Chemistry Form 6 Sem 2 06

Element Chlorine, Cl₂ Bromine, Br₂ Iodine, I₂ Atomic radius (nm) 0.099 0.114 0.133 Physical state at room temperature Electronic configuration Melting point (o_{C) -102 -7.1} --Yellow greenish gas Red brown liquid Black solid 3s2 _3p5 _4s2 _4p5 _5s2 _5p5 Melting point ( C) -102 -7.1 --Boiling point (o_{C) -34.5 58.6 184} Bond energy (kJ/mol) 242.5 191.5 150.0 Electron affinity (kJ/mol) -364 -340 -297 Electronegativity 3.0 2.8 2.5

1. Atomic radius

When goes down to group 17,

atomic radius increase

. This is due

to, when going down to Group 17,

nuclear charge increase

, as the

proton number increase. As number of proton increase, number

of electrons also increased, which caused

more shells are required

to fill in electrons and gradually

increase the screening effect

. As a

result, the

effective nuclear charge decreased

and caused the

atomic radius increase

2. Melting point, boiling point, physical state at room temperature,

colour intensity & density of Group 17

colour intensity & density of Group 17

All Group 17 elements

exist as diatomic molecules

, where

fluorine, chlorine, bromine, iodine and astatine exist as F

,

Cl

,Br

,I

,and At

respectively

The physical state of chlorine, bromine and iodine are different

despite they are from the same classification of elements,

non-metal.

Chlorine

exist as

yellow-greenish gas

,

bromine

exist as

brown liquid

, while

iodine

exist as

purplish-black solid

. These

observations were supported by the data of melting point and

boiling point of Group 17 elements

All these were due to when going down to Group 17,

relative

molecular mass increased

from Cl

to Br

to I

and

caused the

weak Van Der Waals' forces increase

. As a result, the melting

point and boiling point increase down the Group and each

elements of Group 17, hence exist as different physical forms.

At the same time, as the weak intermolecular forces increased,

molecules become more close packed

, causing the

number of

molecules become more close packed

, causing the

number of

molecules per unit volume increased

, hence

increase the density

of elements, which consequently

increased the colour intensity

of halogen. This will cause the colour of halogens to become

darker. This can also be used to explain the volatility of

molecules, where going down Group 17,

volatility decreased

which were also due to

stronger intermolecular forces between

particles

hence harder to vapourise.

3. Electronegativity, bonding energy and electron affinity

Electronegativities measure the ability to pull the bonding pair

towards the direction of the atom.

Electronegativity

of the group 17

decrease

when goes down to

group due to the

increase of covalent radius of the elements.

Fluorine has the highest electronegativity among all Group 17

elements. This will influence the bonding length of the Group 17

which directly causing the decrease in bonding energy.

Bonding energy measure the heat (energy) required to break the

Bonding energy measure the heat (energy) required to break the

bond formed between 2 atoms.

Example :

Cl - Cl (g) → 2 Cl (g) ∆H = + 242.5 kJ mol

Bonding energy decreased, as the

bond length increased

from

chlorine to iodine. As the bond length of the covalent molecules

increased, it is getting

easier to break the bond

, thus lower the

bonding energy.

Fluorine has a lower bonding energy

compare to chlorine as the

fluorine atom is too small

, which result the

lone pair electrons

between F-F repel each other,

causing the bond length increased.

(since lone pair - lone pair electrons repulsion is stronger than

lone pair - bond pair electrons repulsion).

Electron affinity is the heat liberated when one mole of gaseous

atom accept one mole of electron to form a gaseous anion under

standard condition.

standard condition.

Equation : F (g) + e- → F

(g) ∆H

= - 364 kJ mol

The electron affinity of Group 17 become less negative when goes

down to group 17. This is due to, as atomic size increased,

nucleus is further away from the outermost shell. As a result,

electron received is further away and the electrostatic forces of

attraction is weaker, hence less exothermic

4. Solubility of halogens in water

Solubility of halogen decrease when goes down to Group 17.

Fluorine is the most soluble halogen among the Group 17

because enthalpy change of hydration of fluorine is very

exothermic. Furthermore, fluorine can form hydrogen bond with

water.

Both chlorine and bromine are considerably soluble in water and

form an acid and acid halide via disproportionation reaction

form an acid and acid halide via disproportionation reaction

(oxidation and reduction occur simultaneously on the same

substance), while iodine is sparingly soluble in water.

Equation :

Cl

(g) + H

O (l) → HCl (aq) + HOCl (aq)

Br

(g) + H

O (l) → HBr (aq) + HOBr (aq)

Iodine aqueous solution can be prepared by dissolving iodine in

potassium iodide (KI) aqueous solution and form a brown

solution of I

_.

Equation :

I

(s) + I

_{(aq) ↔ I}

The hypohalous acid formed can be used as ……… agent.

When hypohalous is exposed to sunlight, it will decomposed and

oxygen is evolved. Hypochlorous acid and hypobromous acid both give different reaction.

2 HOCl 2 HCl (aq) + O₂ (g)

Overall : 2 Cl₂ + 2 H₂O 4 HCl (aq) + O₂ (g)

As extra notes, halogens also dissolve in organic solvent such as bleaching

As extra notes, halogens also dissolve in organic solvent such as tetrachloromethane, CCl₄, or benzene solution.

Halogen

Chlorine, Cl

Bromine, Br

Iodine, I

Colour in aqueous

solution

Pale yellow

Brown

Brown

Colour in organic

6.2

Chemical properties of Group 17 elements

Halogens are strong oxidising agent. This is supported with the

value of standard reduction potential where

½ F

_{(aq) + e-
F}

_(aq)

_E

_{= + 2.87 V}

½ Cl

_{(aq) + e-
Cl}

_(aq)

_E

_{= + 1.36 V}

½ Br

_{(aq) + e-
Br}

_(aq)

_E

_{= + 1.07 V}

½ I

_{(aq) + e-
I}

_(aq)

_E

_{= + 0.54 V}

A displacement reaction can take place between halogens. Using

oxidising agent

A displacement reaction can take place between halogens. Using

this method, the reactivity of halogen can also be deduced. The

following mixture of halide and halogen are added and

Observation Equation Chlorine in Tetrachloromethane is added to aqueous potassium bromide (KBr). Colourless solution in CCl₄ turned brown Reactants : Cl₂ + Br

-So the half equation for reaction: Cl₂ + 2e-
_{2 Cl}- _E0 _{= + 1.36 V} 2 Br-
_Br 2 + 2e- E0 = - 1.07 V ---Cl₂ + 2 Br-
_{2 Cl}- _{+ Br} 2 E_cell = + 0.29 V Since the E_cell is positive, the reaction is spontaneous. Bromine in tetrachloromethane is added to aqueous potassium iodide (KI) cell reaction is spontaneous. Brown colour in CCl₄ turn purple colour. Reactants : Br₂ + I

-So the half equation for reaction: Br₂ + 2e-
_{2 Br}- _E0 _{= +1.07 V} 2 I-
_I 2 + 2e- E0 = - 0.54 V ---Br₂ + 2 I-
_{2 Br}- _{+ I} 2 E_cell = + 0.53 V Since the E_cell is positive, the reaction is spontaneous.

Iodine is tetrachloromethane is added to aqueous potassium chloride Purple colour in CCl₄ remain unchanged Reactants : I₂ + Cl

-So the half equation for reaction: I₂ + 2e-
_{2 I}- _E0 _{= + 0.54 V} 2 Cl-
_Cl 2 + 2e- E0 = - 1.36 V ---I₂ + 2 Cl-
_{2 I}- _{+ Cl} 2 potassium chloride (KCl) unchanged I2 + 2 Cl 2 I + Cl2 E_cell = – 0.82 V Since the E_cell is negative, the reaction is non-spontaneous.

5.2.1 Changing of iron (II) ion to iron (III) ion and vice versa

Given the half equation of iron (II) ion and iron (III) ion and its standard reduction potential.

Fe3+ (aq) + e-
Fe2+ (aq) E∅ = + 0.77 V

If the halogen aqueous solution is added to iron (II) ion, Fe2+, solution

respectively

Chlorine in iron (II) ion aqueous solution.

Reactants : Cl₂ + Fe2+ _{; So the half equation for reaction:} Cl₂ + 2e-
_{2 Cl}- _E0 _{= + 1.36 V}

Fe2+
_Fe3+ _{+ e}- _E0 _{= – 0.77 V} Cl + 2 Fe2+
_{2 Cl}- _{+ 2 Fe}3+ _{E = + 0.59 V}

Bromine in iron (II) ion aqueous solution. Cl₂ + 2 Fe2+
_{2 Cl}- _{+ 2 Fe}3+ _E

cell = + 0.59 V

Since the E_cell is positive, the reaction is spontaneous.

Obs : pale yellow solution dissolve in green solution to form yellow solution

Reactants : Br₂ + Fe2+ _{; So the half equation for reaction:} Br₂ + 2e-
_{2 Br}- _E0 _{= + 1.07 V}

Fe2+
_Fe3+ _{+ e}- _E0 _{= – 0.77 V} Br₂ + 2 Fe2+
_{2 Br}- _{+ 2 Fe}3+ _E

cell = + 0.30 V

Since the E_cell is positive, the reaction is spontaneous.

Obs : brown solution first dissolve in green solution and form yellow solution

Iodine in iron (II) ion aqueous solution.

Reactants : I₂ + Fe2+ _{; So the half equation for reaction:} I₂ + 2e-
_{2 I}- _E0 _{= + 0.54 V}

Fe2+
_Fe3+ _{+ e}- _E0 _{= – 0.77 V} I₂ + 2 Fe2+
_{2 I}- _{+ 2 Fe}3+ _E

cell = – 0.23 V

Since the E_cell is negative, the reaction is non-spontaneous. Obs : no changes occur

*Conversely when iron (III) ion is added to iodide ion, 2 I- _{+ 2 Fe}3+
_I

2 + 2 Fe2+ Ecell = + 0.23 V

iodide ion (which act as reducing agent) are able to reduce iron (III) ion to form iron (II) ion and iodine is formed.

It can be conclude that

………..…… form iron (II) ion and iodine is formed.

Chlorine and bromine can act as oxidising agent while iodine is not a good oxidising agent. Conversely, iodide ion can act as a good reducing agent.

The oxidation of thiosulphate ions, S2O32-, by halogen is quantitative.

Using volumetric analysis, the exact concentration of the reagent (S₂O₃2-_{) used can be found.}

Since chlorine and bromine are strong oxidising agent, when chlorine

and bromine react with thiosulphate solution, thiosulphate ion is oxidise to its highest oxidation state, which is ………sulphate ion, SO4

2-S₂O₃2- _{+ 5 H}

2O
2 SO42- + 8 e– + 10 H+ Cl₂ + 2 e–
_{2 Cl}–

---S O 2- _{+ 5 H O + 4 Cl
2 SO} 2- _{+ 8 Cl}– _{+ 10 H}+

But when it react with a weaker oxidising halogen like iodine, it only

slightly oxidised. Thiosulphate ion will be oxidised to become

tetrathionate ion, S₄O₆2-_{, which can be used in the volumetric analysis}

S₂O₃2- _{+ 5 H} 2O + 4 Cl2
2 SO42- + 8 Cl– + 10 H+ 2 S₂O₃2-
_S 4O62- + 2 e– I₂ + 2 e–
_{2 I}– ---2 S₂O₃2- _{+ l} 2
S4O62- + 2 l–

Example : The carbon monoxide in a sample of polluted air can readily be

determined by passing it over solid iodine(V) oxide, I₂O₅, to give carbon dioxide and iodine.

(a) Write a balanced equation reaction between carbon monoxide and iodine(V) oxide. ... (b) The iodine produced is removed and titrated with aqueous sodium thiosulphate:

2 Na₂S₂O₃ (aq) + I₂ (s) → Na₂S₄O₆ (aq) + 2 NaI (aq)

A 1.0 dm3 _{sample of air produced iodine that required 20.0 cm}3 _{of 0.10 mol dm}-3

sodium thiosulphate to discharge the iodine colour. Calculate the mass of carbon monoxide in this sample of polluted air. [4]

I₂O_{5 + 5 CO
5 CO2} + I₂

monoxide in this sample of polluted air. [4]

Mol of Na₂S₂O₃ = MV / 1000 ; mol = (0.10)(20.0) / 1000 Mol of Na₂S₂O₃ = 0.002 mol

From eq (b) since 2 mol of Na₂S₂O₃ = 1 mol of I2 Mol of I₂ = 0.001 mol

In (a) since 1 mol of I₂ = 5 mol of CO Mol of CO = 0.005 mol

Mass of CO = mol x RMM @ mass = 0.005 x (12.0 + 16.0) Mass of CO = 0.14 g

6.3 Hydrogen halide and hydrides.

When chlorine is mixed with hydrogen gas under the presence of

sunlight, an explosive reaction occur and yellow-greenish (chlorine) gas turned to white fumes (hydrogen chloride).

Equation : Cl₂ (g) + H₂ (g) → 2 HCl (g)

The reactivity decreased when going down to Group 17 (as suggested by

Halides HF HCl HBr HI

Boiling point (o_{C) 23 - 85 - 66 - 35}

Bonding enthalpy (kJ/mol) +560 +433 +366 +298

The reactivity decreased when going down to Group 17 (as suggested by

their respective E0 _{value) where bromine react with hydrogen upon}

heating to 2000_{C to form hydrogen bromide}

While iodine react partially with hydrogen until reaction established an

equilibrium between hydrogen, iodine and hydrogen iodide, at 450o_C

All hydrogen halide exist as gas under room temperature and pressure,

except hydrogen fluoride, which exist as a volatile liquid at room temperature. Hydrogen fluoride exist as volatile liquid because H-F has strong hydrogen bond between them, which are stronger than the weak Van Der Waals forces that exist between HCl, HBr and HI. The boiling points eventually increased from HCl < HBr < HI, since the relative molecular mass increased, which increase the weak

intermolecular forces.

The thermal stability of hydrogen halide decreased down Group 17. This can be explained by the fact of the increasing bond length of H This can be explained by the fact of the increasing bond length of H -X. When going down Group 17, the atomic radius increased. As a result, the bond length between hydrogen atom and halogen atom increased, which eventually caused lesser heats were required to break the

covalent bond between H - X. When heated at high temperature, all hydrogen halide will decomposed to form back their respective

elements. Some hydrogen iodide can even decomposed to form hydrogen and iodine under room temperature pressure until equilibrium is established.

The factor of bond length can also be used to explain the acidity of

hydrogen halide. When dissolved in water, acidity increase from HF <<HCl<HBr< HI.

Hydrofluoric acid is a weak acid, which dissociate partially in water,

where

HF (g) + H₂O (l) F- _{(aq) + H}

3O+ (aq)

while hydrochloric acid, bromic acid, iodic acid dissociate completely

in water. [where X may be Cl, Br, I]

HX (g) + H₂O (l) X- _{(aq) + H}

3O+ (aq)

2 3

The acidity increase down Group 17 since the bond length of H-X

increase. As a result, H-X is easier to dissociate in water, hence causing the acidity increased.

Even though hydrogen halide can be prepared using direct reaction

between hydrogen and halogen, but in laboratory / industrial process, they are prepared in different way. Hydrogen fluoride and chloride can be prepared using fluoride / chloride salt with concentrated sulphuric acid (H₂SO₄), while hydrogen bromide and iodide is prepared via

5.3.1 Displacement from salt

When react with concentrated sulphuric acid is added to halides salt and heated, hydrogen halide fumes will be released

MX (s) + H₂SO_{4 (aq)
HX (g) + MHSO4} (aq)

Using this method, hydrofluoric acid can be produced, when fluorspar

(calcium fluoride) react with hot concentrated sulphuric acid to form hydrogen fluoride, which is a super acid.

Hydrogen chloride is produced when sodium chloride is heated and

2 H

SO

(aq) + CaF

_(s)
Ca(HSO4

)

(s) + 2HF (aq)

Hydrogen chloride is produced when sodium chloride is heated and concentrated acid is added slowly to the salt solution.

Equation :

Hydrogen bromide and hydrogen iodide cannot be prepared using

this method as concentrated sulphuric acid is a strong oxidising

agent, and this will result the hydrogen halide produced to further oxidise to become the individual halogen involved.

For bromide salt ;

NaBr (s) + H

SO

_{(aq)
HBr (g) + NaHSO}

(aq)

2 HBr (g) + H

SO

_(aq)
Br

(l) + SO

(g) + 2 H

O (l)

For iodide salt ;

NaI (s) + H

SO

_{(aq)
HI (g) + NaHSO}

(aq)

8 HI (g) + H

SO

_{(aq)
4 I}

(l) + H

S (g) + 4 H

O (l)

Observation : black spot is formed on the wall of the tube. A bad egg smell gas evolved from the mixture.

6.3.2

Hydrolysis of covalent halides

In order for hydrolysis of phosphorous halide to take place,

phophorous (III) tribromide and phosphorous triiodide must be

first prepared by direct reaction between phosphorous and halogen,

where

Equation :

P

(s) + 6 X

(l) → 4 PX

(l) [X may be Br ot I]

[However, when excess chlorine is reacted directly with

phosphorous, it formed an expand octet of

phosphorouspentachloride, PCl

. As for phosphorous

pentabromide, PBr and phosphorous pentaiodide, PI , cannot be

5

pentabromide, PBr

and phosphorous pentaiodide, PI

, cannot be

formed due to steric hindrance effect]

Bromine and iodine react with phosphorus to form phosphorus(III)

tribromide and phosphorus(III) iodide respective.

PBr

(l) + 3 H

O (l) →

PI

(l) + 3 H

O (l) →

This method is suitable for preparing hydrogen bromide and

hydrogen iodide since the phosphorous acid, H

PO

, produced is

not an oxidising agent

3 HBr (aq) + H₃PO₃ 3 HI (aq) + H₃PO₃

5.4 Oxoacid of chlorine.

All halogens (except fluorine) can exhibit various oxidation states in their compound

In chlorine, it can exhibit 6 oxidation states

Chemical Hydrochloric_acid Chlorine Hypochlorous acid Chlorous acid Chlorite acid Perchlorite acid

Formula HCl Cl₂ HClO HClO₂ HClO₃ HClO₄

Oxidation

– 1

0

+1

+3

+5

+7

The hydrogen chloride which contain oxygen atom in it, is called as ………. All oxoacids are soluble in water. When dissolve in water Oxidation

state

– 1

0

+1

+3

+5

+7

HClO (Hypochlorous acid) HClO₂ (Chlorous acid) HClO₃ (chloric acid) HClO₄ (Perchloric acid) HClO + H₂O
ClO– _{+ H} 3O+ HClO₂ + H₂O
ClO₂– _{+ H} 3O+ HClO₃ + H₂O
ClO₃– _{+ H} 3O+ HClO₄ + H₂O
ClO₄– _{+ H} 3O+

As oxygen is more ………. than chlorine, O – H bond in the

oxoacid are weakened by the inductive effect (electron withdrawing effect) of the oxygen atoms which are bonded to chlorine atom. As a result, the acidity ………. with the increase of number of oxygen atom.

Explain this : HI is a stronger acid than HBr but HIO3 is a much weaker

acid than HBrO₃.

………. electronegative

increase

HI is stronger acid than HBr since the bond length is longer in H–I

………. ………. ………. HI is stronger acid than HBr since the bond length is longer in H–I

compare to H–Br. However, HBrO₃ is a stronger acid than HIO₃ since Br is more electronegative than I, causing the inductive effect of HBrO₃ is stronger, hence the dissociation is stronger.

5.5 Preparation of Chlorine

In laboratory, chlorine gas is produced by using manganese base compound

When hot concentrated HCl react with potassium permanganate:
When hot concentrated HCl react with manganese oxide :

In industrial process, chlorine gas, together with sodium metal, is

prepared using molten sodium chloride (brine) using

mercury-16 HCl + 2 KMnO

_{2 KCl + 2 MnCl2}

+ 8 H

O + 5 Cl

4 HCl + MnO

_MnCl2

+ 2 H

O + Cl

prepared using molten sodium chloride (brine) using mercury-cathode cell.

In industrial process, chlorine gas, together with sodium metal, is

prepared using molten sodium chloride (brine) using mercury-cathode cell

Half equation occur at anode :
Half equation occur at cathode :

Mercury is specially used to attract the sodium formed in cathode and

form an alloy named amalgam

This method is not environment friendly as the mercury used is

poisonous.

2 Cl–
_{2 e}– _{+ Cl} 2 Na+ _{+ e}–
_Na

Despite on the disadvantage given by the mercury, a diaphragm cell is used.

Similar to the mercury-cathode cell, the electrolytes used in diaphragm cell is also ………..

The process inside the diaphragm cell is known as the chlor-alkali process

When sodium chloride dissociates under the effect of an electric current, the chloride ions are discharged.

At anode : 2 Cl- _{(aq) → Cl}

2 (g) + 2e-.

Titanium is chosen as the anode because it resists corrosion by the very reactive chlorine

At cathode, since the sodium ion (Na+_{) is attracted to cathode}

through the diaphragm, the selectivity to discharge is between sodium ion and water molecule.

Standard reduction potential of sodium :

Na+ _{+ e}-
_{Na E}0 _{= – 2.71 V}

Na _{+ e
Na} E =

Standard reduction potential of water : 2 H₂O (l) + 2 e-
_H

2 (g) + 2 OH- (aq) E0 =

Since BBBB has a higher SRP value, BBBB is discharge and BBBBBBBBBBB...B is produced.

The level of brine (left or anode position) is always placed higher

than the water (right or cathode position) to prevent a back pressure of the cell and form back flow. If back flow occurs, water will dilute the concentrated brine and chlorine will not be selected.

– 2.71 V – 0.83 V

water water

5.6 Reaction with silver nitrate

When aqueous halide ions (such as,sodium chloride solution) are added to silver nitrate solution,coloured precipitates are formed.

NaCl (aq) + AgNO₃ (aq) → AgCl (s) + NaNO₃ Ag+ (aq) + Cl- (aq) → AgCl (s) (white precipitate)

Silver chloride is a white precipitate which dissolves in excess aqueous ammonia to form a soluble complex of silver

AgCl (s) + 2 NH₃ (aq) → [Ag(NH₃)₂]+_Cl- _{(aq) (soluble complex)}

If chloride ion is replaced by bromide ion, the precipitate formed is cream

If chloride ion is replaced by bromide ion, the precipitate formed is cream colour (AgBr)

Ag+ _{(aq) + Br}- _{(aq) → AgBr (s) (cream precipitate)}

Unlike silver chloride, silver bromide dissolve only in concentrated ammonia aqueous solution according to the equation :

AgBr (s) + 2 NH₃ (aq) → [Ag(NH₃)₂]+Br- (aq) (soluble complex)
If iodide ion is replaced, silver iodide (AgI) produced is yellow colour

5.7 Reaction with aqueous sodium hydroxide

Chlorine reacts with dilute sodium hydroxide solution to form sodium hypochlorite and NaCl.(low temp)

Cl_{2 (g) + 2 NaOH (aq)}
NaCl (aq) + NaClO (aq) + H₂O (l) * When NaClO is exposed to uv (sunlight) it will undergoes

decomposition to form sodium chloride and oxygen gas. 2 NaClO (aq)
2 NaCl (aq) + O2 (g)

0 –1 +1

2 NaClO (aq)
2 NaCl (aq) + O2 (g)

The reaction above can also be describe as a futher decomposition upon heating hypochlorite ion

3 NaClO (aq)
NaClO₃ (aq) + 2 NaCl (aq)

When concentrated sodium hydroxide is heated to a high temperature of 70o_{C, sodium chlorate(v) is formed.}

3 Cl_{2 (g) + 6 NaOH (conc)
5NaCl (aq) + NaClO3} (aq) + 3 H₂O (l)

+1 +5 –1

For bromine, it behave the same way as chlorine. When bromine

is added to aqueous sodium hydroxide

3 Br

_{(l) + 6 NaOH (aq)
NaBrO}

(aq) + 5 NaBr + 3 H

O

Iodine reacts with NaOH in the same ways Cl

and Br

.

Disproportionation occurs even at 0

_C.

3 I (l) + 6 NaOH (aq)
NaIO (aq) + 5 NaI + 3 H O

0 _{+5 –1}

3 I

_{(l) + 6 NaOH (aq)
NaIO}

(aq) + 5 NaI + 3 H

O

It is therefore impossible to prepare the IO

_{ions without the}

presence of IO

5.8

Uses of Halogen and Their Compounds

Fluorine

As oxidising agents.

Hydrogen fluoride is used to etch glass.

Sodium fluoride is added to the water and toothpaste to prevent

dental caries.

Fluorine in industry is the production of polymer plastic-teflon.Teflon

is an insulator of electricity and as linings in non- stick pans

.Medically,teflon is an important material in surgical transplants as it

can be used to replace broken cartilages.

NaF or SnF

are important constituents in toothpaste and fluoride

ion,F-, forms a protective layer

Chlorine

Used as bleaching agents. Bleach contains sodium hypochlorite or

sodium chlorate(I) and prepared by passing chlorine gas into a cold

sodium hydroxide solution.

Cl

(g) + 2NaOH(aq)

_{NaCl (aq) + NaClO (aq) + H}

O (l)

Add to water and swimming pool to kill germs and bacteria. In water

forms hydrochloric and hypochlorus acids and ClO- has the

disinfecting effect.

Cl

(g) + H

O (l)

_{HCl (aq) + HClO (aq)}

Dichlorodiphenyltrichloroethane (DDT) is used as insecticide.

Chloroethene is the monomer for PVC.

Chlorine dioxide,ClO

is used to bleach wood pulp for paper

making.

Potassium chlorate(V) is used as an oxidant in fireworks and

matches

Bromine

Bromine occurs naturally in some marine organisms.

1,2-dibromoethene,BrCH=CHBr is used in insecticides.

1,2-dibromoethane,BrCH

CH

Br is added to leaded petrol to

remove the lead oxide sticking to the piston surface.

Silver bromide is used in photography. The surface of the paper is

spread with a layer of silver bromide. When exposed to sunlight,

the bromide ion will be oxidised to bromine gas, leaving a grey

the bromide ion will be oxidised to bromine gas, leaving a grey

layer of silver on the surface of the paper

This is a redox reaction

Ag

_{+ e- Ag [reduction]}

2Br-  Br2

+ 2 e- [oxidation]

Overall Reaction

Iodine

Iodine in alcohol is a popular antiseptic used medicinally to treat

wounds.

The iodine vapour is used in electric bulbs.