HALOGENS

The halogens are the group VII elements. The family comprises of fluorine, chlorine, bromine, iodine and astatine. Each member of the family contains 7 electrons at the valence shell. Halogens are the most reactive non-metals; hence, they don’t occur freely in nature but combine with metals to form salts.

Fluorine is the most reactive element in the family and is in many ways unique. Bromine and iodine are chemically very similar to chlorine. Astatine is a radioactive element and has been made artificially. It does not occur naturally.

 

Electronic configuration and properties of halogen

The electronic configuration of halogens is one electron short of the noble gas structure. Each atom of the elements in the group complete its octet structure by either:;

– Gaining one electron usually from group I and II to form a univalent negative ion, e.g. F-, Cl-, Br-.

Or

– Sharing a pair of electrons in a single covalent with another atom having a fairly similar electronegative value as in gaseous chlorine; Cl – Cl, and hydrogen chloride; H – Cl.

The electronic configurations of halogens are:

9F – 2, 7

17Cl – 2, 8, 7

35Br – 2, 8,18,7

53I – 2,8,18,18,7

85At – 2, 8, 18, 32,18,7

 

General properties of halogens

1. They are all diatomic molecules; hence, they have low melting and boiling points and are soluble in covalent non-polar liquids.
2. At room temperature, fluorine and chlorine are gases, bromine is liquid and iodine is solid. The boiling point increases down the group.
3. They are coloured with typical penetrating odours. The colours deepen down the group. Fluorine is pale yellow, chlorine is greenish-yellow, bromine is reddish-brown and iodine is purple black.
4. Their reactivity decreases down the group. The reactivity series such as volatility, solubility, electronegativity, oxidizing ability, and tendency to enter into chemical reactions decrease down the group. This gradation in properties is due to the increasing complexity of the atoms as their atomic numbers increase from fluorine to iodine.
5. All halogens dissolve to some extent in water. Fluorine reacts vigorously with water to produce oxygen and hydrogen fluoride.
6. They are all strong oxidizing agents. Fluorine oxidizes water to peroxide. Chlorine reacts slowly with water to give oxochlorate (I) acid, (HClO). The oxochlorate (I) acid decomposes to give hydrochloric acid and oxygen.

2HClO (aq)     →         2HCl(aq) + O2(g)

7. They all combine readily with most metals (except gold, platinum and titanium) and non-metals (except oxygen, carbon and nitrogen) to form ionic and covalent compounds, respectively. These compounds contain halide ion, X- in which they have an oxidation state of -1. E.g.

2Na(s) + X2   →          2NaX(s)

2P(s) + 3X2  →         2PX3(s)

 

Note: Under certain conditions, the less reactive halogens react more vigorously than chlorine. For example, bromine is capable of corroding gold because it is a liquid and has more concentration than chlorine which is a gas.

 

8. They combine with hydrogen to form gaseous covalent compounds of formula HX. The halide ion has an oxidation state of -1.

9. They react with silver to form silver halide. Only silver fluoride is soluble. Silver chloride is white, silver bromide is pale yellow and silver iodide is bright yellow. This characteristic can be used to test for halides.

10. They all react with sodium hydroxide to give a mixture of halide salt and oxochlorate (I) salts e.g.

Cl2(g) + 2NaOH(aq)  →  NaCl(aq) + NaClO(aq) + H2O(l)

11. More reactive halogen displaces (displacement-reaction) less reactive halogens from their salts. This is a typical redox reaction e.g. Br2(aq) + 2KI (aq)              2KBr(aq) + I2(aq)

 

CHLORINE AND ITS COMPOUNDS

Chlorine is the most important halogen. It does not occur freely in nature but in a combined state as chlorides, mostly sodium chloride.

 

Laboratory preparation of chlorine

Chlorine is usually prepared by the oxidation of concentrated hydrochloric acid with a strong oxidizing agent such as manganese (IV) oxide – MnO2, potassium tetraoxomanganate (VII)-KMnO4 or lead (IV) Oxide – PbO2.

– Using MnO2: heat a mixture of MnO2 and conc. HCl . the acid is oxidize to liberate chlorine.

MnO2(s) + 4HCl(aq) heat     MnCl2(aq) + 2H2O(l) + Cl2(g)

– Using KMnO4 – heat is not required. The conc. HCl is oxidized upon addition of KMnO4.

2KMnO4(aq) + 16HCl(aq)               2MnCl2(aq) + 2KCl(aq) + 8H2O(l) + 5Cl2(g)

 

 

 

Industrial preparation of chlorine

Chlorine is manufactured industrially by the electrolysis of (a) brine (ii) the molten chlorides of sodium, magnesium or calcium. The chlorine is then liquefied and stored under pressure in steel cylinder.

Physical properties of chlorine

1. It is a greenish – yellow gas with an unpleasant choking smell.
2. It is moderately soluble in water
3. It is about 2.5 times denser than air
4. It can easily be liquefied under a pressure of about 6atm.
5. It is poisonous

 

Chemical properties of chlorine

– Displacement of other halogens (except fluorine) – chlorine displaces other halogens below it in

the Periodic Table from their respective acids and salts e.g

Cl2(g) + 2NaBr(aq)            2NaCl(aq) + Br2(l)

Cl2(g) + 2HI(aq)             2HCl(aq) + I2(s)

– Direct combination with other elements:

(i) Metals react readily with chlorine especially when heated to form the corresponding chloride e.g.

2Fe(s) + 3Cl2(g)             2FeCl3(s)

Note: where a metal forms more than one chloride, the higher chloride is usually formed due to the oxidative nature of chlorine. For example, Iron(III) chloride and not iron (II) chloride is formed in the reaction above.

(ii) All non-metals (except the noble gases, oxygen, nitrogen and carbon) burn in chlorine to produce the corresponding chlorides e.g. hydrogen and chlorine combine explosively in bright sunlight to produce hydrogen chloride. The reaction is slower in diffuse light.

H2(g) + Cl2(g)    →       2HCl(g)

– Reaction with hydrogen – chlorine reacts readily with molecular hydrogen as well as compounds containing hydrogen to form hydrogen chloride e.g.

(i) With hydrocarbons, chlorine removes hydrogen from burning hydrocarbons like wax, candle, petrol or methane to form hydrogen chloride fumes and soot. Also, a filter paper saturated with turpentine (C10H16) will rapidly catch fire in a jet of chlorine to form soot and misty fumes of hydrogen chloride gas.

C10H16(l) + 8Cl2(g)     →        10C(s) + 16HCl(g)

Under controlled conditions, substitution reaction occurs. In this reaction, chlorine reacts with saturated hydrocarbons to produce chlorinated hydrocarbons and hydrogen chloride e.g.

CH4(g) + Cl2(g)   →           CH3Cl(g) + HCl(g)

                                     Chloromethane

CH3Cl(g) + Cl2(g)     →        CH2Cl2(g) + HCl(g)

          Dichloromethane

CH2Cl2(g) + Cl2(g)     →       CHCl3(g) + HCl(g)

                                      Trichloromethane

CHCl3(g) + Cl2(g)      →        CCl4(g) + HCl(g)

                                       Tetrachloromethane

(ii) with ammonia – chlorine reacts with ammonia to form nitrogen and hydrogen chloride. The hydrogen chloride then reacts with excess ammonia to form ammonium chloride.

2NH3(g) + 3Cl2(g)      →        N2(g) + 6HCl(g)

HCl(g) + NH3(g)   →         NH4Cl(s)

(iii) with hydrogen sulphide – when hydrogen sulphide and chlorine are mixed, a yellow deposit of sulphur is formed as a result of oxidation of sulphide by chlorine.

H2S(g) + Cl2(g)     →       2HCl(g) + S(l)

(iv) With water–chlorine water (formed by bubbling chlorine through water) gives off oxygen when exposed to sunlight. This is due to the formation of oxochlorate (I) acid, which decomposes slowly to yield hydrogen chloride and oxygen.

Cl2(g) + H2O(l)      →        HCl(g) + HOCl(aq)

2HOCl(aq) sunlight →   2HCl(aq) + O2(g)

 

– As an oxidizing agent, –  chlorine is a powerful oxidizing agent because of: (i) its ability to remove

hydrogen (ii) its readiness to accept electrons from reducing agents to from chloride ions e.g. (i) 2FeCl2(aq) + Cl2(g)     →          2FeCl3(aq)

     Green yellow

(ii) when bubbled through a freshly prepared solution of H2SO3, chlorine oxidizes the trioxosulphate (IV) ion to tetraoxosulphate (VI) ion.

H2SO3(aq) + H2O(l) + Cl2(g)   →         H2SO4(aq) + 2HCl(aq)

– As a bleaching agent – in the presence of water, chlorine bleaches most dyes and inks except those containing carbon like printer’s ink. The bleaching action of chlorine is due to its ability to react with water to form oxochlorate (I) acid. The oxochlorate (I) acid is unstable and decomposes to release oxygen, which oxidizes the dye to form a colourless compound.

HOCl(aq)       →        HCl(aq) + [O]

Dye + [O] (dye + O)

Coloured colourless

 

– Reaction with alkali

(i) with dilute alkali – when chlorine is bubbled through a cold dilute solution of sodium hydroxide, a pale – yellowish mixture of the oxochlorate (I) and chloride of the metal is formed

e.g. Cl2(g) + 2NaOH(aq) NaOCl(aq) + NaCl(aq) + H2O(l)

Sodium oxochlorate (I)

Bleaching powder is prepared by bubbling chlorine through a freshly prepared solution of slaked lime; Cl2(g) + Ca(OH)2(s) CaOCl2. H2O(s)

Slaked lime bleaching powder

(ii) with conc. Alkali solution – if chlorine is bubbled through a hot conc. Solution of sodium hydroxide, a mixture of trioxochlorate (V) and chloride of the metal is formed. 3Cl2(g) + 6NaOH(aq) NaClO3(aq) + 5NaCl(aq)

Sodium trioxochlorate (V)

Test for chlorine

1. Drop a piece of damp blue litmus paper into the gas jar of the unknown gas. If the litmus paper turns pink and then becomes bleached, the gas is chlorine.
2. Drop a piece of damp starch-iodide paper into the unknown gas jar. If the paper turns dark blue, then the gas is chlorine.

Uses of chlorine

– As a powerful germicide, it is used in the sterilization of water

– It is used as a bleaching agent for cotton, linen and wood-pulp. It is too strong for bleaching animal fibre.

– It is used in the manufacture of;

Important organic solvents like CHCl3, CCl4 etc

Polyvinyl chloride (PVC) a widely used plastic

Hydrochloric acid

KClO3 used for making matches and in fireworks and NaClO3 which is a weed killer.

Bleaching powder and NaClO used in dye works and laundry.

Domestic antiseptic e.g. acidified NaClO

 

 

HYDROGEN CHLORIDE

Hydrogen chloride is a hydride of chlorine. It exists as a gas at s.t.p. and dissolves in water to form

hydrochloric acid. It was first prepared by priestly in 1772.

Laboratory preparation of hydrogen chloride

Hydrogen chloride is prepared in the laboratory by the action of concentrated tetraoxosulphate (VI) acid

on sodium chloride. The gas is dried by passing it through concentrated H2SO4 and is collected by downward delivery.

2NaCl(s) + H2SO4(aq)                 Na2SO4 + 2HCl(g)

 

Industrial preparation of hydrogen chloride

It is prepared industrially be the direct combination of hydrogen and chlorine gases obtained from the electrolysis of brine.

H2(g) + Cl2(g)                  2HCl(g)

Physical properties of hydrogen chloride

– It is colourless with a sharp irritating smell

– It is slightly denser than air

– It turns moist blue litmus paper red (property of an acidic oxide)

– It is very soluble in water forming an aqueous solution of hydrochloric acid

– It fumes strongly in moist air forming droplets of hydrochloric acid

– It does not burn or support combustion

– Dry hydrogen chloride dissolves readily in non-polar solvents like chloroform and toluene, but does not ionize as it does in polar solvents.

Chemical properties of hydrogen chloride

1. Reaction with ammonia – hydrogen chloride gas reacts with ammonia gas to form dense white fumes of solid ammonium chloride.

HCl(g) + NH3(g)    →         NH4Cl(s)

 

2. Reaction with metals – hydrogen chloride gas on heating reacts with many metals, forming the corresponding chlorides and hydrogen. If the metal can form two chlorides, the lower chloride is formed.

Zn(s) + 2HCl(g)     →        ZnCl2(s) + H2(g)

Fe(s) + 2HCl(g)    →         FeCl2 (s) + H2(g)

 

Assignment

Describe an experiment to show that hydrogen chloride gas in readily soluble in water. Name another gas which shows this property.

 

Tutorial question

– Name any five gases that are collected by downward delivery

– Name any three gases that are collected by upward delivery

– Draw and label the laboratory set up for the preparation of chlorine gas

– Show the reaction of chlorine with saturated hydrocarbon to produce chlorinated hydrocarbon.

What condition is required for the reaction to take place? What type of reaction is it?

– With the aid of balanced equations, show the reaction of chlorine with (i) dilute alkali (ii) conc.

 

Alkali solution

– State five physical properties of hydrogen chloride

– State what happens when chlorine reacts with iron (II) chloride

– State five general properties of halogens

– How do we test for chlorine

– State three uses of chlorine