School Science Lessons
Topic 12C
2024-02-02
Halogens, Astatine, Bromine, Chlorine
(topic12C) Contents
Halogens
Astatine, At
Bromine, Br
Chlorine, Cl
Halogens
See: Table 12.19.1.0
CFCs, chlorofluorocarbons, "Freons": 12.19.5.0
Composition of the atmosphere and greenhouse gases: 37.42.1
Halogenation of benzene: 12.4.6.1
Halogens, Group 17 of the periodic table: 12.19.0
Halogen compounds, haloalkanes: 12.4.17
Halogen experiments: 12.19.2
Halogen oxyacids: chloric acid HClO3 | chlorous acid HClO2, weak unstable acid, | hypochlorous acid HClO | perchloric acid HClO4 |
Properties of halogens: 12.19.2.5
Tests for halides, Cl-, Br-, I-: 12.11.10.
12.19.2 Halogen experiments
12.19.2.1 Compare halogens, chlorine, bromine, iodine
12.19.2.2 Displace a less reactive halogen from halogen compounds
12.19.2.3 Halide salts with hot concentrated sulfuric acid
12.19.2.4 Halide vapour over hot iron wire forms iron halides
12.19.2.5 Properties of halogens
12.19.2.6 Reactions of silver halides, photography.
Astatine, 85At
Astatine, At2, (Greek astatos unstable), radiation element, halogen group
Astatine, Table of Elements
Astatine, RSC
Astatine has the longest life of any radioactive isotope.
Bromine, Br
Bromide ion: Br-
See: Bromine, Br, Table of elements
See: Bromine, RSC
Bromine compounds: 12.19.3
Bromine experiments: 12.19.6
Bromides hazards: 3.7.1
Bromine products, (used in swimming pools): 18.7.19 HBrO
Bromine water, bromine solution, Br2 (aq), Toxic by all routes
Brominated flame retardants (BFRs): 3.0.5
Bromination of acetone: 12.19.9.7
Opening a bromine gas ampoule: 12.19.9.1.2, (Safety)
Reactions of bromo-compounds: 12.18.0
BCDMH, Bromine products, HBrO
Bromoform, CHBr3, Tribromomethane
Bromine
Bromine is toxic by all routes, extremely irritant vapour, Not permitted in schools, vapour is highly toxic by inhalation, liquid causes severe burns to eyes and skin, aqueous solution attacks lungs, eyes and nose
Bromide (bromo), Br-, monodentate ligand
Tests for bromides
"Bromide", bromine + less electronegative substance, common name for sedative:
Bromic acid, hydrobromic acid, HBr, hydrobromic acid solution, hydrogen bromide solution, Harmful
Brominated vegetable oil stabilizer has been removed from the sports drink "Gatorade":
2-Bromofluorene, 9,9'-Spirobifluorene, 2-Bromo-9,9'-spirobifluorene, 4-Bromo-9,9-spirobifluorene.
12.19.3 Bromine compounds
Bromoform, Tribromomethane: 12.18.21
Ethyl bromide, Bromoethane: 12.18.7
Ethylene bromide, Cassava project, 17, Cassava, Pests and diseases
The different fungi that cause storage rot of tubers can be treated with ethylene bromide.
Ethylene dibromide, 1,2-Dibromoethane: 12.18.16
Halon-1211, bromochlorodifluoromethane CBrClF2: Table 12.19.5.0, RODP
Halon-1301, bromotrifluoromethane CBrF3: Table 12.19.5.0, RODP
16.14.1 PBDE, pentabromodiphenyl ether, C12H5Br5O, PBB, polybrominated biphenyls, 4,4'-dibromobiphenyl, C12H8Br2.
12.19.6 Bromine experiments
Bromides, Tests for bromides: 12.11.6
Bromination of acetone: 12.19.9.7
Bromine as an oxidizing agent: 15.2.2.1
Bromine catalyses the oxidation of sulfur to sulfuric acid: 17.3.4
Bromine water oxidizes sodium thiosulfate to sodium sulfate and sulfur: 12.1.44
Halide salts with hot concentrated sulfuric acid: 12.19.2.3, (See 2.)
Heat hydrogen bromide: 12.19.9.3a
Potassium bromate with propanedioic acid, double autocatalytic reaction, oscillating reaction: 17.3.8
Prepare bromine water: 12.19.9.1.1
Prepare hydrogen bromide: 12.19.9.3
Prepare potassium bromide, KBr: 12.19.9.5
Reactions of 2-bromopropane with hydroxide, ions, nucleophilic reactions: 12.19.9.8
Reactions of bromides, Br: 12.19.9.6
Reactions of bromine, Br2: 12.19.9.1
Reactions of bromine water (bromine solution): 12.19.9.2
Reactions of hydrogen bromide, HBr: 12.19.9.4.
Tests for saturated hydrocarbons, bromine water test: 9.4.7.0
12.1.44 Bromine water oxidizes sodium thiosulfate to sodium sulfate and sulfur
Bromine water, bromine solution, Br2 (aq), Toxic by all routes
Na2S2O3 + Br2 + H2O --> Na2SO4 + 2HBr + S
Chlorine, Cl
See: Chlorine, Table of Elements
See: Gas, Molecular weight, Density, (Table 1)
See: Chlorine, RSC
Chloride ion: Cl-
Chlorides, List of chlorides, 1.9
Chlorates
Chlorides
Chlorine compounds: 12.19.9
Chlorine, properties: 7.2.2.12
Chlorine as an oxidizing agent: 15.2.2
Chlorine deficiency in soils: 6.13.2
Chlorine hazards: 3.8.4
Chlorine used in swimming pools: 18.2.0
Chlorine with water (redox reaction, disproportionation): 12.2.6.4
Direct union of elements to form compounds: 8.0.0
Prepare chlorine: 12.19.7
Reactions of chlorine: 12.19.8
Tests for chlorine: 3.40.1
Tests for chlorine levels in swimming pools, test kit: 18.7.21.0.
12.19.7 Prepare chlorine
Prepare chlorine: 12.4.1.0
Prepare chlorine water: 12.4.2
Prepare chlorine with bleaching powder or bleach solution: 12.4.1.1
Prepare chlorine with hydrochloric acid / potassium permanganate: 12.4.1.2
Prepare chlorine with sodium chloride: 12.4.1.3
Prepare chlorine water: 12.4.2
12.19.8 Reactions of chlorine
Chlorine reagent test kit
Activated carbon, (See: A. and C. chlorine removal ability): 2.3.0
Burn substances in chlorine: 12.4.8
Chlorine used in swimming pools: 18.2.0
Chlorine water oxidizes sodium thiosulfate to sodium sulfate and sulfur: 12.1.43
Heat metals in chlorine: 7.1.3
Pass chlorine through iodine solution: 12.4.12
Pass chlorine through iron (II) chloride solution: 12.4.13
Pass chlorine through water: 12.4.9
Reactions of chlorine with alkalis, bleaching powder: 12.4.15
Reactions of chlorine with copper: 12.4.14
Reactions of chlorine with benzene: 12.4.6.0
Reactions of chlorine with group 3 elements, Na, Mg, Al, Si, P: 12.4.7.1
Reactions of chlorine with sodium: 12.4.7
Reactions of chlorine with steel wool: 12.4.16.
12.19.9 Chlorine compounds
(2,4-D), (2,4-dichlorophenoxyacetic acid), C8H6Cl2O3
(2,4,5-T), (2,4,5-trichlorophenoxyacetic acid), C8H5Cl3O3
Acyl chlorides, acid chlorides (acyl = RC=O-)
Arachidonoyl chloride, C20H31ClO
Bifenox herbicide: 4.4.4.
Calcium hypochlorite: 18.7.2.2.1
CFCs, chlorofluorocarbons, "Freons": 12.19.5.0
Chloral: 16.3.2.1
Chloramines in swimming pools: 18.7.23.0
Chloramphenicol, C11H12Cl2N2O5: 7.9.14.2
Chlorate ion: ClO3-
Chlorides, List of chlorides, 1.9
Chlorates
Chlordane, C10H6Cl8, pesticide: 4.4.01 (Agriculture)
Chloric acid, HClO3
Chlorides
Chlorine dioxide, ClO2: 12.4.5
Chlorine monoxide, dichlorine oxide, Cl2O
Chlorites, ClO2-
Chloro-compounds, List of chloro-compounds: 1.10
Chloroacetic acid, CH2ClCOOH
Chlorofluorocarbons, CFCs, "Freons": 12.19.5.0
Trichloromethane, (chloroform), CHCl3
Chloromethane, CH3Cl, methyl chloride
Chlorophenol red, C19H12Cl2O5S, (acid-base indicator): 9.0
Chlorophenols
Chloroplatinic acid, H2PtCl6.6H2O
Chlorothalonil fungicide, C8Cl4N2: 46.6.5
Chloroethene (vinyl chloride) CH2:CHCl
CS gas, C10H5ClN2, C6H4(Cl)CHC(CN)2, (2-chlorobenzalmalononitrile), orthochlorobenzylidene malononitrile, "tear gas".
Dichloroacetic acid, CHCl2COOH
Dichlorobenzene, C6H4Cl2
Dicofol C14H9Cl5O.
Ethanoyl chloride, (Acetyl chloride, CH3COCl
Ethylene dichloride (1.2-dichloroethane, Freon 150), C2H4Cl2, ClH2C-CH2Cl, toxic, highly flammable.
Freons, CFCs, chlorofluorocarbons, "Freons": 12.19.5.0
Phosgene, carbonyl chloride, COCl2 : 13.39.2
Thiophosgene, thiocarbonyl dichloride, CSCl2, (chlorine + sulfur)
Tetrachlorvinphos: 4.4.5.
Trichlorophon: 4.4.6.
Chlorine experiments
3.5.11A Aluminium chloride in fireproof cloth
12.1.8 Aluminium chloride with water
15.5.12 Electrolysis of sodium chloride solution:
12.4.4 Bleaching powder
16.5.1 Prepare ethyl chloride, chloroethane, C2H5Cl
12.19.6.4 Compare silver chloride, silver bromide and silver iodide
12.19.8.4 Concentrated sulfuric acid with potassium chlorate
14.0 Electrophiles and nucleophiles, hydrogen chloride
11.4.1 Electric writing, sodium chloride with litmus paper
15.6.3 Electrolysis of saturated sodium chloride solution
12.4.6.1 Halogenation of benzene
Hydrochloric acid (List)
3.42.0 Hydrogen chloride, HCl (List)
3.42.1 Hydrogen chloride, Tests for hydrogen chloride (List)
3.7.7 Hypochlorites, hazards
16.5.1.3 Methane with chlorine, (Dangerous experiment!)
4.5a Organochlorine insecticides, chlorinated hydrocarbons
12.19.8.8 Phosphorus pentachloride with water
16.5.1 Prepare ethyl chloride, chloroethane, C2H5Cl
12.19.8.7 Prepare chromyl chloride, CrO2Cl2
3.42.0 Prepare hydrogen chloride / hydrochloric acid
12.1.30 Prepare hydrochloric acid with sodium chloride
12.19.8.3 Prepare iron (III) chloride
12.19.8.5 Prepare potassium perchlorate
12.18.3.2 Prepare thionyl chloride, SOCl2
16.1.14 Prepare trichloromethane, (chloroform)
12.19.8.1 Reactions of sodium chloride
12.19.8.6 Recover silver from silver chloride, AgCl2
3.43.3 Reduce iron (III) chloride with hydrogen sulfide
3.51.3 Reduce iron (III) chloride with sulfur dioxide
7.2.5.1 Silicon tetrachloride with water
3.42.1 Tests for hydrogen chloridezzz
12.1.43 Chlorine water oxidizes sodium thiosulfate to sodium sulfate and sulfur
Na2S2O3 + Cl2 + H2O --> 2HCl + Na2SO4 + S
Sodium thiosulfate destroys unreacted chlorine in the process of bleaching, so it acts as an antichlor.
Chlorates
Chlorates, chlorate ion ClO3-, perchlorate ion ClO4-: Hazards: 3.7.2
Decomposition of chlorates: 3.7.4, potassium chlorate
Chlorates hazards: 3.7.2
Chlorides
Chlorides hazards: 3.7.3
Decomposition of chlorides: 3.7.5
Reactions of sodium chloride: 12.19.8.1
List of chlorides: 1.9
Tests for chlorides: 12.11.8
Tests for chlorides in groundwater: 18.2.2.1.
12.4.8 Burn substances in chlorine
12.4.8.5 Burn candles in chlorine
12.4.8.2 Burn copper in chlorine
12.4.8.7 Burn ethylene in chlorine
12.4.8.3 Burn magnesium in chlorine
12.4.8.8 Burn methane in chlorine
12.4.8.4 Burn paraffin wax tapers in chlorine
12.4.8.1 Burn steel wool in chlorine
12.4.8.6 Burn wood splints in chlorine.
12.19.0 Halogens, Group 17 of the periodic table
Halogens, Group 17 Periodic Table: 1.6.0
The halogens are yellow gas fluorine F2, green-yellow gas chlorine Cl2, red-brown liquid bromine Br2, violet solid iodine I2, and the rare element Astatine, At.
Halogens form halide ions F-, Cl-, Br- and I-, are strong oxidizing agents, react with alkali metals to form salts, react with hydrogen in the decreasing order F-, then Cl-, then Br-, then I-.
Halogens react with most elements, with alkali metals to form salts.
Halogens are toxic (poisonous) and are more soluble in hydrocarbon solvents than in water.
The hydrogen halides are gases that dissolve in water to form acidic solutions that conduct electric current.
12.19.2.4 Halide vapour over hot iron wire forms iron halides
12.4.8.1 Burn steel wool in chlorine.
12.19.2.1 Compare halogens, chlorine, bromine, iodine
1. Add drops of 0.02 M chlorine solution to Universal indicator paper on a white tile and note any change.
The chlorine solution bleaches the Universal indicator paper.
Chlorine in water forms hydrochloric acid and bleach (chloric (I) acid).
* Add drops of 0.02 M bromine solution, to Universal indicator paper on a white tile and note any change.
The bromine solution bleaches the area on the Universal indicator paper where it touched it.
* Add drops of 0.02 M iodine solution to Universal indicator paper on a white tile and note any change.
The bromine solution faintly bleaches the area on the Universal indicator paper where it touched it.
2. Add 0.02 M chlorine solution to 0.2 M potassium bromide solution
Cl2 (aq) + 2KBr (aq) --> 2KCl (aq) + Br2 (aq) orange red-brown solution
chlorine + potassium bromide --> potassium chloride + bromine
Cl2 (aq) + 2Br- (aq) --> 2Cl- (aq) + Br2 (aq)
* Add 0.02 M chlorine solution to 0.2 M potassium iodide solution
Cl2 (aq) + 2KI (aq) --> 2KCl (aq) + I2 (aq) black precipitate in a brown solution
chlorine + potassium iodide --> potassium chloride + iodine
Cl2 (aq) + 2I- (aq) --> 2Cl- (aq) + I2 (aq).
3. Add 0.02 M bromine solution to 0.2 M potassium chloride solution - No reaction
* Add 0.02 M bromine solution to 0.2 M potassium iodide solution
Br2 (aq) + 2KI (aq) --> 2KBr (aq) + I2 (aq) black precipitate in a brown solution
bromine + potassium iodide --> potassium bromide + iodine
Br2 (aq) + 2I- (aq) --> 2Br- (aq) + I2 (aq).
4. Add 0.02 M iodine solution to 0.2 M potassium chloride solution - No reaction
* Add 0.02 M iodine solution to 0.2 M potassium bromide solution - No reaction.
Chlorine is a more powerful oxidizing agent than either bromine or iodine, so it can take electrons from both bromide ions and iodide ions.
Bromine is a more powerful oxidizing agent than iodine, so it can remove electrons from iodide ions to give iodine.
Oxidizing ability falls as you go down the Group, chlorine then bromine then iodine.
Halogens get less reactive as the group is descended.
The halogen molecule is the electron acceptor (the oxidizing agent) and is reduced by electron gain to form a halide ion.
The halide ion is the electron donor (the reducing agent) and is oxidized by electron loss to form a halogen molecule.
12.19.2.2 Displace a less reactive halogen from halogen compounds.
More reactive halogen displaces less reactive halogen from its compound.
1. Add iodine solution to colourless potassium bromide solution.
No reaction, because iodine is less active than bromine.
Pass chlorine gas through colourless potassium bromide solution.
The more active chlorine displaces the less active bromine and the solution turns orange.
2KBr (aq) + Cl2 (g) --> 2KCl (aq) + Br2 (aq).
2. Pass chlorine gas through potassium iodide solution.
The more active chlorine displaces the less active iodine and the solution turns deep brown.
12.19.2.3 Halide salts with hot concentrated sulfuric acid
The oxidizing power of F2 > Cl2 > Br2 > I2.
The reducing power of I- > Br- > Cl- > F-.
1. Sodium chloride with hot concentrated sulfuric acid
Sodium and potassium halides with concentrated sulfuric acid have similar reactions.
In each case the metal ion is a spectator ion.
A displacement reaction (not a redox reaction), forms a white residue of sodium hydrogen sulfate or sodium sulfate.
NaCl (s) + H2SO4 (l) --> NaHSO4 (s) + HCl (g)
2NaCl (s) + H2SO4 (l) --> Na2SO4 (s) + 2HCl (g)
Cl- (s) + H2SO4 (l) --> HSO4- (s) + HCl (g).
* Similarly, sodium or potassium fluoride with concentrated sulfuric acid forms hydrogen fluoride gas, HF.
KF (s) + H2SO4 (l) ---> KHSO4 (s) + HF (g) (the HF formed can etch glass!).
2. Sodium bromide with hot concentrated sulfuric acid
In this redox reaction, 1. the bromide ion acts as a reducing agent and is oxidized to form bromine gas, 2. the sulfuric acid is reduced to sulfur dioxide (sulfur (IV) oxide), 3. a white residue remains.
NaBr (s) + H2SO4 (l) --> NaHSO4 (s) + HBr (g)
2HBr (g) + H2SO4 (l) --> Br2 (g) + SO2 (g) + 2H2O (l)
2NaBr (s) + 2H2SO4 (l) --> NaHSO4 (s) + Br2 (g) + SO2 (g) + 2H2O (l)
2Br- --> Br2 + 2e- (oxidation, loss of 2 electrons, oxidation state of bromine changes from -1 to 0 )
H2SO4 + 2H+ + 2e- --> SO2 + 2H2O (reduction, gain of 2 electrons, oxidation state of sulfur changes from +6 in sulfuric acid to +4 in sulfur dioxide )
2Br- + H2SO4 + 2H+ --> Br2+ SO2 + 2H2O (redox equation).
3. Sodium iodide with hot concentrated sulfuric acid
In this redox reaction, the iodide ion acts as a stronger reducing agent and is oxidized to form iodine gas as a purple vapour to form iodine crystals when cool.
The sulfuric acid is reduced to hydrogen sulfide gas with a rotten egg gas smell and hydrogen iodide fumes.
A white residue remains of sodium sulfate is formed.
The iodine will crystallize out on the upper cooler parts of the test-tube.
NaI (s) + H2SO4 (l) --> NaHSO4 (s) + HI (g)
8HI (g) + H2SO4 (l) --> 4I2 (g/s) + H2S (g) + 4H2O (l)
8NaI (s) + 5H2SO4 (l) --> 4Na2SO4 (s) + 4I2 (gas / solid crystals) + H2S (g) + 4H2O (l)
2I- --> I2 + 2e- (oxidation, gain of 2 electrons, oxidation state of iodine changes from -1 to 0)
H2SO4 + 8H+ + 8e- --> SO2 + 4H2O (reduction, gain of 8 electrons).
The oxidation state of sulfur changes from +6 in sulfuric acid to -2 in hydrogen sulfide.
8I- + H2SO4 + 8H+ --> 4I2 + H2S + 4H2O (redox equation).
12.19.2.6 Reactions of silver halides, photography
Add silver nitrate solution to solutions of potassium chloride, potassium bromide and potassium iodide.
Silver chloride (white) silver bromide (pale yellow) and silver iodide (deep yellow) are all sensitive to light and are used in photography.
Carl Wilhelm Scheele (1742 - 1786) exposed silver chloride beneath water to light.
He added silver nitrate to precipitate new silver chloride, then added ammonia solution to the blackened chloride to produce diamine silver ion solution.
He also noticed that violet rays of the spectrum blackened the silver chloride much more than red rays.
Ag+ + Cl- + light energy --> Ag+ + Cl + e-, i.e. one electron lost from chlorine, oxidation of chlorine
Ag+ + e- --> Ag (metal) i.e. one electron gained by silver, reduction of silver, to form a dark image on film
AgCl + 2NH3 --> Ag(NH3)2+ + Cl-.
12.19.6.1 Prepare hydrogen iodide
See diagram 12.19.6.1: Prepare hydrogen iodide
1. Grind a 2 cc each of dry red phosphorus and iodine in a mortar and introduce into a boiling tube.
Add four drops of water and fit the boiling tube.
P4 + 6I2 --> 4PI3
PI3 + 3H2O --> H3PO3 + 3HI (g)
Heat the boiling tube to produce further quantities of hydrogen iodide.
2. Pass the gas into silver nitrate solution in a test-tube.
Note the yellow precipitate of silver iodide.
3. Pass the gas into a test-tube containing drops of 880 ammonia.
Note the white fumes of ammonium, iodide.
HI + NH3 --> NH4I.
4. Pass the gas into a test-tube containing drops of concentrated nitric acid.
The hydrogen iodide is easily oxidized to iodine and the nitric acid reduced to nitrogen dioxide.
2HNO3 + 2HI --> 2H2O + 2NO2 + I2.
5. Pass the gas for some time into a test-tube containing concentrated sulfuric acid.
The acid is reduced to sulfur dioxide, hydrogen sulfide or sulfur, showing that hydrogen iodide is a powerful reducing agent.
6. Heat the delivery tube with a Bunsen burner.
Note the violet vapours of iodine.
2HI <--> H2 + I2.
12.19.6.2 Reactions of iodides, I-
Use one crystal the size of a match head or 1 mL of a 10% solution for each reaction.
1. Grind the iodide with a small quantity of manganese dioxide and add 1 mL of concentrated sulfuric acid to the mixture in a test-tube.
Heat gently and observe the violet vapours of iodine.
MnO2 + 2KI + 2H2SO4 --> MnSO4 + K2SO4 + 2H2O + I2
With concentrated sulfuric acid alone iodine is also obtained on heating, because hydrogen iodide is a powerful reducing agent.
2. Add some silver nitrate solution to potassium iodide solution.
Note the yellow precipitate of silver iodide that is insoluble in both dilute nitric acid and ammonium hydroxide solution.
Ag+ + I- --> AgI (s).
3. Add drops of chlorine water to potassium iodide solution.
Note that iodine is given off.
The iodine turns the solution brown, and some black crystals may be seen at the surface.
Cl2 + 2I- --> I2 + 2Cl-.
4. Add drops of lead acetate solution to potassium iodide solution.
Note the yellow precipitate of lead iodide.
Pb2+ + 2I- --> PbI2 (s).
12.19.6.3 Prepare iodic acid and potassium iodate
Potassium iodate, KIO3
I2O5 + H2O --> 2HIO3
iodine pentoxide (iodic anhydride) + water --> iodic acid
1. To prepare iodic acid, use 5 g of iodine in a retort and add a measured 40 mL of fuming nitric acid.
Heat the retort on a sand tray, keeping the temperature high enough to promote action.
Collect any nitric acid that distils over and return it to the retort.
When the iodine has all been oxidized to white crystals of iodic acid, pour the contents of the retort into an evaporating basin and heat almost to dryness on a water bath.
Collect the crystals and dry between filter paper.
I2 + 10HNO3 --> 2HIO3 + 10NO2 + 4H2O
Heat some of the crystals in a dry test-tube, gently then strongly.
Note the formation of moisture to leave iodine pentoxide.
2HIO3 --> I2O5 + H2O
This is followed by decomposition to iodine and oxygen.
(Test with glowing splint.)
2I2O5 --> 2I2 + 5O2.
2. To prepare potassium iodate, put 2 g of potassium hydroxide into a test-tube and add 4 cm of water.
When dissolved, slowly add 4.5 g of iodine to the heat solution.
Pour the solution into a watch glass and let cool
6OH- + 3I2 --> IO3- + 5I- + 3H2O
Pour off the solution from the crystals, wash the latter with some water and dry them on a filter paper.
Heat crystals in a dry test-tube and show oxygen forms.
3. Prepare potassium bromate by a similar experiment using 30 drops (1 mL) of bromine in place of the iodine.
12.19.6.4 Compare silver chloride, silver bromide and silver iodide
1. Add silver nitrate solution to 2 cm of potassium chloride solution, potassium bromide solution and potassium iodide.
Ag+ + Cl- --> AgCl (s) white precipitate
Ag+ + Br- --> AgBr (s) slightly yellow precipitate
Ag+ + I- --> AgI (s) yellow precipitate
Divide each solution into two parts, Part 1 and Part 2:
Part 1. Add dilute nitric acid.
Part 2. Add drops of dilute ammonium hydroxide.
All three precipitates are insoluble in nitric acid.
Silver chloride is soluble in ammonia solution forming the soluble complex ion [Ag(NH3)2]+.
Silver bromide is slightly soluble in ammonia solution.
Silver iodide is insoluble in ammonia solution.
AgCl + 2NH3 --> [Ag(NH3)2]+ + Cl-
Silver fluoride is soluble in water, so potassium fluoride solution gives no precipitate with silver nitrate.
2. Repeat the experiment or divide the original precipitate into three parts instead of two and show that all three precipitates are soluble in sodium thiosulfate solution.
Silver chloride dissolves sodium thiosulfate to produce the sodium silver thiosulfate ion, NaAgS2O3.
AgCl + Na2SO3 --> NaAgS2O3 + NaCl.
12.19.6.13 Nitrogen triiodide, NI3
Nitrogen triiodide, ammonium triiodide, triiodoamine, touch powder, contact explosive
Explosive if dry, Nitrogen triiodide is Not permitted in schools.
12.19.7.0 Hydrofluoric acid, HF
Hydrofluoric acid is Not permitted in schools
Hydrofluoric acid, HF, is a colourless gas with a strong irritating odour.
It dissolves in water to form hydrofluoric acid.
Hydrogen fluoride, corrodes most substances except lead, wax, polyethylene, and platinum.
It comes as a, white powder and may be dyed blue for identification.
Hydrogen fluoride gas can cause irritation, muscle spasms, harm the lungs and heart and cause death.
Contact with hydrofluoric acid (even diluted) can burn the eyes (causing blindness) and skin, causing severe burns deep beneath the skin damaging internal tissues.
Australian drinking water guidelines maximum of 1.5 mg / L (i.e. 0.0015 g / L)
Hydrogen fluoride is used to make aluminium, chlorofluorocarbons (CFCs), aluminium fluoride, sodium fluoride and other fluoride salts.
It is used in the petroleum, chemical, and plastics industries.
It is used to separate uranium isotopes, clean metals, bricks, or remove sand from metal castings, etch glass and enamel, polish glass and galvanize iron, in brewing and to cloud light bulbs.
Fluorine combines with nearly all known elements.
Hydrofluoric acid must not be produced as a reaction product, e.g. by reaction of metal fluorides with mineral acids.
Hydrofluoric acid is not used in school science laboratories.
Hydrofluoric acid dissolves glass so it must be stored in polyethylene or Teflon containers.
It is a very hazardous substance.
It is used for tyre cleaning and aluminium brightener solutions.
Dissociation of hydrofluoric acid
HF + H2O <--> H3O+ + F-.
12.19.7.1 Prepare hydrogen fluoride
Hydrogen fluoride, HF, is Not permitted in schools
Use a fume cupboard.
Coat a microscope slide with wax and remove part of the wax by writing on the slide with a pin.
Pour concentrated sulfuric acid over powdered calcium fluoride in the bottom of a lead basin.
Put the microscope slide face downwards over the basin and heat gently.
Blowing across an ammonia bottle in the directions of the gas.
Put a piece of damp blue litmus paper in the gas.
The action of silver nitrate on fluorides in solution is not typical as silver fluoride is soluble in water.
Note the steam-like fumes hydrogen fluoride.
After three minutes, remove the wax from the slide and note the writing etched on the glass, because of the formation of silicon fluoride.
CaF2 + H2SO4 --> CaSO4 + 2HF
SiO2 + 4HF --> SiF4 + 2H2O
(The SiO2 comes from the glass in the microscope slide.)
12.19.7.2 Prepare silicon tetrafluoride
Silicon tetrafluoride, SiF4, is not permitted in schools
See diagram 12.19.7.2: Prepare silicon tetrafluoride
1. All the apparatus must be dry.
Mix 5 cc of fine dry sand and 5 cc of powdered calcium fluoride and put the mixture into a 250 mL flask with a two-hole stopper fitted with a funnel and delivery tube.
Put other end of the delivery tube into a gas jar, to keep the delivery tube dry.
Pour concentrated sulfuric acid on to the mixture in the flask, and shake to moisten the whole mass.
Pour water to a depth of 6 cm and then heat the contents of the flask.
The silicon tetrafluoride, which is a gas, passes into the water and hydrolyses to form hydrated silica as a white precipitate in hydrofluosilicic acid solution (H2SiF6).
3SiF4 + 2H2O --> SiO2 (s) + 2H2SiF6
The precipitate may be regarded as silicic acid with a formula H2SiO3 or H4SiO4 or as hydrated silica (SiO2.xH2O).
Some of the water may be combined and some may be occluded.
2. To obtain a specimen of pure silica from sand, filter off the solution with suspended silica obtained above, and wash the hydrated silica well with four washings of hot distilled water.
Transfer the silica to a crucible and heat to redness.
Leave to cool.
The product is pure silica.
12.19.8.1 Reactions of sodium chloride
1. Sodium chloride dissolves in water to form an aqueous solution containing the metal ion and chloride ion.
NaCl (s) + aq --> Na+ (aq) + Cl- (aq)
2. Add 1 mL of concentrated sulfuric acid to 2 cc of sodium chloride solution. Hydrogen chloride is given off.
NaCl + H2SO4 --> NaHSO4 + HCl (g)
3. Grind 2 cc of sodium chloride with twice its volume of manganese dioxide and transfer the mixture to a boiling tube.
Add 3 mL of concentrated sulfuric acid and heat.
Chlorine forms and some hydrogen chloride forms.
Chlorine bleaches wet litmus paper.
2NaCl + 2H2SO4 + MnO2 --> Na2SO4 + MnSO4 + 2H2O + Cl2 (g).
12.19.8.3 Prepare iron (III) chloride
See diagram 12.19.8.3: Prepare iron (III) chloride
Use a fume cupboard.
This experiment may not be allowed in some school systems.
1. Anhydrous iron (III) chloride and some other chloride, e.g. iron (II) chloride, magnesium chloride, zinc chloride, aluminium chloride, and antimony chloride, cannot be prepared by evaporating the salt solution to dryness, because hydrolysis occurs and the final product is iron (III) oxide or any of the other oxides.
Wind 50 cm of thin iron wire around a pencil and put the wire in a combustion tube.
Pass dry chlorine through the combustion tube over for a minute to displace the air and then heat the tube with a Bunsen burner until the iron wire commences to burn.
After removal of the Bunsen burner flame the wire will continue to burn if the supply of chlorine is sufficient.
Most of the iron (III) chloride condenses as a mass of black crystals in a cooler part of the combustion tube.
2Fe + 3Cl2 (g) --> 2FeCl3 (s)
2. Repeat the experiment using dry hydrogen chloride is used in place of chlorine.
The small colourless scales of iron (II) chloride produced are much less volatile and often stick to the iron.
Fe + 2HCl --> FeCl2 + H2 (g).
12.19.8.4 Concentrated sulfuric acid with potassium chlorate, KClO3
This experiment may not be allowed in some education systems.
Drop a crystal of potassium chlorate the size of half a small split pea into a clean dry test-tube and clamp in a nearly horizontal position.
Make sure that the mouth of the test-tube points away from you in a safe direction.
Drop two drops of concentrated sulfuric acid into the mouth of the test-tube.
Adjust the slant of the test-tube so that the acid runs slowly down onto the potassium chlorate.
The yellow gas given off is chlorine dioxide, ClO2.
Install a safety screen between you and the equipment and slowly heat the test-tube while holding it at arm's length.
A violent reaction occurs as the chlorine dioxide decomposes.
3KClO3 + 2H2SO4 --> KClO4 + 2KHSO4 + 2ClO2 + H2O.
12.19.8.5 Prepare potassium perchlorate
Prepare potassium perchlorate crystals by fractional crystallization
See diagram 12.19.8.5: Solubility of potassium salts
Half fill a crucible with potassium chlorate.
Fit the crucible firmly in a pipe clay triangle.
Heat gently until the potassium chlorate melts then stir the liquid it becomes pasty while supplying heat to keep the mass molten.
Leave to cool, add an equal volume of water and heat gently until all the potassium chlorate has dissolved.
Pour the solution on to a watch glass and let cool.
The crystals that appear are almost pure potassium perchlorate, KClO4, that can be purified further by dissolving in hot water and crystallizing again.
4KClO3 --> 3KClO4 + KCl
potassium chlorate ---> potassium perchlorate + potassium chloride.
12.19.8.6 Recover silver from silver chloride, AgCl2
1. Wash the residues with water, dry them and mix with twice the volume of a mixture of anhydrous sodium and potassium carbonates.
Transfer the mixture to a crucible and heat strongly in a furnace, then leave to cool.
Note a button of silver remaining in the bottom of the crucible.
2. Transfer the residues after washing to an evaporating basin and add sodium hydroxide solution and glucose and heat the mixture.
When a portion of the solid dissolves completely in dilute nitric acid, pour off the liquid from the grey silver, which remains in a finely divided condition.
12.19.8.7 Prepare chromyl chloride, CrO2Cl2
Heat a dry test-tube in the Bunsen burner flame to soften the glass a third of the distance from the open end.
Draw out the glass to reduce the diameter of the test-tube to a 0.5 cm at the heated part and at the same time bend the open end slightly downwards.
The apparatus will then serve as a small retort.
When cold introduce into the test-tube a mixture of not more than 2 cc of finely ground potassium dichromate and half that amount of sodium chloride.
Add just enough concentrated sulfuric acid to cover the mixture.
Grasp the test-tube in one holder and in another hold a dry test-tube to act as a receiver.
Heat the mixture gently and collect drops of the red-brown liquid, chromyl chloride.
K2Cr2O7 + 4NaCl + 3H2SO4 --> K2SO4 + 2Na2SO4 + 2CrO2Cl2 + 3H2O
Add drops of water to the compound.
Test the hydrogen chloride given off with litmus paper and with silver nitrate solution on a glass rod.
The yellow solution contains chromic acid.
Add sodium hydroxide solution until neutral, then acidify with acetic acid and add lead acetate solution.
Note the yellow precipitate that shows the presence of chromate ion.
CrO2Cl2 + 2H2O --> H2CrO4 + 2HCl.
12.19.8.8 Phosphorus pentachloride with water
Add a piece of solid phosphorus pentachloride the size of half a small pea to water.
Note the vigorous reaction.
Phosphorus pentachloride forms an acidic solution of phosphoric (V)) acid in water.
PCl5 (s) + 4H2O (l) --> H3PO4 (aq) + 5H+ (aq) + 5Cl- (aq)
PCl5 + 4H2O --> H3PO4 +5HCl (g).
12.19.9.1 Reactions of bromine, Br2
Be careful! Liquid bromine can cause sores if in contact with the skin.
Also, bromine vapour is painful to the eyes.
Store bromine water in a screw-topped bottle in a refrigerator.
1. Let drops of bromine fall into a test-tube and cover the mouth by 2/ 3 with a stopper.
The bromine evaporates and fills the test-tube.
Fix the stopper firmly in the mouth of the test-tube.
Invert a test-tube of hydrogen gas over a test-tube of bromine and let them mix.
Apply a flame and note the weak explosion.
H2 + Br2 --> 2HBr.
2. Invert a test-tube of hydrogen sulfide over a test-tube of bromine.
Note the sulfur precipitate.
Misty fumes of hydrogen bromide replace the colour of bromine.
Br2 + H2S --> 2HBr + S (s).
3. Dip a filter paper in an alcoholic solution of fluorescein and let dry.
Put it in a gas jar of bromine vapour, when the paper turns red, because of the formation of eosin.
2.19.9.6 Reactions of bromides, Br-
Use one crystal the size of a match head or 1 mL of 10% potassium bromide for each experiment, chlorine water, carbon tetrachloride.
1. Grind the bromide with a small quantity of manganese dioxide, add 1 mL of concentrated sulfuric acid to the mixture in a test-tube and heat gently.
The red vapour of bromine may condense to small drops of liquid bromine on the sides of the test-tube.
MnO2 + 2KBr + 2H2SO4 --> MnSO4 + K2SO4 + 2H2O + Br2.
2. Add drops of silver nitrate solution to potassium bromide solution. Note the pale yellow precipitate of silver bromide that is insoluble in dilute nitric acid, but dissolves in excess ammonium hydroxide, i.e. it is sparingly soluble.
Ag+ + Br- --> AgBr (s).
3. Add drops of chlorine water to a potassium bromide solution.
Bromine is liberated, which turns the solution light brown or red.
Cl2 + 2Br- --> Br2 + 2Cl-.
12.19.9.1.1 Prepare bromine water
Bromine water, < 0.1% solution, Not hazardous, but not be ingested, use eye and skin protection
Bromine water, a yellow orange solution of bromine in water, is usually available from laboratory suppliers.
Bromine gas, Br2 (g), and bromine liquid, Br2 (l), are too dangerous to prepare in schools so only bromine water may be prepared.
Neutralize bromine water residue with 10% sodium carbonate solution, and dilute with water.
Flush to foul water drain.
1. Pass chlorine gas through a bromide salt solution (Collect soluble gases in water)
Cl2 + 2Br --> 2Cl- + Br2 (aq).
2. Cool a 1 mL ampoule of bromine water, Br2 (aq) in a refrigerator, break with forceps under 200 mL water.
Be careful! Use safety glasses and nitrile chemical-resistant gloves.
3. Put a 1 mL ampoule of bromine water inside a test-tube of an appropriate size.
Fix a short piece of wide bore rubber tubing to the end of the test-tube.
Seal the other end of the rubber tubing with a wide bore glass rod.
Move the ampoule inside the test-tube so that the top is inside the rubber tubing.
Break the ampoule by squeezing the rubber tubing with pliers.
Put in water and carefully cut the rubber tubing to release the bromine.
4. Prepare bromine water by adding a small volume of bromine to a large volume of water.
The bromine dissolves to a limited extent, colouring the water yellow-orange.
Use the relatively harmless bromine water to demonstrate reactions of bromine with various organic and inorganic materials.
Reaction is shown by loss of colour from the solution.
5. Mix 5 g of potassium bromide and 5 g of manganese dioxide in a 250 mL conical flask.
Add 5 mL of concentrated sulfuric acid and seal with a cork fitted with a distillation tube.
Heat with a Bunsen burner to distil the bromine into a 200 mL conical flask containing 50 mL of water.
Collect the gas in water solution.
Make sure that the end of the distillation tube is just below the surface of the water.
Do this reaction in a fume cupboard.
6. Add hydrogen peroxide and a few drops of concentrated sulfuric acid to a dilute solution of sodium or potassium bromide.
A yellow colour develops after a few minutes.
7. Shake 0.5 mL of bromine in a fume cupboard with 100 mL of water.
Do this in the fume cupboard, wearing rubber gloves.
The solution deteriorates with time.
Store in a tightly stoppered brown bottle.
12.19.9.1.2 Opening a bromine gas ampoule
Open an ampoule of bromine gas with extreme care in a fume cupboard while wearing safety glasses.
Cool the bottom of the ampoule in an ice / water mixture to reduce the vapour pressure of the bromine before opening.
Use a glass knife (ceramic impregnated with diamond dust) or the edge of a new file to score around the outside of the neck.
Wrap the ampoule in a cloth towel before cracking the ampoule open.
Transfer the bromine to a container with a tightly fitting glass stopper or a Teflon-lined screw cap.
This container should be placed in a larger container with some packing material to ensure it cannot leak or break.
Both containers should be fully labelled.
Always store the bromine in a cool secure store area.
12.19.9.2 Reactions of bromine water (bromine solution)
Bromine dissolves slightly in water forming a 4% yellow-orange solution.
A red vapour remains above the saturated solution.
1. Add iron filings to 3 cm of bromine water and shake the mixture.
Note the pale green solution of iron (II) bromide if iron is in excess, or the yellow solution of iron (III) bromide if bromine is in excess.
Tests for the presence of iron (II) or iron (III) iron by adding sodium hydroxide solution to give a black precipitate of Fe3O4.
Fe + Br2 --> FeBr2
Fe + Br2 --> Fe2+ + 2Br-
2Fe + 3Br2 --> 2FeBr3
2Fe + 3Br2 --> 2Fe3+ + 6Br- .
2. Hold a piece of blue litmus paper in the vapour above bromine water.
The litmus paper turns red and becomes bleached.
3. Add drops of bromine water to 3 cm of sulfurous acid.
Test the solution for sulfate by adding dilute hydrochloric acid, followed by barium chloride.
SO32- + Br2 + H2O --> SO42-+ 2Br- + 2H+.
4. Add drops of sodium hydroxide solution to 3 cm of bromine water until the colour disappears.
The remaining solution contains the hypobromite ion, BrO-.
2OH- + Br2 --> BrO- + Br- + H2O
The hypobromite solution can precipitate manganese dioxide from manganese sulfate solution and precipitate lead dioxide from lead nitrate solution.
5. Add 2 cc of red phosphorus to 3 cm of bromine water.
Shake the mixture and leave to stand.
Note that the colour of the bromine water disappears.
The bromine and phosphorus combine and the resulting bromide of phosphorus decomposes to give phosphorous or phosphoric acids.
P4 + 6Br2 --> 4PBr2
P4 + 10Br2 --> 4PBr5
2PBr3 + 6H2O --> 2H3PO3 + 6HBr
PBr2 + 4H2O --> H3PO4 + 5HBr
Add bromine water to potassium iodide solution.
Iodine is displaced.
2I- + Br2 --> I2 (s) + 2Br-.
6. Prepare the following solution in two test-tubes: 2 drops of bromine water with 10 drops of n-hexane.
Shake the test-tubes and add stoppers.
Wrap the first test-tube in aluminium foil to exclude light.
Place the second test-tube next to an artificial light source.
Hexane and other saturated hydrocarbons can be brominated with the right wavelength of light shining into the test-tube or by adding hydrogen peroxide or by heating the reaction .
C6H14 + Br2 --> C6H13Br + HBr
A substitution reaction where energy from the light source is used to substitute one hydrogen atom for bromine to form bromohexane.
Also, a hydrogen atom bonds with bromine to form hydrogen bromide.
Zinc dust decolorizes bromine water.
12.19.9.3 Prepare hydrogen bromide
See diagram 12.19.9.3: Prepare hydrogen bromide
1. Add drops totalling 2 mL of acid to 1 mL of water in a boiling tube.
Add 2 cc of potassium bromide and heat gently.
Be careful!
KBr + H2SO4 --> KHSO4 + HBr (g).
2. Put a paste of 5 g of red phosphorus with water and sand into the flask.
The sand is to moderate the action.
Slowly let drops of bromine fall from the tap funnel.
The first few drops react with a flash of light.
Pass the gases through a U-tube containing beads smeared with damp red phosphorus to remove the bromine volatilized by the heat of the reaction.
Collect the hydrogen bromide by displacement of air or by passing it through an inverted funnel over water.
12.19.9.3a Heat hydrogen bromide
Show the action of heat on hydrogen bromide by filling a boiling tube with hydrogen bromide, inserting a loose stopper, and heating strongly with a Bunsen burner.
Hold a piece of white paper behind the boiling tube as soon as decomposition starts.
2HBr <--> H2 + Br2.
12.19.9.4 Reactions of hydrogen bromide, HBr
Test the misty fumes of hydrogen bromide by slowly lowering a drop of the following reagents on the end of a glass rod into the gas.
1. Silver nitrate solution: A pale yellow precipitate of silver bromide forms.
Ag + + Br- --> AgBr (s)
2. Concentrated 880 ammonia: white ammonium bromide fumes form.
NH3 + HBr --> NH4Br
3. Litmus solution: Litmus turns red.
4. Chlorine water: Yellow coloration, because of bromine.
2Br- + Cl2 --> Br2 + 2Cl-
5. A drop of water: This may be removed and tested by dipping the rod into drops of silver nitrate solution in a test-tube.
The positive reaction shows the high solubility of the gas in water.
6. Concentrated nitric acid: The hydrogen bromide is rapidly oxidized to bromine.
2HNO3 + 2HBr --> 2H2O + 2NO2 + Br2
7. Attach a stopper with a delivery tube bent at right angles and heat strongly.
Hold a piece of white paper behind the test-tube.
Strong heat decomposes the hydrogen bromide into bromine and hydrogen gas.
If the bromine is not easily visible, put in the bromine vapour a filter paper dipped in an alcoholic solution of fluorescein and let dry.
The filter paper turns red, because of the formation of eosin.
12.19.9.5 Prepare potassium bromide
1. Use 100 mL of distilled water in the flask.
Measure 5 mL of bromine in a measuring cylinder, which already contains 3 mL of water.
Pour the bromine and water into the flask (and wash out the cylinder without delay).
Weigh 8 g of iron filings and add to the solution in portions of 0-5 g, shaking well on each addition.
If this operation of adding the iron is hurried, much heat is generated and some iron (III) bromide forms and persists throughout the preparation.
Heat the flask on a water bath for ten minutes and filter quickly.
Fe + Br2 --> FeBr2.
2. Prepare potassium carbonate solution, 20 g in 50 mL of water.
Add this solution to the green iron (II) bromide solution, mix and heat on a water bath for ten minutes.
The white later green precipitate is iron (II) carbonate.
FeK2 + K2CO3 --> 2KBr + FeCO3
Filter quickly and evaporate the colourless solution to crystallization.
Observe a drop of solution under a microscope for cubic crystals of the bromide.
12.19.9.7 Bromination of acetone
The term "bromination" may simply mean to combine a substance with bromine or a bromine compound.
However, it may have a more complicated meaning in chemistry.
In a fume cupboard, a solution of acetone in water + hydrochloric acid reacts with saturated bromine water to form bromoacetone and hydrogen bromide.
The reaction goes through several stages.
Bromoacetone is a poisonous tear gas and was used as such in the First World War.
CH3COCH3 + Br2 --> CH3COCH2Br + HBr
acetone + bromine --> bromoacetone + hydrogen bromide.
12.19.9.8 Reactions of 2-bromopropane with hydroxide ions, nucleophilic reactions
Heat 2-bromopropane in a reflux condenser with concentrated solution of sodium hydroxide in ethanol.
1. Elimination reaction
The hydroxide ion acts as a base to remove hydrogen as a hydrogen ion from the first carbon atom.
This reaction expels the bromine as a bromide ion to form propene, leaving sodium ions, bromide ions and water.
CH3CHBrCH3 + NaOH --> CH2=CHCH3 + NaBr + H2O
2-bromopropane + sodium hydroxide --> propene + sodium bromide + water.
2. Nucleophilic substitution reaction
The nucleophilic -OH group replaces the bromine to form the alcohol propan-2-ol, leaving sodium ions and bromide ions.
CH3CHBrCH3 + NaOH --> CH3CHOHCH3 + NaBr
2-bromopropane + sodium hydroxide --> propan-2-ol + sodium bromide
This example of nucleophilic substitution reaction is the hydrolysis of an alkyl bromide, R-Br, under alkaline conditions, where the attacking nucleophile is the OH-, leaving the group Br-.
R-Br + OH- --> R-OH + Br-
The nucleophile OH- donates an electron pair to an electrophile, RBr.
Nucleophiles are electron rich so they find electron deficient sites in an organic molecule, e.g. H2O, NH3, OH-, Cl-, Br-, CN-.
The organic chemistry term nucleophile (nucleophilic reagent) forms a bond to its reaction partner (the electrophile) by donating an electron pair to it to form a chemical bond in the reaction.
12.4.1.0 Prepare chlorine
See diagram 1.13a: Simple fume hood
Be careful! Prepare chlorine gas only in a fume hood or fume cupboard.
Chlorine gas is poisonous and damages the respiratory organs.
Do not prepare chlorine in an open room.
Use small quantities only.
See diagram 1.13: Smelling a gas
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.
Chlorine
Chlorine is a greenish yellow gas with an irritating and choking odour and can be poured from one container to another.
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.
However, the best treatment for inhaling chlorine gas is plenty of fresh air.
12.4.1.1 Prepare chlorine with bleaching powder or bleach solution.
Bleaching powder is a mixture of calcium chloride, calcium hydroxide and calcium chlorate (I).
Bleaching powder is manufactured by the reaction of chlorine with solid calcium hydroxide.
See diagram 1.13a: Simple fume hood
Be careful! Prepare chlorine gas only in a fume hood or fume cupboard.
Do not prepare chlorine in an open room.
Use small quantities only.
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. Solutions of bleach (sodium hypochlorite), and solid bleaching powder produce small amounts of chlorine gas when exposed to the air, with its characteristic smell.
Adding acid causes a vigorous production of chlorine gas.
The most convenient ways to prepare chlorine gas are to use the reaction of dilute acid with bleaching powder or the reaction of potassium permanganate with concentrated hydrochloric acid.
Collect the chlorine gas by displacement of air.
3. 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.
4. Add dilute sulfuric acid to bleaching powder.
After collecting a small amount of chlorine gas put a stopper in the receiving test-tube and put the end of the delivery tube into sodium thiosulfate solution to absorb excess chlorine.
CaOCl2 + H2SO4 (aq) --> CaSO4 (s) + H2O (l) + Cl2 (g)
Repeat the experiment with dilute hydrochloric acid.
CaOCl2 + 2HCl (aq) --> CaCl2 (s) + H2O (l) + Cl2 (g).
5. 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).
6. Prepare chlorine with bleaching fluid.
Do this experiment near an open window or in a fume hood.
Add drops of an acid solution, e.g. citric acid, tartaric acid, dilute sulfuric acid, to a few drops of bleaching fluid in a test-tube.
The green gas chlorine forms without any heating.
Sniff the gas very cautiously and test it with a piece of wet litmus paper.
The litmus paper is bleached by the chlorine.
12.4.1.2 Prepare chlorine with hydrochloric acid / potassium manganate (VII)
See diagram 1.13a: Simple fume hood
Be careful! Prepare chlorine gas only in a fume hood or fume cupboard.
Do not prepare chlorine in an open room.
Use small quantities only.
Experiments
1. Add 5 M hydrochloric acid to sodium hypochlorite.
Keep a 5M alkali solution nearby to stop the reaction.
2. 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).
3. Prepare chlorine, concentrated hydrochloric acidwith potassium manganate (VII)
See diagram 12.19.8.2: Prepare chlorine
Add drops of concentrated hydrochloric acid to a few crystals of potassium permanganate in a flask that can be stoppered.
Collect the gas by upward displacement of air in a fume hood or fume cupboard.
16HCl (aq) + 2KMnO4 (s) --> 5Cl2 (g) + 2MnCl2 (aq) + 8H2O (l) + 2KCl (aq)
6HCl + 2KMnO4 + 2H+ --> 3Cl2 + 2MnO2 + 4H2O + 2K+.
4. Put 5 cc of potassium permanganate into a flask.
Fill the dropping funnel with concentrated hydrochloric acid and allow the acid to run on to the permanganate to produce chlorine.
16H + + 2MnO4- + 10Cl- --> 2Mn2+ + 8H2O + 5Cl2 (g)
Pass the gas through water to remove hydrogen chloride and pass the gas through concentrated sulfuric acid to dry it.
Collect the gas by downward displacement of air.
5. 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) --> 5Cl2 (g) + 2MnCl2 (aq) + 8H2O (l) +2KCl (aq).
12.4.1.3 Prepare chlorine with sodium chloride
See diagram 1.13a: Simple fume hood
Be careful! Prepare chlorine gas only in a fume hood or fume cupboard.
If using small quantities only, do this experiment next to an open window or in a fume hood.
Add 2 g of black manganese dioxide to 2 g of sodium chloride, powdered alum (hydrated potassium aluminium sulfate), iron (II) sulfate, and sodium hydrogen sulfate (sodium bisulfate).
Heat the mixtures over a Bunsen burner flame while holding damp litmus paper in the mouths of the test-tubes.
Chlorine forms in the test-tube as green gas with a choking smell.
Sniff the gas very cautiously and test it with the damp litmus paper.
The litmus paper is bleached white, the usual test for chlorine.
12.4.2 Prepare chlorine water
See diagram 1.13a: Simple fume hood
Be careful! Prepare chlorine gas only in a fume hood or fume cupboard.
Use small quantities only.
Chlorine water, chlorine solution, Toxic by all routes, forms lung irritant chlorine gas (Use small quantities, < 10 mL or 10 g in well-ventilated area).
Chlorine is available from chemical suppliers 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).
3.40.1 Tests for chlorine
See diagram 1.13a: Simple fume hood
Be careful! Prepare chlorine gas only in a fume hood or fume cupboard.
Use small quantities only.
1. Bleaching test for chlorine
Chlorine bleaches moist red or blue litmus paper, flowers and some dyes in cloth.
Use chlorine gas to bleach flower petals, leaves and hair suspended in the gas.
2. Lighted splint test for chlorine
Chlorine extinguishes a lighted splint, but hot steel wool burns in it.
3. Pass chlorine through water
Chlorine is available from chemical suppliers 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).
12.4.4 Bleaching powder
Bleaching powder is a mixture of calcium chloride, calcium hydroxide and calcium chlorate (I).
Bleaching powder is manufactured by the reactions of chlorine with solid calcium hydroxide.
Domestic bleach is manufactured by mixing chlorine solution with sodium hydroxide solution.
A bleach is a chemical used to make white or remove the colouring from fibres, textiles and hair ("peroxide blondes"), e.g. hydrogen peroxide, sodium hypochlorite.
Bleaching powder from chemical suppliers is probably a mixture of calcium chlorate (I), calcium chloride and calcium hydroxide.
The value of bleaching powder depends on its available chlorine.
Bleaching may occur by oxidation using free chlorine or combined chlorine or by exposure to sunlight, e.g. "fading" of the curtains.
Bleaching is important in the textile and paper industries.
Domestic bleach is manufactured by mixing a solution of chlorine with sodium hydroxide solution.
Hypochlorous acid, HClO, "bleach", is used as a bleach and a disinfectant.
Cl2 (g) + 2OH- (aq) --> Cl- (aq) + ClO- (aq) + H2O.
Experiments
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.
2. Test with wet red or blue litmus paper that becomes colourless, because of the bleaching action of chlorine.
12.4.5 Chlorine dioxide, ClO2
Chlorine dioxide, chlorine peroxide; chloroperoxyl, chlorine (IV) oxide, is a strongly oxidizing, yellow to reddish-yellow gas or liquid with a pungent, sharp odour like chlorine and nitric acid.
It is soluble in water, alkaline, and sulfuric acid solutions.
It is usually a < 10% solution in cold water.
It is used as an industrial bleach, cleaning and de-tanning of leather, bactericide, fungicide and algicide for drinking water systems, milking equipment, food processing and handling as an alternative to chlorine because of lesser problems with disinfection by-products.
Chlorine dioxide gas is flammable, and is violently explosive in air at concentrations > 10%.
It can be ignited by sunlight, heat, or sparks.
Chlorine dioxide is strongly oxidizing, and reacts violently with organic chemicals and can be detonated by sunlight, heat, or contact with mercury or carbon monoxide.
12.4.6.0 Reactions of chlorine with benzene
Chlorine reacts with benzene, C6H6, to form three liquid aromatic isomers by chlorinating benzene with an iron filings catalyst, then separate isomers by fractional distillation.
The 3 separate compounds, isomers, called dichlorobenzene are 1,2-dichlorobenzene (ortho-), 1,3-dichlorobenzene (meta-), and 1,4-dichlorobenzene (para-).
* 1,2-dichlorobenzene, C6H4Cl2, (ortho-dichlorobenzene), BP 179oC, used in insecticides and to make dyes, chloroben (used for termite treatment), o-dichlorobenzene, Toxic, avoid inhalation of vapour, Solution < 5% Not hazardous.
* 1,3-dichlorobenzene, C6H4Cl2, colourless liquid, insoluble in water, toxic to aquatic life
* 1,4-dichlorobenzene, C6H4Cl2, (para-dichlorobenzene), BP 174oC, colourless solid, strong odour, used as a deodorant and an insecticide substitute for naphthalene in the common household "mothballs" or "moth crystals".
C6H4Cl2, p-dichlorobenzene, paradichlorobenzene, PDB, p-DCB, 1,4-dichlorobenzol, "dichlorobenzene", Toxic if ingested, may be carcinogenic
Solution < 25%, Not hazardous
May react with sulfide groups to form the polymer poly(p-phenylene sulfide) and sodium chloride
C6H4Cl2 + Na2S --> 1/n (C6H4S)n + 2NaCl.
The two isomers above can react to form (meta-dichlorobenzene) 1,3-dichlorobenzene
C6H6 (l) + Cl2 (g) --> C2H4Cl2 (l)
benzene + chlorine --> dichlorbenzene.
12.4.6.1 Halogenation of benzene
Halogenation is the substitution of -H by -Cl or -Br
Benzene reacts with bromine or chlorine in the presence of aluminium bromide/ aluminium chloride or iron, as electrophilic substitution reactions.
Iron takes part in the reaction so it is actually the bromide or chloride formed that acts as the catalyst, in the same way as does aluminium chloride.
2Fe + 3Br2 --> 2FeBr2
2Fe + 3Cl2 --> 2FeCl3
Chlorinating benzene with aluminium chloride catalyst, AlCl3.
Bubble chlorine into a mixture of benzene and anhydrous aluminium chloride.
C6H6 (l) + Cl2 (g) --> C6H5Cl (l) + HCl (g)
With iron catalyst
C6H6 + Br2 --> C6H5Br + HBr.
12.4.7 Reactions of chorine with sodium
See diagram 13.4.7: Reactions of chlorine with sodium
Be Careful! The reaction is very vigorous. Do this experiment in a fume cupboard.
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 leave to stand.
When the reaction stops, remove the spoon, allow it to cool and place it in a small amount of alcohol.
Let excess chlorine diffuse away in the fume cupboard.
Let the mixture of alcohol and solid stand until no further reaction takes place.
Wash the crystals with alcohol and let them cool and dry.
sodium (s) + chlorine (g) --> sodium chloride (s).
12.4.7.1 Reactions of chlorine with group 3 elements, Na, Mg, Al, Si, P
Use a fume cupboard.
Heat the element with a Bunsen burner plunge it into a gas jar of chlorine.
Collect the metal chlorides as solids at room temperature, but cool the non-metal chlorides.
From the equations below, note how the proportion of chlorine in the compounds increases
2Na (s) + Cl2 (g) --> 2NaCl (s)
Mg (s) + Cl2 (g) --> MgCl2
2Al (s) + 3Cl2 (g) --> 2AlCl3 (s)
Si (s) + 2 Cl2 (g) --> SiCl4 (l)
2P (s) + 5 Cl2 --> 2PCl5 (s).
12.4.8.1 Burn steel wool in chlorine
Ignite steel wool held by tongs in a Bunsen burner flame or butane lighter, then put into chlorine gas.
Be Careful!
The reaction occurs with strong combustion to form a brown-red cloud that condenses to black flakes of anhydrous iron (III) chloride.
2Fe (s) + 3Cl2 (g) --> 2FeCl3 (s).
12.4.8.2 Burn copper in chlorine
Heat copper foil in a burner flame and put into chlorine gas.
The reaction forms a layer of brown copper (II) chloride that turns green in the presence of moisture.
Cu (s) + Cl2 (g) --> CuCl2 (s).
12.4.8.3 Burn magnesium in chlorine
A burning magnesium ribbon burns violently in chlorine gas to form magnesium chloride.
Mg (s) + Cl2 (g) --> MgCl2 (s).
12.4.8.4 Burn paraffin wax tapers in chlorine
Reaction of chlorine with non-metals.
Chlorine has such a strong attraction for hydrogen that it removes all the hydrogen from the hydrocarbon paraffin leaving behind the carbon as a residue.
A mixture of chlorine and hydrogen gas does not react in the dark, but if heated or exposed to strong sunlight the mixture reacts explosively to form hydrogen chloride.
Put a burning paraffin wax taper plunged in a jar of chlorine.
It burns with a small dull red flame and produces clouds of black carbon and white fumes of hydrochloric acid.
The chlorine removes the hydrogen from the hydrocarbon to form hydrogen chloride and leaving the carbon.
The chlorine does not combine directly with the carbon.
Be Careful! Do not mix chlorine and hydrogen gas.
H2 (g) + Cl2 (g) --> 2HCl (g) + energy.
12.4.8.5 Burn candles in chlorine
Burn a small birthday candle in chlorine.
The taper keeps burning with a dull red flame and forms a thick smoke black carbon particles (soot) and hydrogen chloride gas.
nCl2 + (CH3-CH2-CH2-CH2-) --> nHCl + nC
12.4.8.6 Burn wood splints in chlorine
A burning wood taper continues to burn in chlorine gas, forming hydrogen chloride and organic chloro compounds.
12.4.8.7 Burn ethylene in chlorine
A mixture of 1 volume of ethylene and 2 volumes of chlorine, when ignited with a taper, burns with a red flame to form hydrogen chloride and black clouds of carbon
C2H4 + 2Cl2 --> 4HCl + 2C
Ethylene combines with chlorine on exposure to light to form an oily liquid.
12.4.8.8 Burn methane in chlorine
A mixture of 1 vol.
of methane and of 2 vols.
of chlorine, prepared out of direct sunlight, when ignited with a taper, burns with a weak whistling noise, to form hydrochloric acid and black clouds of carbon.
CH4 + 2Cl2 --> C + 4HCl.
12.4.912.4.9 Pass chlorine through water
Chlorine is available from chemical suppliers for school laboratory use as chlorine water.
Hypochlorous acid HClO, a bleach and a disinfectant, is an aqueous solution of chlorine (I) oxide that forms salts called hypochlorites.
Hypochlorous acid is a weak acid that easily decomposes back to chlorine gas and water.
When chlorine passes through water, a mixture of HCl and HClO forms.
The chlorine is oxidized and reduced.
Cl2 (g) + H2O (l) <--> HCl (aq) + HClO (aq).
12.4.12 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).
12.4.13 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).
12.4.14 Reactions of chlorine with copper
Be Careful! Do not breath this poisonous gas.
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).
12.4.15 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).
12.4.16 Reactions of chlorine with steel wool
Be Careful! Do not breath this poisonous gas.
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).
12.4.17 Halogen compounds
Halogen compounds: haloalkanes (alkyl halides), halogen derivatives (organic chemistry)
| See diagram 16.2.2: Chlorinated hydrocarbons, methyl chloride, methylene chloride, chloroform, carbon tetrachloride
| See diagram: 16.13.5: Halogen compounds, bifenox, dicofol, naled, trichlorophon, tetrachlorvinphos
| See diagram 16.13.11: MCPA, (2,4-D), (2,4,5-T), picloram
Acyl halides (acid halides; RCOX, where X = halide group and R = organic group, e.g. acetyl = CH3CO-
Haloforms, e.g. trihalomethanes CHX3
Alkanes react with chlorine and bromine in ultraviolet light to produce haloalkanes, e.g. 2-chloropropane.
Reactions of chlorine or bromine in ultraviolet light
CH4 + Br2 --> CH3Br + HBr
CH3Br + Br2 -->CH2Br2 + HBr
CH2Br2 + Br2 --> CHBr3 + HBr
CHBr3 + Br2 -->CBr4 + HBr.
12.19.2.5 Properties of halogens
Be careful! Chlorine and bromine are harmful when inhaled or when they contact the skin.
Fluorine, F, is a yellow gas.
Chlorine, Cl, is a green gas.
Bromine, Br, is a red-brown liquid.
Iodine, I, is a grey-black crystal.
All halogens are slightly soluble in water to form weak acidic solutions that are bleaches.
Chlorine is the most soluble and the strongest bleach.
Test each solution with universal indicator.
Universal indicator turns red then is bleached.
1. Compare the colours and states of the halogen elements at room temperature.
At room temperatures fluorine is a pale yellow gas, chlorine is a green-yellow gas, bromine is a red-brown liquid that gives a brown vapour and iodine is a grey-black solid.
2. Compare the colours, states and solubility in water of sodium fluoride, chloride, bromide and iodide.
They are all soluble, white, crystalline solids.
3. Compare the activity of fluorine, chlorine, bromine and iodine by investigating, which will displace another from their compounds.
Prepare, sodium fluoride solution, sodium chloride solution, sodium bromide solution and sodium iodide solution.
Add chlorine solution to each solution.
Chlorine has no visible effect on sodium fluoride solution or sodium chloride solution.
Chlorine turns sodium bromide solution brown-yellow.
Chlorine turns sodium iodide solution deep brown.
So chlorine displaces bromine from sodium bromide and chlorine displaces iodine from sodium iodide.
However, chlorine does not displace fluorine from sodium fluoride.
So chlorine is more active than bromine and iodine, but chlorine is less active than fluorine.
4. Compare the colours and solubility of silver fluoride, chloride, bromide and iodide by adding silver nitrate solution to sodium fluoride solution, sodium chloride solution, sodium bromide solution
and sodium iodide solution.
Silver fluoride is soluble.
Silver chloride, silver bromide and silver iodide are insoluble.
Silver chloride is white.
Silver bromide is very pale yellow.
Silver iodide is a deep yellow.
Bromine, Br
Bromine experiments: 12.19.6
Bromine, Br, (Greek bromos stench), bromide ion Br-, gas Br2, is a red-brown, fuming, volatile, poisonous, non-metal liquid between 19oC and 27oC, suffocation odour, vapour irritates eyes and throat, strong oxidizing agent, used for many chemical compounds including "anti-knock" petrol additive.
Silver bromide was important for photography.
Bromoform, tribromomethane, CHBr3 used to separate minerals.
Bromothymol blue indicator, pH 6.0 to 7.6.
Potassium bromide, was formerly used as sedative and was said to be put in army tea to lessen soldiers' sexual urges.
Bromochlorodifluoromethane, CHBrClF2, low toxic fire extinguisher for confined spaces.
The fat soluble fire retardant PBDE, polybromyldiphenyl ether, in the deca, octa and penta forms, has been detected in mothers' milk, fish and the environment, (Poison COR 1744).
Br2 (3.6% bromine), RD 3.12 gm cm-3, bp 58.7oC, solidifies -7oC (swimming pool sanitation, products from bromine, e.g. BCDMH).
Bromine is a dense red-brown liquid with a powerfully irritant vapour.
Bromine water, a yellow-orange solution of bromine in water, is usually available from chemical suppliers.
Handle pure liquid bromine in small quantities in a fume cupboard.
The liquid is unexpectedly dense, so increasing the chance of containers being dropped by inexperienced people.
Breakage of a bottle of bromine outside a fume cupboard will require evacuation of the area until the vapour dissipates.
Bromine reacts violently with active metals, e.g. aluminium / magnesium and sodium.
Do not allow active metals to contact liquid bromine.
Always store the bromine in a cool secure store area.
Atomic number: 35, Relative atomic mass: 79.904, RD 3.12, MP = -7.2oC, BP = 58.78oC.
Specific heat capacity: 448 J kg-1 K-1.

Chlorine, Cl
Reactions of chlorine: 12.19.8
Chlorine, Cl, chlorine gas Cl2, Chloride (chloro), Cl-, monodentate ligand, (Greek khlōros green), Cl2, dichlorine, molecular chlorine
Chlorine, Cl2, Highly toxic, granular, liquid, powder, tablets, highly irritant to lungs, chloride Cl-, chloro -Cl, Solution < 3%, Highly toxic by all routes
Chlorine gas, < 3%, Not Hazardous if small volume in cross ventilation
For the reactions of chlorine with metals, solid non-metals, and hydrocarbons, use small quantities only.
Chlorine is a green-yellow, dense diatomic gas, soluble in water, alcohols, and alkalis, evaporates into the air very quickly.
Chlorine is a powerful oxidizing agent.
Chlorine can react to cause fires or explosions upon contact with turpentine, ether, ammonia gas, illuminating gas, hydrocarbon, hydrogen gas and powdered metals.
Chlorine dissolves readily in water forming highly corrosive solutions.
Chlorine causes rapid corrosion of metals and destruction of plastics.
Chlorine directly combines with hydrogen gas in bright light or ignition of the mixture by lighted taper or electric spark.
Chlorine is a very reactive non-metal and free chlorine never occurs naturally.
Chlorine occurs in the minerals halite, sylvite, and carnallite and in 1.9% of sea water as chloride ions.
Chlorine is a poisonous, irritating smell gas at room temperature and pressure.
Chlorine is a powerful lung irritant, causing coughing and diminishing lung efficiency.
It attacks the mucous membrane linings of the eyes, nose, throat and lungs, causes the lungs to fill with fluid and the victim drowns.
Do not prepare chlorine in an open room, but use a fume cupboard for all reactions that may result in evolution of chlorine.
Sodium chloride, a constituent of gastric juice, which is about 0.03 M HCl.
Adults require a daily minimum, of 750 mg of chloride.
Chlorine gas was discovered in Sweden by Carl Wilhelm Scheele, 1742-1786, and identified as an element by Humphry Davy in 1810.
It was prepared by action of hydrochloric acid on manganese dioxide, "black oxide of manganese".
Chlorine is prepared by chemical suppliers by electrolysis of concentrated sodium chlorine solution (brine).
Chlorine is used in a wide range of disinfecting and cleaning products, bleaching powder and bleaches of wood pulp, paper pulp, and shrink- proofing of wool.
Chlorine kills most living things and is used to sterilize drinking water and disinfect swimming pools.
Chlorine was used as a poison gas in the First World War, as a chemical weapon.
Chlorine is used to manufacture PVC plastic and DDT insecticide.
DDT New IUPAC name: 1,1'-(2,2,2-Trichloroethane-1,1-diyl)bis(4-chlorobenzene).
Chlorine is used to prepare organic compounds, but many of these substances cannot be broken down in the environment (biodegraded), so avoid using them.
Chlorofluorocarbon refrigerant, is an aerosol now phased out, because of damage to the ozone layer in atmosphere.
Chlorophenol red pH 4.8 to 6.4 indicator,
Chloral hydrate sedative,
Chloric (V) acid, HClO3 and its salts chlorates (V) powerful oxidizing agents, plastics
Chlorinated lime is used for water purification, flame retardant compounds and batteries, metal fluxes, detinning and dezincing iron swimming pool algicide.
Tetrachloroethene CCl2.CCl2, solvent
Tetrachloromethane (carbon tetrachloride), CCl4, solvent
Trichloromethane CHCl3, chloroform, anaesthetic
1,1,1-trichloroethane CH3CCH3, safer solvent
Atomic number: 17, Relative atomic mass: 35.453, RD 1.56 (238 K), MP = -101oC, BP = -34.7oC.
Specific heat capacity: 477 J kg-1 K-1.