Topic 13 Gases
Updated 2008-09-04 R
Please send comments to: J.Elfick@uq.edu.au
See also: Interesting websites
Table of contents
3.32 Prepare gases with a gas generation apparatus
3.32.1 Composition of the atmosphere
3.8.0 Hazards associated with gases
13.1.5 Gases, relative molecular mass of gases
13.1.6 Molar volume of oxygen prepared with hydrogen peroxide
13.0.0 Prepare gases
13.01 Gas bags
13.1 Ammonia, anhydrous (NH3)
13.2 Bromine, Br2
13.3 Butane, C4H10
13.4 Carbon dioxide, CO2
3.39 Carbon monoxide, CO
13.6 Chlorine, Cl2
13.7 Ethane, Ethene, Ethyne
13.8 Fluorine, F
13.9 Hexane, C6H14, Heptane, C7H16
13.10 Hydrogen, H2
13.11 Hydrogen bromide, HBr
13.12 Hydrogen chloride, HCl
13.13 Hydrogen sulfide, H2S
13.14 Methane, CH4
13.15 Nitrogen, N2
13.16 Nitrogen dioxide, NO2
13.17 Nitric oxide, NO
13.18 Nitrous oxide, N2O
13.19 Octane, C8H18
13.20 Oxygen, O2
13.21 Ozone, O3
13.22 Pentane, C5H12
13.23 Petroleum gas
13.24 Propane, C3H8
13.25 Sulfur dioxide, SO2, sulfur (IV) oxide
13.26 Trichloromethane, CCl3
13.27 Triodomethane, CHI3
13.0.0 Prepare Gases
13.01 Gas bags
See diagram 13.01: Gas bag, cable tie
1. Party balloons casn be inflated with gas only from a high pressure source, e.g. a gas cylinder.
2. Snap lock resealable polythene bags. can be resealed with a finger press sealing strip to give a gas-tight seal. The closure system which reseals an opened bag includes a pressure sensitive adhesive on the front side and a defined release surface on the back of the bag. The top portion of the bag is folded so that the defined release surface comes into contact with the adhesive to reseal the opened bag.
3. The plastic bag used in a 2 litre wine cask can be washed out and used to stire gas. Use a cork borer to insert a glass tube through a one hole rubber stopper. Be Careful! Leave 1 cm of glass tube to protrude from the top of the stopper. Pull tight around the neck of the plastic bag around the rubber stopper and securely it tightly with a cable tie. To check for leaks, close the end of the glass tube with a rubber cap, immerse the bag in water and squeeze the bag. To fill the bag, squeeze it flat then fill it from a gas cylinder or chemical generator. Refill the bag and squeeze out the gas more than once to ensure that any air is flushed out. When the bag is finally filled, close the glass tube with a rubber cap. To get a gas sample, inject the hypodermic needle of a syringe through the rubber cap and suck gas into the syringe.
A cable tie usually consists of a Nylon tape with a gear rack and a ratchet within a small open case. When the pointed tip of the cable tie has been pulled through the case and past the ratchet, it cannot be pulled back, but the loop formed may be pulled tighter. Cable ties are used to bind several cables together, e.g. around a motor car engine.
13.1 Ammonia, anhydrous (NH3)
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.
3.33 Prepare ammonia, NH3
3.33.1 Tests for ammonia, ammonia fountain experiment
12.11.3 Ammonia and the ammonium ion, NH3, NH4+
13.6.5 Tests for ammonia and hydroxyl ions (hydroxide ions)
13.6.6 Catalytic oxidation of ammonia with red-hot platinum wire
13.6.6.1 Catalytic oxidation of ammonia with chromium (III) oxide catalyst
13.6.7 Reduce copper (II) oxide to copper with ammonia
13.2 Bromine, Br2
12.19.9.1 Bromine, Br2: Reactions of bromine
13.3 Butane, C4H10
16.1.7 Butane (C4H10): Prepare butane
13.4 Carbon dioxide, CO2
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.
3.34 Prepare carbon dioxide, CO2
3.34.1 Tests for carbon dioxide
3.34.2 Test the breath for carbon dioxide
3.34.3 Solubility of carbon dioxide in water
3.34.4 Reduce carbon dioxide with burning magnesium
3.34.5 Frozen carbon dioxide ("dry ice", "hot ice")
3.34.5.1 Dry ice in water
3.34.6 Soda-acid fire extinguisher
3.35 Carbon dioxide in the home
3.36 Carbon dioxide and photosynthesis
3.37 Carbon dioxide and respiration
3.38 Carbon dioxide and fermentation for brewing
13.7.6 Prepare carbon dioxide by heating carbonates
13.7.7 Prepare carbon dioxide by heating hydrogen carbonates
13.7.13 Simulated boiling
13.6 Chlorine, Cl2
Chlorine This 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.
3.40 Prepare chlorine, Cl2
3.40.1 Tests for chlorine
3.40.2 Pass chlorine through water
13.4.4 Prepare chlorine with dilute hydrochloric acid and domestic bleach solution
13.4.5 Prepare chlorine with concentrated hydrochloric acid and manganese (IV) oxide
13.4.6 Prepare chlorine with concentrated acid and potassium manganate (VIII)
12.19.8.2 Prepare chlorine with concentrated hydrochloric acid and potassium manganate (VIII)
13.4.7 Reactions of chlorine with sodium
13.4.8 Burn steel wool in chlorine
13.4.9 Burn copper in chlorine
13.4.10 Burn a wax taper in chlorine, reaction of chlorine with non-metals
13.4.11 Pass chlorine through water
13.4.12 Pass chlorine through iodine solution
13.4.13 Pass chlorine through iron (II) chloride solution
13.4.14 Reactions of chlorine with heated copper and steel wool
13.4.15 Reactions of chlorine with alkalis, bleaching powder
12.19.5.0 CFCs, chlorofluorocarbons
18.7.2.1 Swimming pool chemistry, test for chlorine
13.7 Ethane, Ethene, Ethyne
16.1.5 Ethane: Prepare ethane, C2H6
16.1.1.2.1 Ethene: Prepare ethene (ethylene) gas, C2H4
16.1.1.3 Ethyne (acetylene), C2H2
13.8 Fluorine, F
12.19.7.0 Fluorine, F
13.9 Hexane, C6H14, Heptane, C7H16
16.1.9 Hexane (C6H14)
16.1.10 Heptane (C7H16)
13.10 Hydrogen, H2
Hydrogen, H2
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.
3.41 Prepare hydrogen, H2
3.41.1 Tests for hydrogen
3.41.2 Prepare hydrogen bubbles
3.41.3 Reduce metal oxides to metals with hydrogen
13.2.3.0 Prepare hydrogen with iron filings and sulfuric acid, or sodium hydrogen sulfate
13.2.3.1 Prepare hydrogen with aluminium foil and sodium carbonate, washing soda
13.2.3.2 Prepare hydrogen with iron filings and alum
13.2.5 Reduce metal oxides to metals with hydrogen
13.11 Hydrogen bromide, HBr
12.19.9.3 Hydrogen bromide: Prepare hydrogen bromide, HBr
13.12 Hydrogen chloride, HCl
Hydrogen chloride gas is corrosive. Do not prepare hydrogen chloride in an open room. Use fume cupboard.
3.42 Prepare hydrogen chloride, HCl
3.42.1 Tests for hydrogen chloride
12.19.7.1 Hydrogen fluoride: Prepare hydrogen fluoride, HF
12.19.6.1 Hydrogen iodide: Prepare hydrogen iodide, HI
13.13 Hydrogen sulfide, H2S
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.
3.43 Prepare hydrogen sulfide, H2S
3.43.1 Tests for hydrogen sulfide solution
3.43.2 Reduce potassium manganate (VII) with hydrogen sulfide
3.43.3 Reduce iron (III) chloride with hydrogen sulfide
13.13.8 Dry hydrogen sulfide and dry sulfur dioxide will not react
13.14 Methane (CH4)
16.1.3 Methane, natural gas (CH4)
16.1.4 Methane (CH4): Prepare methane gas
13.15 Nitrogen, N2
Atmospheric nitrogen cannot be used directly by the body. Liquid nitrogen is used to freeze tissues for microscopic examination.
3.46 Prepare nitrogen, N2
13.9.0 | Nitrogen, N2
13.9.3 Nitrogen gas generated in a motor car airbag
13.16 Nitrogen dioxide, NO2
3.47 Prepare nitrogen dioxide, NO2, nitrogen (IV) oxide, dinitrogen tetroxide, Also: nitrogen tetroxide, N2O4
3.47.1 Pass nitrogen dioxide through water, NO2
3.48 Acid rain and nitrogen oxides, NOx
13.10.2 Prepare nitrogen dioxide from lead (II) nitrate crystals
13.10.2 Prepare nitrogen dioxide (C8H18)from lead (II) nitrate crystals, NO2
13.17 Nitric oxide, NO, nitrogen monoxide
Nitric oxide has many physiological functions in the body, e.g. dilation of blood vessels.
3.44 Prepare nitrogen monoxide (nitric oxide) NO
3.44.1 Catalytic conversion of nitrogen monoxide (nitric oxide)
13.18 Nitrous oxide, N2O
3.45 Prepare dinitrogen oxide (nitrous oxide) N2O
3.45.1 Tests for dinitrogen oxide (nitrous oxide)
13.19 Octane, C8H18
16.1.11 Octane C8H18
16.1.11.1 Octane number
13.20 Oxygen, O2
3.49 Prepare oxygen, O2
3.49.1 Tests for oxygen
13.1.6 Molar volume of oxygen prepared with hydrogen peroxide
13.3.1 Prepare oxygen foam with hydrogen peroxide
13.3.2 Burn sulfur in oxygen
13.3.3 Burn steel wool in oxygen, burn iron filings
13.3.4 Burn magnesium ribbon in oxygen
13.3.5 Prepare an oxygen absorbent
13.21 Ozone, O3
3.50 Ozone, O3
13.22 Pentane, C5H12
16.1.8 Pentane, C5H12
13.23 Petroleum gas
16.1.12 Fractional distillation of crude oil
13.24 Propane, C3H8
16.1.6 Propane, C3H8
13.25 Sulfur dioxide, SO2, sulfur (IV) oxide
Sulfur dioxide is a colourless gas that irritates the lungs.
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
12.18.4 Properties of sulfur dioxide and sulfites
13.13.3 Prepare sulfur dioxide with dilute sulfuric acid and sodium sulfite
13.13.4 Prepare sulfur dioxide with sulfuric acid and copper
13.13.8 Dry hydrogen sulfide and dry sulfur dioxide will not react
13.26 Trichloromethane, CCl3
16.1.14 Trichloromethane: Prepare trichloromethane (chloroform), CCl3
13.27 Triodomethane, CHI3
16.1.13 Triodomethane: Prepare triodomethane (iodoform), CHI3
13.1.5 Relative molecular mass of gases
See diagram 13.1.5 | See also: Saturated vapour pressure over water
The relative molecular mass, M, of a compound is the ratio of the average mass of molecules of the substance to 1 / 12 of the mass of one atom of C-12. Number of moles = volume in litres / 22.4 litres / mol. At s.t.p., 1 mol of most gases occupies 22.4 L at s.t.p.. At 25oC, 1 mol of most gases occupies 24.45 L
Weigh a gas container. Collect 1 litre of gas in an inverted measuring cylinder over water. The levels of water inside and outside the measuring cylinder must be the same. Weigh the gas container again. Calculate the loss in weight (about 2 g). Note the temperature and atmospheric pressure. For propane, if loss in weight of gas container = 1.8 g, 1.8 X 24.45 = 44 = relative molecular mass of propane. Use other sources of gas, e.g. a cigarette lighter. Hold it under water below the measuring cylinder with the valve kept open with a rubber band.
13.1.6 Molar volume of oxygen prepared with hydrogen peroxide
See diagram: 13.1.6 | See also: Table of saturated vapour pressure over water, Psvp
Put 15 mL of 3 % w/w (3 g H2O2  /100 g solution) hydrogen peroxide solution into flask A. Put 0.05 g of yeast in a small test-tube then lower the test-tube into flask A. Weigh flask A and its contents, W1. Attach Plastic tube 1 and Plastic tube 2 to flask B only. Put water into flask B leaving a space in the neck of the flask. Add water to a beaker until it is one third full. Siphon water into the beaker by blowing into the open end of Rubber tube 1 or by using a pipette bulb. Raise and lower the beaker to remove any air bubbles from Plastic tube 2. Adjust the height of the beaker so that the levels of the water in the beaker and in flask B are the same. Connect Plastic tube one to flask A. Raise the beaker to check for leaks in the apparatus. Again, adjust the height of the beaker so that the levels of the water in the beaker and in flask B are the same. Close the pinch clamp. Replace the glass tube in the beaker and open the pinch clamp to allow some water to flow into the beaker. With the pinch clamp still open, tip flask A so that the yeast falls into the hydrogen peroxide solution. Swill flask A until the reaction is completed when the water level in the beaker does not change. Again, adjust the height of the beaker so that the levels of the water in the beaker and in flask B are the same. Close the pinch clamp. Remove the stopper in flask A, insert a thermometer and note the temperature of the gas inside, T1. Repeat this measurement with flask B, T2. Disconnect Plastic tube 1 from flask A and again weigh flask A and its contents, W2. Measure the oxygen produced, by measuring the volume or weight water in the beaker, V. Find the vapour pressure of water at that temperature from the Table of saturated vapour pressure over water, Psvp. Note the room temperature. Note the barometric pressure from a barometer or ask the weather bureau or local airport.
Calculate the volume at s.t.p. of 32 g, one mole, of oxygen gas.
W2-W1 = Loss in weight
VO2 = volume of water in the beaker = volume of oxygen collected
Tf = average temperature in the flasks = (T1 + T2) /2
Patm = atmospheric pressure = pressure of oxygen in the flask, PO2 + saturation vapour pressure of water at that temperature, Psvp
so PO2 = Patm - Psvp
VO2 = volume of oxygen in the flask
Tstp = temperature at s.t.p. = 0oC, 273 K
Pstp = pressure at s.t.p. = 760 mm Hg = 101325 Pa
Vstp = volume at s.t.p.
P1V1 / T1 = P2V2 / T2 (Boyle's law and Charles's law)
(P1 X V1) / T1 = (P2 X V2) / T2
(PO2 X VO2) / Tf = (Pstp X Vstp) / Tstp
So Vstp = VO2 [(PO2 / Pstp) X (Tstp / Tf)]
Relative molecular mass of oxygen = 32 g
So number of moles of oxygen = (W2-W1) / 32
So the molar volume of oxygen at stp = Vstp / number of moles of oxygen = litres / mole
A mole of an ideal gas occupies 22.4 litres at s.t.p.
If barometric pressure = 1016 kPa, average temperature = 20oC, loss in weight of flask = 0.2 g, volume of oxygen collected at average temperature = 140 mL
Pressure of oxygen in apparatus = (barometric pressure - SVP at 20oC) = 1016 - 2.3 = 1013.7 kPa
Vstp = 140 [(1013.66 / 101.325) X (293 / 273)] = 503.17 mL
Number of moles = 0.2 / 32 = 0.00625 moles
Molar volume = 140 / 0.00625 = 22400 = 22.4 litres = 22.4 L / mole
13.2.3.0 Prepare hydrogen with iron filings and citric acid, sulfuric acid, or sodium hydrogen sulfate
Put 1 cm depth of iron filings in a test-tube. Just cover the iron filings with a dilute acid solution. Warm the test-tube until frothing starts. Hydrogen is colourless and odourless but any impurities in the iron filings give a nasty smell. To test for hydrogen remove from heating, place your thumb over the end of the test-tube, count to five, apply a lighted paper to the end of the test-tube, the hydrogen explodes with a loud pop sound. Never test more than a test-tube full of hydrogen gas!
13.2.3.1 Prepare hydrogen with aluminium foil and sodium carbonate, washing soda
Cut into small pieces some aluminium foil or aluminium milk bottle top and put into a test-tube. Add 5 mL sodium carbonate solution (Na2CO3.10H2O, washing soda). Heat until effervescence.
13.2.3.2 Prepare hydrogen with iron filings and alum
Put 5 g of iron filings in 1 cm depth of alum solution [Al2(SO4)3.K2(SO4).24H2O, potash alum, alum] [also shown as KAl(SO4)2.12H2O] in a test-tube. Heat the solution until effervescence occurs.
13.2.4 Prepare hydrogen bubbles
Hydrogen is much lighter than air and was used in airships. It has now been replaced by helium because hydrogen ignites easily and is thus too dangerous to use.
Put soapy water or detergent into the gas bubbler. Prepare hydrogen as described above. Bubbles of hydrogen form as the gas passes through the soapy water. Shake them gently to make them float up. Try to ignite the bubbles with a lighted splint.
13.2.5 Reduce metal oxides to metals with hydrogen
See diagram 13.2.4
Pass hydrogen over copper (II) oxide, or lead (II) oxide (lithage), or iron (III) oxide. Hydrogen reduces metal oxides to metals. The products are the metal and water.
CuO(s) + H2(g) ---> Cu(s) + H2O(l)
13.3.1 Prepare oxygen foam with hydrogen peroxide
Pour dilute hydrogen peroxide into a measuring cylinder. Add drops of detergent. Add manganese (IV) oxide (manganese dioxide) powder. The reaction forms oxygen gas as a foam of bubbles. Use the oxygen foam for combustion experiments with burning twine, burning iron wire and burning magnesium. Test the gas in the space above the liquid.
13.3.2 Burn sulfur in oxygen
Dip a wire loop into sulfur powder. Ignite the sulfur in a burner flame and then put it into a test-tube of oxygen. The sulfur burns with a bright blue flame.
13.3.3 Burn steel wool in oxygen, burn iron filings
Collect oxygen in test-tubes with stoppers. Store test-tubes in a test-tube rack and remove the stoppers just before inserting the burning element. Fasten steel wool to wire. Heat the steel wool in a burner flame. Put it into a test-tube of oxygen. The steel wool burns with bright sparkles to form grey-black iron oxide, Fe3O4(FeO.Fe2O3) magnetite, lodestone, iron ore.
Repeat the experiment by sprinkling iron filings into a Bunsen burner flame. A shower of sparks occurs as in some fireworks.
13.3.4 Burn magnesium ribbon in oxygen
Wrap a 3 cm piece of magnesium ribbon around the loop at the end of a wire. Ignite it in a burner and put it quickly in the oxygen. Magnesium burns with a very bright flame.
BE CAREFUL! Do not look directly at the flame because its brightness can cause injury to eyes. The white smoke is magnesium oxide. Its toxicity is low, but inhalation should be avoided.
Put the ash on a watch glass and add 3 mL of deionized water to wet the ash thoroughly and leave it lying in a small pool of water. Add one small drop of phenolphthalein solution and leave to stand for two minutes. Magnesium oxide has a low solubility in water, so you will not see any visible evidence that any of the solid has dissolved. Add one drop of dilute hydrochloric acid solution and leave to stand until the solution around the solid ash will turn pink, showing that the solution has become alkaline. This is evidence that some magnesium oxide has dissolved. Oxide ions in the solid react with water to form aqueous hydroxide ions. When no further change occurs, add a second drop of dilute hydrochloric acid. The pink colour disappears almost instantly, showing that the hydroxide ions have been neutralized very quickly, and replaced by an excess of hydrogen ions. During the next 2 to 15 minutes, depending on the size and concentration of the drop of acid added, the mixture changes slowly back to pink as the excess acid being neutralized slowly by solid magnesium oxide, followed by slow dissolving of remaining magnesium oxide to make the solution. When no more changes occur, add a second small drop of dilute hydrochloric acid. The same cycle of discharge and reappearance of pink colour can be repeated for as long as any solid magnesium oxide remains. Magnesium oxide is a metallic oxide, and is therefore basic. Magnesium oxide has a low solubility in water. Dilute hydrochloric acid reacts rapidly with aqueous magnesium hydroxide, but slowly with solid magnesium oxide. Magnesium oxide dissolves slowly in water. Phenolphthalein is an indicator that shows changes in alkalinity of the solution. An equilibrium is established between solid magnesium oxide and dissolved magnesium ions. The addition of acid disrupts the equilibrium by removing hydroxide ions from the solution. An equilibrium is established between solid magnesium oxide and dissolved magnesium ions. The addition of acid disrupts the equilibrium by removing hydroxide ions from the solution. Equilibrium is restored by slow dissolving of more magnesium oxide. Addition of larger drops or higher concentration of acid causes a larger initial excess of acid in the solution. Because the reaction of acid with the solid magnesium oxide is slow, it will take a much longer time for the pink colour to return to the mixture. The magnesium oxide formed from combustion of magnesium ribbon forms a hard mass with a small surface area for reaction. The rate of reaction with acid, and the rate of solution of the solid to form an alkaline solution, would be increased by crumbling the ash.
13.3.5 Prepare an oxygen absorbent
Dissolve 300 g of ammonium chloride in 1 litre of water and add 1 litre of concentrated ammonia solution. Shake the solution. Pass the gas through the solution after adding half the volume of copper turnings.
13.4.0 Chlorine
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. It dissolves readily in water forming highly corrosive solution. Do not prepare chlorine in an open room but use a fume cupboard. Direct combination of chlorine and hydrogen occurs in bright light or ignition of the mixture by lighted taper or electric spark. It reacts with metals, solid non-metals, hydrocarbon. Use small quantities only. Chlorine is a yellow-green, dense gas that causes rapid corrosion of metals and destruction of plastics. It is also a dangerous gas because it attacks the mucous membrane linings of the eyes, nose, throat and lungs, causes the lungs to fill with fluid and the victim drowns. During the First World War it was used as a chemical weapon. Chlorine is prepared commercially from electrolysis of concentrated sodium chlorine (brine) solution. Chlorine is a very reactive non-metal and free chlorine never occurs naturally. Do all chlorine experiments in a fume cupboard. Chlorine kills most living things and is used to sterilize drinking water and disinfect swimming pools. Chlorine is used to manufacture the plastic PVC, to bleach wood pulp and to prepare organic compounds such as solvent tetrachloroethene CCl2.CCl2, solvent tetrachloromethane (carbon tetrachloride) CCl4, safer solvent 1,1,1-trichloroethane CH3CCH3 and the insecticide DDT (C6H4Cl)2CH-CCl3 [Former name: dichlorodiphenyltrichloroethane, New IUPAC name: 1,1,1-trichloro-2,2-bis (4-chlorophenyl)ethane] However, many of these substances cannot be broken down in the environment (biodegraded), so you should avoid using them. Chlorine is a powerful oxidizing agent.
13.4.4 Prepare chlorine with dilute hydrochloric acid and domestic bleach solution
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)
13.4.5 Prepare chlorine with concentrated hydrochloric acid and manganese (IV) oxide
Put some manganese (IV) oxide in a boiling tube and add drops of concentrated hydrochloric acid from the reservoir. Heat the test-tube gently. Observe the slight green colour in the tube. The wider the tube, the easier this is to see.
Use potassium manganate (VII) instead of manganese (IV) oxide to prepare chlorine because the reaction does not require heating avoiding hot concentrated hydrochloric acid.
Be careful! Prepare chlorine only in a fume cupboard
4HCl(aq) + MnO2(s) ---> MnCl2(aq) + 2H2O(l) + Cl2(g)
13.4.6 Prepare chlorine with concentrated acid and potassium manganate (VII)
Connect a conical flask by means of a delivery tube to a collection vessel. Put about 5 g of solid potassium manganate (VII) (potassium permanganate) in a conical flask and add concentrated hydrochloric acid drop by drop.
16HCl(aq) + 2KMnO4(s) --->2KCl(aq) + 2MnCl2(aq) + 8H2O(l) + 5Cl2(g)
13.4.7 Reactions of chlorine with sodium
See diagram. 13.4.7
BE CAREFUL! THE REACTION IS VERY VIGOROUS! Do this experiment in a fume cupboard.
1. Dry a small piece of sodium with absorbent paper. Grip a piece of sodium with a pair of tongs. File the sodium and let the obtained sodium filings fall into chlorine gas collected in a test-tube. The sodium filings react violently with the chlorine, sparks flying off, to form many smoke particles of sodium chloride and a crust of sodium chloride on what is left of the piece of sodium. When the reaction has stopped, wash the residue in methylated spirit to remove unreacted chlorine. Let chlorine leave the test-tube by diffusion. Crystals of sodium chloride remain in the test-tube.
2Na(s) + Cl2(g) ---> NaCl(s)
2. Put a pin head volume of sodium in the bowl of a deflagrating spoon. In a fume cupboard, put the spoon into a test-tube of chlorine and allow to stand. When the reaction stops, remove the spoon, allow it to cool and place it in a small amount of alcohol. Let excess chlorine diffuse away in the fume cupboard. Let the mixture of alcohol and solid stand until no further reaction takes place. Wash the crystals with alcohol and let them cool and dry.
sodium(s) + chlorine(g) ---> sodium chloride(s)
13.4.8 Burn steel wool in chlorine
Ignite steel wool held by tongs in a burner flame, then put into chlorine gas.
BE CAREFUL! The reaction occurs with strong combustion to form a red-brown cloud that condenses to black flakes of anhydrous iron (III) chloride.
2Fe(s) + 3Cl2(g) ---> 2FeCl3(s)
13.4.9 Burn copper in chlorine
Heat copper foil in a burner flame and put into chlorine gas. The reaction forms a layer of brown copper (II) chloride that turns green in the presence of moisture.
Cu(s) + Cl2(g) ---> CuCl2(s)
13.4.10 Burn a wax taper in chlorine, reaction of chlorine with non-metals
1. Chlorine has such a strong attraction for hydrogen that it removes all the hydrogen from the hydrocarbon paraffin leaving behind the carbon as a residue. A mixture of chlorine and hydrogen does not react in the dark but if heated or exposed to strong sunlight the mixture reacts explosively to form hydrogen chloride.
BE CAREFUL! Do not mix chlorine and hydrogen!
H2(g) + Cl2(g) ---> 2HCl(g) + energy
2. A mixture of chlorine and methane explodes violently in direct sunlight forming hydrogen chloride and free carbon.
BE CAREFUL! Do not mix chlorine and Methane!
CH4(g) + 2Cl2(g) ---> C(s) + 4HCl(g) + energy
3. Chlorine reacts with benzene, C6H6, to form a substitution product dichlorobenzene, C2H4Cl2, called "mothballs" or "moth crystals". Nowadays people are advised to use camphor instead of moth balls to protect their clothes from being eaten by moths.
C6H6(l) + Cl2(g) ---> C2H4Cl2(l)
4. Burn a paraffin wax taper or a small birthday candle in chlorine. The taper keeps burning with a dull red flame and forms black carbon particles (soot) and hydrogen chloride gas.
nCl2 + (CH3-CH2-CH2-CH2-) ---> nHCl + nC
13.4.11 Pass chlorine through water
Chlorine is available commercially for school laboratory use as chlorine water. Hypochlorous acid HClO is a bleach and a disinfectant. Hypochlorous acid is an aqueous solution of chlorine (I) oxide that forms salts called hypochlorites. Hypochlorous acid is a weak acid that easily decomposes back to chlorine gas and water. When chlorine passes through water, a mixture of HCl and HClO forms. The chlorine is oxidized and reduced.
Cl2(g) + H2O(l) <---> HCl(aq) + HClO(aq)
13.4.12 Pass chlorine through iodine solution
The more reactive chlorine displaces iodine from its salt. The colourless potassium iodide solution turns red then black as iodine is displaced from the solution. Tests for iodine with starch solution.
Cl2(g) + 2KI(aq) ---> I2(aq) + 2KCl(aq)
13.4.13 Pass chlorine through iron (II) chloride solution
Pass chlorine through iron (II) chloride solution. The chlorine oxidizes iron (II) chloride to iron (III) chloride. The solution changes from green to brown.
2Fe2+(aq) + Cl2(g) ---> 2Fe3+(aq) + 2Cl-(aq)
FeCl2(aq) + Cl2(g) ---> 2FeCl3(aq)
13.4.14 Reactions of chlorine with heated copper and steel wool
BE CAREFUL! Do not breath this poisonous gas.
1. In a fume cupboard, put a heated spiral of copper wire into a small test-tube of chlorine. The heated copper is immediately covered with brown copper (II) chloride that turns green in the presence of water.
Cu(s) + Cl2(g) ---> CuCl2(s)
2. Heat a lump of steel wool and plunge it into the chlorine gas. Brown fumes form that condense to black flakes of anhydrous iron (III) chloride.
2Fe(s) + 3Cl2(g) ---> 2FeCl3(s)
13.4.15 Reactions of chlorine with alkalis, bleaching powder
1. Pass chlorine slowly through test-tubes containing dilute sodium hydroxide solution, dilute potassium hydroxide solution and solid calcium hydroxide. Add dilute sulfuric acid to the products of any reactions. Chlorine reacts with cold alkali solutions to form chloride ions, Cl- and hypochlorite ions, ClO-, powerful bleaching agents.
Cl2(g) + 2OH-(aq) ---> Cl-(aq) + ClO-(aq) + H2O(l)
2. Pass excess chlorine into hot alkali solutions to form chloride and chlorate ions. If passed into potassium hydroxide, the less soluble potassium chlorate can be separated from the less soluble potassium chloride by fractional distillation.
3Cl2(g) + 6OH-(aq) ---> Cl-(aq) + ClO3-(aq) + 3H2O(l)
3. Chlorine reacts with strongly basic hydroxides, e.g. calcium hydroxide, and strongly basic oxides, e.g. calcium oxide, in the solid state. The products have a variable composition. Reactions of chlorine with calcium hydroxide produces bleaching powder, a convenient source of chlorine and a powerful bleaching agent in dilute acid solutions. Bleaching powder reacts with sulfuric acid to give off chlorine.
Bleaching powder(s) + sulfuric acid(aq) ---> calcium sulfate(s) + chlorine(g) + water(l)
13.6.5 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. Tests 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 Catalytic oxidation of ammonia with red-hot platinum wire
See diagram 13.06.6A
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 redhot 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)
13.6.6.1 Catalytic oxidation of ammonia with chromium (III) oxide catalyst
See diagram 13.06.6B
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 IT IN! AVOID USING GLASS WOOL!
Heat the catalyst tube for about 2-3 minutes to raise the temperature of the catalyst to above 500oC. 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 redhot, 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 red-brown 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 13.06.7
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)
13.7.6 Prepare carbon dioxide by heating 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 by heating 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)
13.7.13 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.
13.9.0 Nitrogen
Nitrogen gas N2 is colourless, odourless, tasteless, neutral and unreactive. Nitrogen does not support combustion. Magnesium and calcium will continue to burn in nitrogen to form nitrides. Nitrogen is manufactured by fractional distillation of air. Air contains about 78% of nitrogen.
13.9.3 Nitrogen gas generated in a motor car airbag
A gas generator containing a mixture of sodium azide, NaN3, potassium nitrate, KNO3 and silica, SiO2 is ignited electrically to allow a slow detonation so that nitrogen fills the airbag.
2NaN3 ---> 2Na + 3N2 (300oC)
10Na + 2KNO3 ---> K2O + 5Na2O + N2
K2O + Na2O + SiO2 ---> alkaline silicate
13.10.2 Prepare nitrogen dioxide from lead (II) nitrate crystals
Heat lead (II) nitrate crystals. The decomposition may be noisy. Nitrogen dioxide and oxygen form, leaving yellow lead oxide.
Pb(NO3)2(s) ---> 4NO2(g) + 2PbO(s) + O2(g)
lead (II) nitrate ---> nitrogen dioxide + lead oxide + oxygen
13.12.0 Sulfur dioxide
Sulfur dioxide, SO2, is a colourless gas that irritates the lungs. Sulfur dioxide dissolves in water to form, mainly, sulfurous acid (H2SO3). Sulfur dioxide is one component of acid rain.
SO2(g) + H2O(g) ---> H2SO3(l)
13.13.3 Prepare sulfur dioxide with dilute sulfuric acid and sodium sulfite
Na2SO3(s) + H2SO4(l) ---> Na2SO4(aq) + H2O(l) + SO2(g)
13.13.4 Prepare sulfur dioxide with sulfuric acid and copper
Add hot concentrated sulfuric acid to copper to form copper (II) sulfate, water and sulfur dioxide. BE CAREFUL!
Cu(s) + 2H2SO4(l) ---> CuSO4(aq) + 2H2O(l) + SO2(g)
13.13.8 Dry hydrogen sulfide and dry sulfur dioxide will not react
Collect sulfur dioxide in a dry test-tube after passing the gas slowly through concentrated sulfuric acid to dry it. Collect a test-tube of hydrogen sulfide, after passing it over calcium chloride tube to dry it. Invert the test-tube containing sulfur dioxide over the test-tube containing the hydrogen sulfide. No reaction occurs. Leave to stand then pour drops of water into the lower test-tube and quickly replace the upper test-tube. sulfur immediately precipitates in the test-tubes.
2H2S + SO2 ---> 2H2O + 3S (s)