bài tập và lý thuyết Halogens

• Halogens gain an additional electron to form the halide ions  combine with metals to form metal halides  held together by ionic bonding Bonding and Oxidation State... • Aqueous ch

41.1 Characteristic Properties of the Halogens 41.2 Variation in Properties of the Halogens 41.3 Comparative Study of the Reactions of

Halide Ions 41.4 Acidic Properties of Hydrogen Halides and

the Anomalous Behaviour of Hydrogen Fluoride

41

Characteristic Properties of the

Halogens

• Group VIIA elements include

• Astatine

 radioactive

Introduction

• Group VIIA elements

 also called halogens

Introduction

The halogens

• In the free elemental state

 form diatomic molecules

 complete their octets by sharing their

single unpaired p electrons

Introduction

• When halogens react with other elements

 complete their octets

 depending on the electronegativity of

the element

Introduction

• Either

 gaining an additional electron to form

halide ions

 or sharing their single unpaired p

electrons to form single covalent bonds

Introduction

Appearances of halogens at room temperature and pressure: chlorine

chlorine

bromine Appearances of halogens at room

temperature and pressure: bromine

Appearances of halogens at room temperature and pressure: iodine

iodine

Electronegativity is the relative tendency

of an atom to attract bonding electrons towards itself in a covalent bond.

High Electronegativity

• All halogens

 high electronegativity values

 high tendency to attract an additional

electron to achieve the stable octet electronic configuration

 highest among the elements in the

same period

High Electronegativity

Electronegativity values of halogens

Halogen Electronegativity value

FClBrIAt

4.03.02.82.52.2

Electron affinity is the enthalpy change

when one mole of electrons is added to

one mole of atoms or ions in the gaseous state

High Electron Affinity

• All halogens

 negative values of electron affinity

 high tendency to attract an additional

electron to form the respective halide ions

High Electron Affinity

Electron affinities of halogens

Halogen Electron affinity (kJ mol –1 )

FClBrIAt

–348–364–342–314–285

• Halogens

 gain an additional electron to form

the halide ions

 combine with metals to form metal

halides

 held together by ionic bonding

Bonding and Oxidation State

• The oxidation states of the halogens = –1

Bonding and Oxidation State

• The halogen atoms

 share their unpaired p electrons with

a non-metallic atom

 form a covalent bond

Bonding and Oxidation State

• Halogens (except fluorine )

 exhibit an oxidation state of –1 or +1

in the covalent molecules formed

 depend on the electronegativity of the elements that are covalently

Bonding and Oxidation State

• All halogens (except fluorine )

 can expand their octets of electrons

by utilizing the vacant , low-lying d

orbitals

Bonding and Oxidation State

• Each of these halogen atoms

 have variable numbers of unpaired

electrons to pair up with electrons from other atoms

 able to form compounds of different

oxidation states

Bonding and Oxidation State

“Electrons-in-boxes” diagrams of the electronic configuration of a halogen atom of the ground state

+7

Cl2O7 H5IO6HClO4 HIO4ClO4– IO4–

• Fluorine

 cannot expand its octet

 no low-lying empty d orbitals

available

 the energy required to promote

electrons into the third quantum shell

is very high

Bonding and Oxidation State

• Fluorine

 the most electronegative element

 only one unpaired p electron

available for bonding

 oxidation state is limited to –1

Bonding and Oxidation State

Colour

• All halogens

 the absorption of radiation in the

visible light region of the

electromagnetic spectrum

Colour

• The absorbed radiation

 the excitation of electrons to higher

energy levels

Colour

• Fluorine atom

 smaller size

 absorb the radiation of relatively high

frequency (i.e blue light )

 appears yellow

Colour

• Iodine

 absorbs the radiation of relatively low

frequency (i.e yellow light )

 appears violet

Colour

 different colours when dissolved in

different solvents

Colours of halogens in pure form and in solutions

in pure form in water in 1,1,1-trichloroethane

F2 Pale yellow Pale yellow Pale yellow

Cl2 Greenish yellow Pale yellow Yellow

Br2 Reddish brown Yellow Orange

I2 Violet black Yellow (only

slightly soluble) Violet

Colours of halogens in water: (a) chlorine; (b) bromine; (c) iodine

Colours of halogens in 1,1,1-trichloroethane:

(a) chlorine; (b) bromine; (c) iodine

Check Point 41-1

• All halogens

 exist as diatomic molecules

• In the diatomic molecules

 the halogen atoms are held together

by strong covalent bonds

 only held together by weak van der

Waals’ forces (i.e instantaneous dipole-induced dipole interaction)

• The physical properties of halogens

 strongly affected by the way that the

atoms are joined together

 the interactions that hold the

molecules together

Some physical properties of the halogens

–

0.0720.0990.1140.133–

0.1330.1810.1950.216–

Some physical properties of the halogens

Halogen Melting point (°C) Boiling point (°C) Density at 20 °C

–188–34.758.8184380

1.111.563.124.93–

Variation in Physical Properties

 exist as non-polar diatomic molecules

1 Melting Point and Boiling Point

1 Melting Point and Boiling Point

 the melting points and boiling points

of halogens increase

1 Melting Point and Boiling Point

• These physical properties depend on

 the strength of van der Waals’ forces

holding the halogen molecules together

1 Melting Point and Boiling Point

 the molecular size increases

 the electron clouds of the molecules

become larger

 more polarizable

1 Melting Point and Boiling Point

• Instantaneous dipoles

 more readily formed

 the instantaneous dipole-induced

dipole interaction between the molecules is stronger

1 Melting Point and Boiling Point

• A greater amount of energy is required

 separate the molecules in the

processes of melting and boiling

 the melting points and boiling points

increase progressively from fluorine

to astatine

Variations in melting point and boiling point of the halogens

2 Electronegativity

Electronegativity is the relative tendency

of the nucleus of an atom to attract

bonding electrons towards itself in a

covalent bond.

2 Electronegativity

 the electronegativity values of

halogens decrease

2 Electronegativity

 the atomic size increases

 the number of electron shells

increases

 creates a greater screening effect

2 Electronegativity

• The atomic size increases

 The tendency of the nucleus of the

halogen atom attract bonding electrons towards itself in a covalent bond decreases

Variations in electronegativity value

of the halogens

3 Electron Affinity

Electron affinity of halogens is the

enthalpy change when one mole of

electrons is added to one mole of halogen atoms or ions in the gaseous state

3 Electron Affinity

• The electron affinity

 increases from fluorine to chlorine

 decreases from chlorine to astatine

3 Electron Affinity

• The general decrease in electron affinity

 the atomic size increases

 the number of electrons shells down

the group increases

 the effective nuclear charge decreases

 tendency of the nuclei of halogen atoms to attract additional electrons

decreases

3 Electron Affinity

• Fluorine

 abnormally low electron affinity

3 Electron Affinity

• Fluorine atom

 very small atomic size

 energy is required to overcome the

repulsion between the additional electron and the electrons present in the electron shell

Variations in electron affinity

of the halogens

Check Point 41-2A

Variation in Chemical Properties

 the most reactive group of

non-metallic elements

 all halogens have one electron short

of the octet electronic configuration

∴ tend to attract an additional electron

to attain the octet electronic configuration

Variation in Chemical Properties

Variation in Chemical Properties

• Fluorine

 very strong oxidizing agent

Variation in Chemical Properties

• Other elements that combine with

fluorine

 have their highest possible oxidation

numbers

1 Relative Oxidizing Power of Halogens

• All halogens

 combine directly with sodium to form

sodium halides

 the reactivity decreases down the

group from fluorine to iodine

Reactions with Sodium

Reactions with Sodium

• Fluorine

 react explosively to form sodium

fluoride

2Na(s) + F2(g) → 2NaF(s)

• Bromine

 burns steadily in bromine vapours to

form sodium bromide

2Na(s) + Br2(g) → 2NaBr(s)

Reactions with Sodium

• Iodine

 burns steadily in iodine vapours to

form sodium iodide

2Na(s) + I2(g) → 2NaI(s)

Reactions with Sodium

• Aqueous chlorine

 oxidizes green iron(II) ions to

yellowish brown iron(III) ions

Reactions with Iron(II) Ions

2Fe2+(aq) + Cl2(aq)

→ 2Fe3+(aq) + 2Cl–(aq)

= +0.59 V

• Aqueous bromine

 oxidizes green iron(II) ions to

yellowish brown iron(III)

Reactions with Iron(II) Ions

2Fe2+(aq) + Br2(aq)

→ 2Fe3+(aq) + 2Br–(aq)

= +0.30 V

• Iodine

 a mild oxidizing agent

 not strong enough to oxidize iron(II)

ions.

Reactions with Iron(II) Ions

• The spontaneity of a reaction can be

worked out

 adding the standard electrode

potentials of the two half reactions

concerned

Reactions with Iron(II) Ions

Reactions with Iron(II) Ions

 the reaction is predicted to be

spontaneous

• If the overall standard electrode potential

(i.e the standard cell electromotive force,

) is a positive value

• Aqueous chlorine and bromine

Reactions with Iron(II) Ions

 the oxidation reactions of aqueous

iron(II) ions are spontaneous

 the for both reactions are

positive

Standard electrode potentials of some related half reactions

Half reaction Standard electrode potential (V)

• Aqueous iodine

Reactions with Iron(II) Ions

 this reaction is not spontaneous

• Thiosulphate ions

 a reducing agent

 reacts differently with halogens of

different oxidizing power

Reactions with Thiosulphate Ions

• Iodine

 reacts with sodium thiosulphate to

form sodium tetrathionate and

sodium iodide

Reactions with Thiosulphate Ions

• This is a typical reaction

 determine the concentration of iodine

in a solution

 by titration with standard thiosulphate

solution (iodometric titration)

I2(aq) + 2S2O32–(aq)

→ 2I–(aq) + S4O62–(aq)

Reactions with Thiosulphate Ions

• Chlorine and bromine

 more powerful oxidizing agents

 oxidize thiosulphate ions to

sulphate(VI) ions

Reactions with Thiosulphate Ions

Reactions of halogens with sodium

Halogens Reactant

Sodium They react

violently to form sodium chloride

Sodium burns in bromine vapour

to form sodium bromide

Sodium burns in iodine vapour to form sodium iodide

Reactions of halogens with iron(II) ions

Halogens Reactant

Iron(II) ions The green iron(II)

ions are oxidized

to yellowish brown iron(III) ions

The green iron(II) ions are oxidized

to yellowish brown iron(III) ions

The solution remains green since iron(II) ions are not oxidized

by iodine

Reactions of halogens with thiosulphate ions

Halogens Reactant

Thiosulphate

ions

The thiosulphate ions are oxidized

to sulphate(VI) ions

The thiosulphate ions are oxidized

to sulphate(VI) ions

The thiosulphate ions are oxidized

to tetrathionate and iodide ions

• The relative oxidizing power of the

halogens decreases in the order:

F2 > Cl2 > Br2 > I2

Reactions with Thiosulphate Ions

2 Disproportionation of the Halogens in

Water and Alkalis

• Fluorine

 reacts vigorously with water to form

hydrogen fluoride and oxygen

2F2(g) + 2H2O(l) → 4HF(aq) + O2(g)

Reactions with Water

• Chlorine

 less reactive than fluorine

 reacts with water to form hydrochloric

acid and chloric(I) acid (also known

as hypochlorous acid

Reactions with Water

• The oxidation number of chlorine

Reactions with Water

Disproportionation is a chemical change in which oxidation and reduction of the same species (which may be a molecule, atom

or ion) take place at the same time

• Chlorate(I) ion (also known as

hypochlorite ion )

 an unstable ion

 decomposes when exposed to

sunlight or high temperatures to give

chloride ions and oxygen

2OCl–(aq) → 2Cl–(aq) + O2(g)

Reactions with Water

• Bromine

 only slightly soluble in water

 mainly exists as molecules in

saturated bromine water

Reactions with Water

• When the solution is diluted

 hydrolysis takes place

 hydrobromic acid and bromic(I) acid

(also called hydrobromous acid) are formed

Reactions with Water

Br2(l) + H2O(l)

HBr(aq) + HOBr(aq)

• Bromate(I) ion

 also unstable

 forms colourless compounds when

reacting with dyes OBr–(aq) + dye

• Iodine

 does not react with water

 only slightly soluble in water

Reactions with Water

• Iodine

 soluble in potassium iodide solution

 exists as triiodide ions in thesolution

 often called iodine solution

I2(s) + KI(aq) → KI3(aq)

iodine solution

Reactions with Water

• All halogens

 react with aqueous alkalis

Reactions with Alkalis

• The reactions between halogens and

aqueous alkalis

 disproportionation (except fluorine )

Reactions with Alkalis

• Halogens

 react differently under cold / hot and

dilute / concentrated conditions

Reactions with Alkalis

• Fluorine is passed through a cold and

very dilute (2%) sodium hydroxide solution

 oxygen difluoride (OF2) is formed

• When fluorine is passed through a hot

and concentrated sodium hydroxide solution

 oxygen is formed instead

2F2(g) + 4NaOH(aq)

0 –2

hot, concentrated

→ 4NaF(aq) + O2(g) + 2H2O(l) –1 0

Reactions with Alkalis

• Chlorine

 reacts with cold and dilute sodium

hydroxide solution to form sodium chloride and sodium chlorate(I) (also called sodium hypochlorite )

• Chlorine

 reacts with hot and concentrated

sodium hydroxide solution to form

sodium chloride and sodium

chlorate(V)

3Cl2(aq) + 6NaOH(aq)

0 hot, concentrated

→ 5NaCl(aq) + NaClO3(aq) + 3H2O(l)

Reactions with Alkalis

• Bromine

 undergoes similar reactions with

alkalis as chlorine

 sodium bromate(I) formed is unstable

Reactions with Alkalis

• Sodium bromate(I) formed

 disproportionates to form sodium

bromide and sodium bromate(V)

readily at room temperature and pressure

 reversible

Reactions with Alkalis

Br2(aq) + 2NaOH(aq)

cold, dilute

→ NaBr(aq) + NaOBr(aq) + H2O(l)

Reactions with Alkalis

3NaOBr(aq)

2NaBr(aq) + NaBrO3(aq)

• The chemical equation for the overall

• Iodine

 behaves similarly as bromine

Reactions with Alkalis

• Except that the reaction with a cold and

• The backward reaction

 often used to prepare standard iodine

solution for iodometric titrations

Reactions with Alkalis

• Dissolving a known quantity of potassium

iodate(V) in excess potassium iodide solution and dilute sulphuric(VI) acid

 generated a known amount of iodine

solution

KIO3(aq) + 5KI(aq) + 6H+(aq)

→ 3I2(aq) + 3H2O(l) + 6K+(aq)

Reactions with Alkalis

• The iodine generated

 used to oxidize reducing agents

 such as sulphate(IV) ions and

ascorbic acid (vitamin C)

Reactions with Alkalis

• Excess iodine

 can be determined by back titration

with sodium thiosulphate solution

I2(aq) + 2S2O32–(aq)

→ 2I–(aq) + S4O62–(aq)

Reactions with Alkalis