School Science Lessons
Chemistry
Updated: 2009-09-15
Please send comments to: J.Elfick@uq.edu.au
3.32 Prepare, collect and test gases, reactions of gases
3.32.0 Prepare gases with a gas generation apparatus
3.32.1 Composition of the atmosphere and greenhouse gases
3.33.0 Ammonia, NH3
3.34.0 Carbon dioxide, CO2
3.39.0 Carbon monoxide, CO
3.40.0 Chlorine, Cl2
3.41.0 Hydrogen, H2
3.42.0 Hydrogen chloride, HCl
3.43.0 Hydrogen sulfide, H2S
3.44.0 Nitrogen monoxide (nitric oxide), NO
3.45.0 Dinitrogen oxide (nitrous oxide), N2O
3.46.0 Nitrogen, N2
3.47.0 Nitrogen dioxide, NO2
3.49.0 Oxygen, O2
3.51.0 Sulfur dioxide, SO2
3.33.0 Ammonia, NH3
3.8.1 Ammonia, anhydrous, hazards
3.33.0 Prepare ammonia, NH3
3.33.1 Tests for ammonia
3.33.2 Prepare hydroxides with ammonia solution by double decomposition
3.33.3 Ammonia solution in a neutralization reaction
3.33.4 Ammonia with copper sulfate solution
3.33.5 Ammonia with cobalt chloride solution
3.33.6 Ammonium chloride smoke screen
3.33.1 Tests for ammonia
3.33.1.1 Concentrated hydrochloric acid test (hydrogen chloride test)
3.33.1.2 Odour test
3.33.1.3 Moist litmus paper test
3.33.1.4 Solubility test
3.33.1.5 Ammonia fountain test
3.34.0 Carbon dioxide
3.8.2 Carbon dioxide, hazards
3.34.0 Prepare carbon dioxide with acids and carbonates or bicarbonates, e.g. sodium hydrogen carbonate
3.34.1 Tests for carbon dioxide
3.34.2 Tests for carbon dioxide in the breath
3.34.3  Solubility of acidic oxide carbon dioxide in water, acidity of soda water, fizzy drinks
3.34.4 Reduce carbon dioxide with burning magnesium
3.34.5 Frozen carbon dioxide ("dry ice", "hot ice")
3.34.5.1 Dry ice in water
3.34.6 Soda-acid fire extinguisher
3.35.0 Carbon dioxide in the home
3.35.4 Yeast cells
3.36 Carbon dioxide and photosynthesis
3.37 Carbon dioxide and respiration
3.38 Carbon dioxide and fermentation for brewing
3.34.1 Tests for carbon dioxide
3.34.1.1 Lighted splint test for carbon dioxide
3.34.1.2 Limewater test for carbon dioxide
3.34.1.3 Burning charcoal test for carbon dioxide
3.34.1.4 Pouring test for carbon dioxide
3.34.1.5 Litmus test for carbon dioxide
3.35.0 Carbon dioxide in the home
3.35.1 Washing soda
3.35.2 Baking soda
3.35.3 Baking powder
3.39.0 Carbon monoxide, CO
3.8.3 Carbon monoxide, hazards
3.39 Carbon monoxide, properties
3.39.1 Reaction of methane with steam
3.40.0 Chlorine
3.8.4 Chlorine, hazards
3.40 Prepare chlorine, Cl2
3.40.1 Lighted splint test for chlorine
3.40.1.1 Bleaching test for chlorine
3.40.2 Pass chlorine through water
3.41.0 Hydrogen gas
3.8.5 Hydrogen, hazards
3.41 Prepare hydrogen gas, H2
3.41.1.0 Tests for hydrogen gas
3.41.2 Prepare hydrogen gas bubbles
3.41.3 Reduce metal oxides to metals with hydrogen gas
3.41.4 Reduce copper oxide with natural gas, methane
3.41.1.0 Tests for hydrogen gas
3.41.1 Lighted splint test for hydrogen gas
3.41.1.1 Litmus test for hydrogen gas
3.41.1.2 Pouring test for hydrogen gas
3.42.0 Hydrogen chloride
3.8.6 Hydrogen chloride, anhydrous, hazards
3.42.0 Prepare hydrogen chloride, HCl
3.42.01 Prepare hydrochloric acid
3.42.1.0 Tests for hydrogen chloride
3.42.1.7 Hydrogen chloride fountain test
3.42.1.0 Tests for hydrogen chloride
3.42.1.1 Solubility test for hydrogen chloride
3.42.1.2 Moist litmus paper 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
3.42.1.6 Ammonia solution test for hydrogen chloride
3.43.0 Hydrogen sulfide
3.8.7 Hydrogen sulfide, hazards
3.43.0 Prepare hydrogen sulfide, H2S
3.43.1 Tests for hydrogen sulfide solution, ionization of hydrogen sulfide
3.43.2 Reduce potassium manganate (VII) with hydrogen sulfide
3.43.3 Reduce iron (III) chloride with hydrogen sulfide
3.44.0 Nitrogen monoxide (nitric oxide) NO
3.44 Prepare nitrogen monoxide (nitric oxide) NO
3.44.1 Catalytic conversion of nitrogen monoxide (nitric oxide)
3.45.0 Dinitrogen oxide (nitrous oxide, N2O, nitrogen (I) oxide, dinitrogen monoxide)
3.45 Prepare dinitrogen oxide (nitrous oxide) N2O
3.45.1 Tests for dinitrogen oxide (nitrous oxide)
19.4.4.22 Packaging gases, propellants, food additives
3.47.0 Nitrogen dioxide
 3.47 Prepare nitrogen dioxide, NO2
3.47.1 Pass nitrogen dioxide through water
3.48 Acid rain and nitrogen oxides, NOx
3.49.0 Oxygen
 3.49 Prepare oxygen, O2
3.49.1 Tests for oxygen
3.51.0 Sulfur dioxide
 3.51 Prepare sulfur dioxide, SO2
3.51.1 Tests for sulfur dioxide
3.51.2 Reduce potassium manganate (VII) with sulfur dioxide
3.51.3 Reduce iron (III) chloride with sulfur dioxide
3.51.4 Bleach flowers with sulfur dioxide
3.32.0 Prepare gases with a gas generation apparatus
See diagram 3.32: Gas generation apparatus | See: Saturation vapour pressure over water
Collect less dense gas by downward displacement of air, see diagram 1.
Collect more dense gas by upward displacement of air, see diagram 2.
Collect insoluble gas over water, see diagram 3.
Use a borosilicate test-tube that is not cracked. Clamp the test-tube to a stand. Put the solid reagent in the sidearm test-tube and the liquid reagent in the reservoir. Add the liquid reagent very slowly drop by drop. Keep the reservoir tap closed and the reservoir full to prevent gases blowing back. Grease the stopper and insert it so that if an accidental sudden increase in pressure occurs, the stopper blows out of the test-tube. Use rubber tubing to collect the sidearm to a delivery tube that leads into the receiving test-tube. Discard the first gas coming out of the delivery tube because it is mostly air. Never allow a flame near the gas as it comes out of the delivery tube. Some air probably remains in the receiving test-tube. Use the gas bubbler to collect over water insoluble gases with similar density to air. Some water vapour remains in the receiving test-tube. Gases can also be collected in balloons, inflatable footballs, and plastic bags.
3.32.1 Composition of the atmosphere and greenhouse gases
Gas and percentage volume in dry air: N2 78.08%, O2 20.95%, Ar 0.93%, CO2 0.03%, Ne 0.0018%, He 0.00052%, Kr 0.00011%, Xe 0.000009%, Rn 6 X 10-18%. The average formula weight of air is 28.8. The apparent molar mass is 28.96 g / mol. The main greenhouse gases caused by human activities are as follows:
1. Carbon dioxide from burning of fossil fuels, wood and chemical reactions. However, plants remove carbon dioxide from the atmosphere, sequester, during photosynthesis.
2. Methane, CH4, from coal, natural gas, oil, digestion by herbivores and anaerobic decay of plants in rice paddy and solid waste landfills.
3. Nitrous oxide, N2O from combustion of fossil fuels and solid wastes and from chemical reactions and agricultural activities.
4. Fluorinated gases, i.e. hydrofluorocarbons, e.g. tetrafluoroethane (CH2FCF3, R-134a) perfluorocarbons, e.g. tetrafluoromethane (CF4, carbon tetrafluoride, R14) and sulfur hexafluoride (SF6) from chemical reactions. They have high global warming potential but they are not ozone-depleting as are CFCs, e.g. dichlorodifluoromethane (CCl2F2, R-12, "Freon-12") HCFCs, e.g. difluoromonochloromethane (CHClF2, "Freon 22") and halons, e.g. bromochlorodifluoromethane (CF2ClBr, "Halon 1211").
3.33.0 Prepare ammonia
See diagram 3.33.1: Prepare ammonia | See diagram 3.33.2: A fountain experiment
Ammonia is an extremely irritating gas and is flammable in the presence of sufficient oxygen. Do not prepare ammonia in an open room. Use a fume cupboard.
Ammonia is less dense than air. Ammonia solution is a weak electrolyte so the properties of the molecules and the ions in the solution affect its properties. Ammonia (NH3) is produced industrially by the Haber process with a catalyst, with direct synthesis at high pressure and temperature 45oC. Cloudy ammonia is clear ammonia solution with soap added in memory of the days before the Haber Process when ammonia was made from coal tar and had cloudy impurities.
N2 (g) + 3H2 (g) < = > 2NH3 (g) + energy released
1. Put a mixture of calcium hydroxide and ammonium chloride into a test-tube to a depth of 4 cm. Fill a U-tube with lumps of calcium oxide mixed with cotton wool. The cotton wool is to prevent blocking of the U-tube. Gently heat the test-tube. The calcium oxide is to dry the ammonia gas. Collect the gas by downward displacement of air. Test whether the receiver test-tube is full by holding a piece of moist red litmus paper at the opening. Ammonia gas turns red litmus blue. Collect test-tubes of ammonia gas and apply stoppers.
2NH4Cl (aq) + Ca (OH)2 (s) --> 2NH4OH (s) + CaCl2 (aq)
then NH4OH (s) --> NH3 (g) + H2O (l)
2. 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.
3. 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.
4. Prepare ammonia with ammonium chloride and sodium carbonate. Put 5 g of ammonium chloride (sal ammoniac) in 2 cm depth sodium carbonate (washing soda) solution. Heat the test-tube. Tests for ammonia gas and with wet red litmus paper.
5. Prepare ammonia with ammonia solution and sodium hydroxide. Add 15 g of granular sodium hydroxide to 30 mL of concentrated ammonia solution contained in a conical flask. Immediately fix in the flask a stopper with a straight delivery tube inserted in it. A large quantity of ammonia forms quickly. Simultaneously, the temperature of the reaction increases and froth seethes inside the flask.
3.33.1.1 Concentrated hydrochloric acid test (hydrogen chloride test)
Dip one end of a glass rod into concentrated ammonia solution and one end of another glass rod into concentrated hydrochloric acid. Bring the two ends close to each other but do not let them touch. Blue-white smoke of ammonium chloride forms.
NH3 (g) + HCl (g) --> NH4Cl (s)
3.33.1.2 Odour test
Tests for ammonia by very cautious smelling. Use very small amounts of reacting chemicals. Do not inhale directly from a test-tube but fan the air above the test-tube towards you.
3.33.1.3 Moist litmus paper test
Ammonia dissolves in water to form a weak base that turns moist red litmus paper blue.
3.33.1.4 Solubility test
4.1 Ionization reaction, Kb = 1.8 X 10-5
NH3+ H2O <--> NH4+ + OH-
Dip the open end of a test-tube containing ammonia under water. The test-tube fills with water.
Ammonia is the most soluble of all gases. Ammonia dissolves in water to form ammonia solution, NH3 (aq). Do not call it "ammonium hydroxide" because while "NH4+" ions and "OH-" ions can be detected, "NH4OH" cannot be detected.
4.2  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 Ammonia fountain test
5.1 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)
5.2. 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.
3.33.2 Prepare hydroxides with ammonia solution by 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 Ammonia solution 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 you start, make sure your 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.
3.34.0 Prepare carbon dioxide with acids and carbonates or bicarbonates, e.g. sodium hydrogen carbonate
See diagram 3.34: Collecting carbon dioxide, testing when the receiving jar is full
Carbon dioxide gas does not support life so it is a simple asphyxiant. Carbon dioxide and other gases that could accumulate in coal mines to cause choking and suffocation were called choke-damp, after-damp, foul-damp, black damp. Miners used to keep a caged canary bird with them that would die before a concentration of carbon dioxide fatal to humans occurred.
Carbon dioxide is used in photosynthesis. Excess carbon dioxide in the atmosphere from excess burning of fossil fuels causes a greenhouse effect so the temperature of the atmosphere rises, called global warming. An increase of the concentration of carbon dioxide in the atmosphere may increase the rate of photosynthesis.
1. Add dilute hydrochloric acid to 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. Carbonic acid soon decomposes, but it can form stable sodium carbonate, potassium carbonate and hydrogen carbonate salts.
1. Open a bottle of soda water or fizzy lemonade. Bubbles of carbon dioxide appear as the gas leaves the solution under the lower atmospheric pressure. Carbon dioxide leaves the solution. Tests for carbon dioxide by putting a lighted splint in the bottle above the lemonade. Test the pH of soda water at room 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) carbonic acid (equilibrium 2)
H2CO3 (aq) + OH- (aq) <--> H2O (l) + HCO3 (aq)- (hydrogen carbonate ion, hydrogencarbonate ion) (equilibrium 3)
or
H2CO3 (aq) <--> H+ (aq) + HCO3- (aq) (hydrogen carbonate ion, hydrogencarbonate ion) (equilibrium 3)
HCO3- (aq) + OH- (aq) <--> H2O (l) + CO32- (aq) (carbonate ion) (equilibrium 4)
or
HCO3- (aq) <--> H+ (aq) + CO32- (aq) (carbonate ion) (equilibrium 4)
or
CO2 + H2O <--> H3O+ + HCO3-
HCO3- + H2O <--> H3O+ + CO32-
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 -78oC. It changes directly from solid to gas. Hold a piece of dry ice in a gloved hand. Watch it disappear as the carbon dioxide sublimes.
3.34.5.1 Dry ice in water
Fill a 10 cc measuring cylinder water and add universal indicator. Add drops of sodium hydroxide solution. Add a lump dry ice. Note how it sinks to the bottom and gives off bubbles of carbon dioxide to make a fog at the mouth of the measuring cylinder. The universal indicator slowly changes colour from blue, pH 9, to orange, pH 5, as the pH reaches about 4.5.
OH- (aq) + CO2 (g) –> HCO3- (aq)
Repeat the experiment with ammonia solution. The colour change of the universal indicator is more gradual because of the reaction of weak acids with weak bases.
H2O (l) + NH3 (aq) + CO2 (g) –> NH4+ (aq) + HCO3- (aq)
3.34.6 Soda-acid fire extinguisher
Use a plastic drink bottle with a one-hole rubber stopper fitted with a plastic tube. Connect rubber tubing with a nozzle to the tube. Use a test-tube that can fit inside the bottle. Partly fill the bottle with sodium hydrogen carbonate solution. Fill the test-tube with dilute sulfuric acid solution and lower it gently into the bottle so that it rests upright. Fit the stopper and plastic tube. Add a detergent to the acid to produce the blanketing effect of foam. Aim the bottle at the fire and invert the bottle rapidly. A strong reaction forms carbon dioxide. The pressure of the gas pushes the liquid out through the jet to extinguish the fire.
2NaHCO3 (aq) + H2SO4 (l) --> Na2SO4 (s) + H2O (l) + CO2 (g)
To make a foam similar to the foam blanket produced by fire extinguishers, add sodium hydrogen carbonate to warm soapy water in a beaker. Add concentrated aluminium sulfate solution and note the mass of white bubbles that looks like ice-cream soda.
3.35.4 Yeast cells
Yeast cells convert sugar to carbon dioxide gas and alcohol to make bread rise.
See diagram 3.35: Yeast reacting with sugar solution
1. Make a sugar solution and half fill a container with this solution. Add a spoonful of dry yeast and leave to stand for two days. Construct a bubbler to fit on the top of the container. Note whether the yeast forms a gas. Note whether carbon dioxide gas collects in the upper part of the container. Yeast breaks down sugar into ethanol using enzymes that act as catalysts in the conversion:
C6H12O6 --> 2C2H5OH + 2CO2 (g)
glucose --> ethanol + carbon dioxide
3.36 Carbon dioxide and photosynthesis
nCO2 + nH2A --> (CH2O)n + nO2
carbon dioxide + hydrogen donor --> carbohydrate + oxygen 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
3.39 Carbon monoxide, properties
Be careful! Do NOT make carbon monoxide.
See 18.6.3: Danger of vehicle exhausts, tailpipe gases
Air pollution
Carbon monoxide is very toxic. It can cause unconsciousness because of combination of the gas with haemoglobin in the blood. Death can occur from carbon monoxide inhalation. Do not prepare carbon monoxide. Metal oxides are reduced by passing carbon monoxide over the heated oxide. Carbon monoxide is very poisonous and particularly dangerous because it is colourless and has no smell. It kills more people than any other gas. Carbon monoxide is poisonous because it reacts with the haemoglobin in blood and prevents the blood from acting as an oxygen carrier. The gas can form accidentally by leaving a car engine running in a closed garage or by burning a gas fire with restricted ventilation. When carbon or carbon compounds burn in a limited supply of air, the reaction forms carbon monoxide.
2C (s) + O2 (g) --> 2CO (g)
carbon + oxygen gas --> carbon monoxide
Carbon monoxide is insoluble in water, but it is absorbed by potassium hydroxide solution. Carbon monoxide burns with a pale blue flame forming carbon dioxide.
2CO (g) + O2 (g) --> 2CO2 (g)
Carbon monoxide can act as a reducing agent and is the main reducing agent in a blast furnace. At high temperatures, carbon monoxide reduces the oxides of copper, lead and iron to the metal.
CuO (s) + CO (g) --> Cu (s) + CO2 (g)
Fe2O3 (s) + 3CO (g) --> 2Fe (s) + 3CO2 (g)
3.39.1 Reaction of methane with steam, at 700oC and nickel catalyst forms hydrogen and carbon monoxide.
CH4 (g) + H2O (g) --> 3H2 (g) + CO (g)
3.40 Prepare chlorine, Cl2
See diagram 1.13a: Simple fume hood
Chlorine gas is very toxic. Can react to cause fires or explosions upon contact with turpentine, ether, ammonia gas, illuminating gas, hydrocarbon, hydrogen and powdered metals. Dissolves readily in water forming highly corrosive solution. Do not prepare chlorine in open room. Use fume cupboard. Direct combination of chlorine and hydrogen in bright light or ignition of the mixture by lighted taper or electric spark. Reactions of chlorine with metals, solid non-metals, hydrocarbon. Use small quantities only.
Fume cupboards, fume chambers, fume hoods,
Chlorine is a greenish yellow gas with an irritating and choking odour. Be careful! Chlorine gas is poisonous and damages the respiratory organs. Do not inhale gases directly from the test-tube. Fan the gas towards the nose with the hand and sniff cautiously. If no odour is detected, move closer and try again. Prepare chlorine with bleaching powder, bleach solution. Bleaching powder is a mixture of calcium chloride, calcium hydroxide and calcium chlorate (I). Bleaching powder is made commercially by the reaction of chlorine with solid calcium hydroxide. Do the following experiments in a fume cupboard, fume hood or near an open window. Before doing these experiments, make available sodium thiosulfate or calcium hydroxide solution to be used for a chlorine trap to absorb excess chlorine gas. Also prepare ammonia solution because the effect of inhaling chlorine gas may be counteracted by inhaling ammonia vapour. The best treatment for inhaling chlorine gas is plenty of fresh air.
1. With great care, warm bleaching powder and smell it until you notice a choking smell because of chlorine gas being produced by the action of carbon dioxide in the air. Test with wet red or blue litmus paper that becomes colourless because of the bleaching action of chlorine.
2. Put 5 g of bleaching powder (calcium hypochlorite) into a test-tube. Add drops of a weak acid, e.g. citric acid or vinegar. Test with wet red or blue litmus paper. Hold a piece of white paper behind the apparatus to note the green chlorine gas.
3. Add dilute sulfuric acid to bleaching powder. After collecting a small amount of chlorine gas put a stopper in the receiving test-tube and put the end of the delivery tube into sodium thiosulfate solution to absorb excess chlorine.
Bleaching powder + H2SO4 (aq) --> CaSO4 (s) + H2O (l) + Cl2 (g)
4. Domestic bleach is manufactured by mixing a solution of chlorine with sodium hydroxide solution
Cl2 (g) + 2OH- (aq) --> Cl- (aq) + ClO- (aq) + H2O
Add a dilute acid to bleach solution to form chlorine gas.
NaOCl (aq) + HCl (aq) --> NaCl (aq) + H2O (l) + Cl2 (g)
3.40.1 Lighted splint test for chlorine
Chlorine extinguishes a lighted splint, but hot steel wool burns in it
3.40.1.1 Bleaching test for chlorine
Chlorine bleaches moist red or blue litmus paper, flowers and some dyes in cloth.
3.40.2 Pass chlorine through water
Chlorine is available commercially for school laboratory use as chlorine water. Hypochlorous acid HClO, a bleach and a disinfectant, is a solution of chlorine (I) oxide that forms salts called hypochlorites. Hypochlorous acid is a weak acid that easily decomposes back to chlorine gas and water. When chlorine passes through water, a mixture of HCl and HClO forms. The chlorine is oxidized and reduced.
Cl2 (g) + H2O (l) < = > HCl (aq) + HClO (aq)
3.41 Prepare hydrogen gas
See diagram 3.41: Collecting hydrogen gas
Do not allow direct combination of hydrogen and chlorine in bright light or ignition of the mixture by lighted taper or electric spark. You can ignite a jet of hydrogen issuing from a delivery tube. Hydrogen reduces metal oxides.
Hydrogen, H2, is a colourless odourless diatomic gas with the lowest density of any element. Hydrogen does not change the colour of moist litmus. The hydrogen ion, H+, is a proton.
1. Zinc with hydrochloric acid
Do not use a container bigger than a test-tube. Put granulated zinc in a test-tube and cover it with water. Add a crystal of copper (II) sulfate to act as a catalyst. Slowly add dilute hydrochloric acid through a funnel, as in diagram 2.41.1 or through a syringe, as in 2.41.2. Bubbles of hydrogen appear on the surface of the zinc. The test-tube feels hot because the reaction is exothermic. Collect hydrogen gas by downward displacement or over water. Let the reaction continue for some minutes to drive out all the air from the test-tube. Discard the first two test-tubes of hydrogen because they will contain displaced air. Collect test-tubes of the gas and apply stoppers.
Zn (s) + 2HCl (aq) --> ZnCl2 (s) + H2 (g)
2. Iron with sulfuric acid or citric acid, or sodium hydrogen sulfate
Put 1 cm depth of iron filings in a test-tube. Just cover the iron filings with a dilute acid solution. Warm the test-tube until frothing starts. Hydrogen is colourless and odourless but any impurities in the iron filings give a nasty smell.
3. Aluminium with sodium carbonate
Cut into small pieces some aluminium foil or aluminium milk bottle tops and put into a test-tube. Add 5 mL of sodium carbonate solution (Na2CO3.10H2O, washing soda). Heat until effervescence occurs.
4. Iron with alum
Put 5 g of iron filings in a 1 cm depth of alum solution in a test-tube. Heat the solution until effervescence occurs. [Potash alum, "alum" has the formula Al2(SO4)3.K2(SO4).24H2O and is also shown as KAl(SO4)2.12H2O.]
5. Iron with ammonium chloride
Put an equal volumes mixture of iron filings and ammonium chloride in a dry test-tube and heat. Hydrogen gas and ammonia are given off.
6. Calcium with hydrochloric acid
Use forceps to transfer about 0.1 g of calcium metal turnings to dilute hydrochloric acid in a test-tube.
Ca (s) + 2HCl (aq) --> CaCl2 (aq) + H2 (g)
3.41.1 Lighted splint test for hydrogen gas
Be careful! A dangerous explosion may occur if you use anything bigger than a small test-tube when igniting the gas, particularly if the gas is mixed with air. Never test more than a test-tube full of hydrogen gas. Never dry hydrogen gas with concentrated sulfuric acid.
1. Hold a lighted splint or burning taper to the mouth of a test-tube. The gas explodes with a squeaky pop sound. The splint is extinguished. The squeaky pop shows rapid combustion of hydrogen to form water vapour. Look for vapour on the sides of the test-tube. However, as 2 litres of gas forms only about 1 mL of liquid, the liquid on the sides of the test-tube may just show that test-tube was already wet before the experiment.
2H2 (g) + O2 (g) --> 2H2O (l)
hydrogen gas + oxygen gas --> water
3.41.1.1 Litmus test for hydrogen gas
Hydrogen does not change the colour of moist litmus.
3.41.1.2 Pouring test for hydrogen gas
Test whether hydrogen is lighter than air by "pouring" the gas into a test-tube held either above the first test-tube or below it. Use a lighted taper to investigate where the hydrogen has gone.
3.41.2 Prepare hydrogen gas bubbles
Hydrogen is much lighter than air and was formerly used in airships, dirigible balloons. It has now been replaced by helium because hydrogen ignites easily.
Pass hydrogen through soapy water to form soap bubbles full of hydrogen. Shake the bubbles gently to make them float up. The hydrogen bubbles rise into the air, showing the low relative density of hydrogen gas. Try to ignite the bubbles with a lighted splint.
3.41.3 Reduce metal oxides to metals with hydrogen gas
See diagram 3.41.3: Hydrogen over heated copper oxide
Be careful! Use a safety screen and wear eye protection
1. Pass hydrogen over 5 g of copper (II) oxide (CuO, black copper oxide) or lead (II) oxide (lead monoxide, PbO, lithage) or iron (III) oxide (haematite, Fe2O3). Hydrogen reduces metal oxides to metals. The products are the metal and water.
Weigh a reduction tube empty then with copper oxide. Pass hydrogen over the copper oxide and light the gas as it comes out of the hole in the end of the combustion tube. Heat the copper oxide with a Bunsen burner flame until it glows then turns pink. The glow shows that reduction occurs. Remove the Bunsen burner. Let the combustion tube cool then discontinue the supply of hydrogen. When the flame has gone out remove the stopper and weigh the reduction tube and contents again.
CuO (s) + H2 (g) --> Cu (s) + H2O (l)
In the industrial process, blistered copper is heated in a furnace and natural gas is passed through the molten copper oxide until the flame burns green to indicate that almost pure copper remains.
2. Repeat the experiment with 5 g of copper (I) oxide (red copper oxide, Cu2O)
3.41.4 Reduce copper oxide with natural gas, methane
1. Pass natural gas, about 95% methane, over heated copper (II) oxide powder. The reduction reaction is slow and may need twenty minutes of strong heating. The copper does not glow with heating so it is not clear when all the copper oxide has been reduced.
4CuO (s) + CH4 (g) --> 4Cu (s) + 2H2O (l) + CO2 (g)
2. See: Metaldehyde
Repeat the experiment with a 1 cm cubic piece of metaldehyde in the reduction tube. The reduction is quicker.
3. Repeat the experiment with natural gas that has bubbled through ethanol. The reduction is quicker and a slight glow is seen as the copper oxide is reduced.
6CuO (s) + C2H5OH (l) --> 6Cu (s) + 3H2O (l) + 2CO2 (g)
3.42.0 Prepare hydrogen chloride
See diagram 3.42: Collecting hydrogen chloride | See diagram 1.13a: Simple fume hood
Hydrogen chloride gas is corrosive. Do not prepare hydrogen chloride in an open room. Use fume cupboard.
Be careful! Do 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.
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)
2. This 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.
3.42.01 Prepare hydrochloric acid
See diagram 13.5.2: Prepare hydrochloric acid
Repeat the above experiment, but 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 test 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.
3.43.0 Prepare hydrogen sulfide
See diagram 1.13a: Simple fume hood
Hydrogen sulfide gas is both an irritant and an asphyxiant. Do not prepare hydrogen sulfide in an open room. Use fume cupboard. You can ignite a jet of hydrogen sulfide issuing from a delivery tube.
Be careful! Hydrogen sulfide is an extremely poisonous colourless flammable gas with an unpleasant smell of rotten eggs. At less than 1% concentration the smell disappears. So a student may be breathing in this poisonous gas without being aware of it. Do NOT use a Kipp's apparatus for generating hydrogen sulfide.
1. Do this experiment in a fume cupboard, fume hood. Put 5 sodium thiosulfate crystals in a metal screw cap. Heat the metal screw cap gently by holding it with pincers in a Bunsen burner flame until the crystals have melted and solidified again, with steam given off.
Be careful! Do NOT inhale gas directly from the metal screw cap.
With more careful heating, note the "rotten egg" smell of hydrogen sulfide. Allow the metal screw cap to cool. Moisten the white residue with a weak acid, e.g. vinegar. The smell of hydrogen sulfide gas becomes stronger. Dip a strip of clean newspaper in the copper (II) sulfate solution and hold it over the meal screw cap. The paper turns black.
2. Do this experiment in a fume cupboard, fume hood. Add dilute hydrochloric acid to iron sulfide. Collect the gas over warm water by downward displacement.
FeS (s) + 2HCl (aq) --> FeCl2 (aq) + H2S (g)
Ignite the gas as it leaves the delivery tube.
2H2S (g) + 3O2 (g) --> 2SO2 (g) + 2H2O (l)
3.43.1 Tests for hydrogen sulfide solution, ionization of hydrogen sulfide
Be careful! The gas is soluble in water, so use a solution of hydrogen sulfide in water instead of the gas.
1. Odour test
Hydrogen sulfide has the odour of rotten eggs.
Be careful! Do NOT inhale gases directly from the test-tube. Fan the gas towards the nose with the hand and sniff cautiously. If you detect no odour, move closer and try again.
2. Lead (II) nitrate test
Hydrogen sulfide solution turns lead (II) nitrate solution test paper black.
3. Litmus test
Hydrogen sulfide solution turns blue litmus slightly red-pink.
4. Copper (II) sulfate test
Hydrogen sulfide solution turns copper (II) sulfate solution black.
Ionization of hydrogen sulfide
H2S + H2O --> H3O+ + HS-
HS- + H2O --> H3O+ + S2-
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 green-brown precipitate of iron (II) hydroxide..