pressure, dry ice sublimes at -78^oC. It changes directly from solid to gas. Hold a piece of dry ice in a gloved hand. Watch it disappear as the carbon dioxide sublimes.

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

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

3.34.7 Compare the weight of carbon dioxide and air
Put two identical plastic bags on each pan of a scale or attach them to each arm of a simple beam balance. The plastic bags should be open upwards and must balance perfectly. Prepare carbon dioxide by adding vinegar to sodium bicarbonate in a beaker. Hold the beaker above one of the plastic bags and pour the invisible carbon dioxide into the bag without pouring out any froth or chemicals. The plastic bag containing the carbon dioxide sinks down showing that carbon dioxide is heavier than air. The density of carbon dioxide is about 1.98 g per litre which is about 1.5 × as heavy as air. The fact that you can pour the carbon dioxide into the plastic bag shows that it is heavier than air.

3.35.4 Yeast cells convert glucose to carbon dioxide gas and alcohol
See diagram 3.2.39: Yeast reacting with sugar solution
1. Make a sugar solution and half fill a container with this solution. Add a spoonful of dry yeast and leave to stand for two days. Construct a bubbler to fit on the top of the container. Note whether the yeast forms a gas. Note whether carbon dioxide gas collects in the upper part of the container. Yeast breaks down sugar into ethanol using enzymes that act as catalysts in the conversion:
C₆H₁₂O₆ --> 2C₂H₅OH + 2CO₂ (g)
glucose --> ethanol + carbon dioxide

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

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

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

3.38.1 Test carbon dioxide as a greenhouse gas
The carbon dioxide level decreases during the day through the photosynthesis of green plants, and increases at night when these plant respire and release it. So the blanketing effect on heat movement would increase at night impeding radiation of heat away from the surface of the earth and cause higher surface temperatures. If the concentration of carbon dioxide was stable at all times it would impede as much incoming heat to the surface of the earth as outgoing heat. When a cloud cover forms at the end of a clear day, it blankets the movement of heat from the lower atmosphere to the upper atmospheres so the night temperature does not drop as much as on a clear night. If the cloud cover is there in the day and the night sky is clear, the same clouds become the opposite of a greenhouse gas.
Make a calorimeter from a 4 L bottle. Drill a small hole bottom to install a heat source, e.g. a torch bulb or a heating element. Fix a thermocouple in the neck and link it to a recording device, e.g. a device that can draw a graph of the heat changes over a few minutes. Drill another hole in the side of the bottle to add extra carbon dioxide from a hypodermic syringe. Add only about 1.4 cc. of carbon dioxide to double the concentration from the ambient level to 700 parts per million, then seal the hole with adhesive tape. Insulate the apparatus with styrofoam and locate it out of direct sunlight and away from moving air. Investigate the effects of the nature of the cooling surface, e.g. sand, soil, water, plants, effects of ambient temperature, effects of levels of carbon dioxide and free air.

13.7.6 Prepare carbon dioxide, heat carbonates
Lime burning is the thermal decomposition of calcium carbonate as minerals, e.g. limestone and shells to form calcium oxide (quicklime). Lime burning is an important industry with a long history. Sodium carbonate cannot be decomposed by a burner.
Heat zinc carbonate or basic copper (II) carbonate
CuCO₃.Cu(OH)₂.H₂O --> 2CuO (s) + 2H₂O (l) + CO₂ (g)
ZnCO₃ (s) --> ZnO (s) + CO₂ (g)

13.7.7 Prepare carbon dioxide, heat hydrogen carbonates
Commercial baking powders often contain a solid acid that reacts with the sodium hydrogen carbonate only when moist. Baking powder contains sodium hydrogen carbonate (sodium bicarbonate) that reacts with an acid, e.g. 2-hydroxypropanoic acid (lactic acid) from sour milk, to form carbon dioxide. The heat from the oven helps the decomposition of sodium hydrogen carbonate.
2NaHCO₃ (s) --> Na₂CO₃ (s) + CO₂ (g) + H₂O (l)

13.7.8 Prepare carbon dioxide with a spearmint sweet, e.g. "Mentos", and cola
Order online: Cola-Mentos Fountain Kit
Put a sweet, e.g. a spearmint sweets, "Mentos" to a test-tube. Add aerated water, e.g. cola, "Diet cole". Observe the bubbles of carbon dioxide coming from the surface of the sweet. The sweet provide nucleation sites for the formation of carbon dioxide gas from the carbon dioxide in solution in the cola. Diet cola usually works better than other colas because they usually contain corn syrup that inhibits bubble formation. Be Careful! If you put many Mentos tablets in a cola bottle the resulting explosion may injure people and cause damage.

13.7.9 Prepare carbon dioxide with alum and baking soda
Add alum solution (Al₂(SO₄)₃.K₂(SO₄).24H₂O, potash alum) to baking soda or washing soda. The reaction forms carbon dioxide.

13.7.13 Simulated boiling
Heat about 2 cm depth of sodium hydrogen carbonate in a test-tube. Carbon dioxide gas is given off and the sodium carbonate powder left behaves like a liquid. The cushion of gas between the particles allows them to move independently of each other.

3.42 Hydrogen chloride
Hydrogen chloride gas is corrosive. Do not prepare hydrogen chloride in an open room. Use 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
See diagram 3.2.36: Collecting hydrogen chloride | See diagram 1.13a: Simple fume hood
1. Prepare hydrogen chloride gas by gently warming hydrochloric acid in a water bath in a flask with a gas collection tube. Do this experiment only in a fume cupboard. Collect the gas by displacement of air. Hydrogen chloride can be used in place ammonia in the ammonia fountain in the ammonia fountain experiment.
2. 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, 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 this 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 which turns the blue litmus red.
3. 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.
4. Prepare hydrochloric acid gas from sodium chloride. Mix together 2 g of sodium chloride and 2 g of powdered alum (hydrated potassium aluminium 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. Also, 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.
5. 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!

3.42.01 Prepare hydrochloric acid
See diagram 13.5.2: 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 to form the hydrogen chloride, hold the bottle in your other hand and keep the water swirling to dissolve the 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 tests for hydrogen chloride
1. Remove the stopper from the receiving test-tube under water. Note the solubility of hydrogen chloride.
2. Invert a receiving test-tube over water. The gas dissolves immediately to form hydrochloric acid. The water rises almost to the top because collection by upward displacement of air results in some residual air remaining in the test-tube.
3. 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. Describe what you see. Hydrochloric acid is a very strong add, and although the solution made in this experiment is dilute, treat it with care and wash you hands after the experiment. 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 Moist 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 which turns the blue litmus red.

3.42.1.3 Ammonium chloride test for hydrogen chloride
Hold a piece of cotton wool soaked in ammonia solution at the mouth of a bottle of hydrochloric acid. Note the white cloud of ammonium chloride.

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 solution 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 Hydrogen chloride fountain test
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 the alkaline water. Water soon sprays into the flask through the fine jet as the hydrogen chloride gas dissolves in the water and the pressure of hydrogen chloride in the flask decreases. The litmus in the water changes from blue to red.

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 and makes hydrochloric acid.

3.43 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. 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.

3.43.0 Prepare hydrogen sulfide
See diagram 1.13a: Simple fume hood
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) + 3O_{2 (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, ionization of hydrogen sulfide
Be careful! The gas is soluble in water, so use a solution of hydrogen sulfide in water instead of the gas.
1. Odour test
Hydrogen sulfide has the odour of rotten eggs.
Be careful! Do NOT inhale gases directly from the test-tube. Fan the gas towards the nose with the hand and sniff cautiously. If you detect no odour, move closer and try again.
2. Lead (II) nitrate test
Hydrogen sulfide solution turns lead (II) nitrate solution test paper black.
3. Litmus test
Hydrogen sulfide solution turns blue litmus slightly 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⁺ + S^2-

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

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