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Periodicity : Properties of compounds

Physical properties of the oxides and chlorides

As we cross the periods, the melting points of the compounds and their electrical

conductivity in the liquid state generally fall. Both of the trends suggest that the bonding

type is changing from essentially ionic to predominately covalent. The unusually high

melting points of compounds such as beryllium oxide and silicon oxide indicate a

macromolecular rather than a simple covalent structure.

Within any particular group, the compounds of the lower elements tend to be somewhat

more ionic in character. The ease with which the elements form simple positive ions is

inversely related to their ionization energies, which decrease on passing down a group.

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The enthalpies of formation per mole of oxygen or chlorine show that the oxides and

chlorides tend to become less stable with respect to the elements passing from left to right

across the period. Compounds with a large positive enthalpy of formation are

energetically unstable and will often spontaneously decompose, sometimes explosively,

e.g. nitrogen chloride.

2 NCl3(l) Í N2(g) + 3 Cl2(g)

The enthalpy of formation only indicates the energetic stability of a compound in

comparison with its elements. A large negative value does not necessarily mean that it

will be generally unreactive, since it may combine exothermically with other substances.

Thus, although the enthalpy of formation of silicon chloride is -640 kJ mol-1

, it reacts

vigorously with water.

Action of water on the oxides

On passing across the periods, the oxides change in nature from alkaline or basic to acidic.

This trend may be explained in terms of the difference in bonding type in the oxides.

Oxide solubility in water reaction with water

Li2O soluble - alkaline Li2O + H2O Í 2 LiOH

BeO insoluble (amphoteric)

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B2O3 soluble - weakly acidic B2O3 + 3 H2O Í 2 H3BO3

CO2 soluble - weakly acidic CO2 + H2O Í H2CO3

NO2 soluble - acidic 2 NO2 + H2O Í HNO2 + HNO3

OF2 reacts - acidic OF2 + H2O Í O2 + 2 HF

Na2O soluble - very alkaline Na2O + H2O Í 2 NaOH

MgO almost insoluble (basic)

A12O3 insoluble (amphoteric)

SiO2 insoluble (acidic)

P2O5 soluble - acidic P2O5 + H2O Í 2 HPO3

SO2 soluble - acidic SO2 + H2O Í H2SO3

SO3 soluble - strongly acidic SO3 + H2O Í H2SO4

C12O7 soluble - strongly acidic C12O7 + H2O Í 2 HC1O4

The oxides of electropositive metals are ionic, e.g. Na2O. If the oxide is soluble in water

and then the O2-

ion will react to give hydroxide ions, OH- making the solution alkaline.

Insoluble metal oxides also act as bases, e.g. magnesium oxide will accept protons from

acids, forming water.

If the metal ion is small and highly charged, it may also react with water molecules,

releasing protons, i.e. acting as a Bronsted-Lowry acid.

Thus beryllium and aluminium oxides are amphoteric and may act as either bases or acids,

e.g.

Al2O3(s) + 6 HCl(aq) Í 2 AlC13(aq) + 3 H2O(1)

Al2O3(s) + 2 NaOH(aq) + 3 H2O(1) Í 2 NaAl(OH)4(aq)

Non-metal oxides are covalent in character and since the oxygen atom only carries a

slight negative charge it cannot act as a proton acceptor. If soluble in water, non-metal

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oxides hydrate to produce acids, e.g. sulphur dioxide, SO2, may be regarded as the

anhydride of sulphurous acid, H2SO3.

Marcomolecular non-metal oxides such as silicon dioxide, SiO2, are insoluble in water

but will form salts when fused with alkalis,

e.g. Na2O(l) + SiO2(1) Í Na2SiO3(1)

The acidic nature of non-metal oxides is more pronounced in high oxidation states. Since

the central atom carries a greater partial positive charge it accepts a lone pair of electrons

from a water molecule very readily. For example, carbon dioxide, CO2, behaves as a

weak acid, whereas carbon monoxide is insoluble in water and is generally regarded as

neutral.

Action of water on the chlorides

The chlorides become increasingly acidic on passing across the Periodic Table. The

chlorides of Group 1 are essentially ionic and dissolve in water giving virtually neutral

solutions containing hydrated ions, e.g. sodium chloride.

NaCl(s) Í Na+(aq) + Cl

-(aq)

Chloride solubility in water reaction with water

LiCl soluble - weakly acidic

BeCl2 soluble - acidic

BC13 soluble - acidic BC13 + 3 H2O Í B(OH)3 + 3 HCl

CC14 insoluble

NC13 soluble - acidic NC13 + 3 H2O Í NH3 + 3

HOCl

C12O7 soluble - very acidic C12O7 + H2O Í 2 HClO4

ClF soluble - acidic ClF + H2O Í HF + HOCl

NaCl soluble - neutral

MgC12 soluble - weakly acidic MgC12 + H2O Í Mg(OH)Cl + HCl

AlC13 soluble - very acidic AlC13 + H2O Í Al(OH)3 + 3

HCl

SiC14 soluble - very acidic SiC14 + 4 H2O Í Si(OH)4 + 4 HCl

PC15 soluble - very acidic PC15 + H2O Í POC13 + 2

HCl

S2C12 soluble - acidic S2C12 + 2 H2O Í 2 HCl + SO2 + H2S

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As the charge on the cation rises and its size decreases, the solutions become increasingly

acidic. The metal ion attracts the electrons on the water molecule more strongly, making

the release of hydrogen ions more likely.

Non-metal chlorides are essentially covalent but the bonds may have some polarity

depending upon the difference in electronegativity between chlorine and the element

concerned. It is largely this bond polarity which makes non-metal chlorides liable to

attack by water. Silicon chloride, SiC14 dissolves exothermically in water giving strongly

acidic solution. The silicon atom carries a slightly positive charge and can accept a lone

pair of electrons from the oxygen atom on a water molecule into an empty ‘d’ orbital in

its valency shell.

It is interesting to note that not all non-metal chlorides react readily with water in this

way.

Carbon tetrachloride has slightly positively charged carbon but it has no empty d orbital

in its valency shell which can accept a lone pair of electrons from the water molecule.

Hydrides

Elements in the first two short periods also form hydrides which exhibit periodicity in

their nature of bonding and in their chemical behaviour.

Period 2 LiH BeH2 BH3 CH4 NH3 H2O HF

Period 3 NaH MgH2 AlH3 SiH4 PH3 H2S HCl

Ionic Covalent Covalent Polar covalent

with ionic

character

Cl

SiCl

ClCl

HO

H

Cl

SiCl

ClOH

+ HCl

OH

SiCl

ClOH

OH

SiCl

OHOH

OH

SiOH

OHOH

further attack by water molecules

c�c�

c�

c�

c�

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Nature of bonding

Ionic hydrides in Group I are essentially ionic and they have hydride ions, H- in their

compounds. They have the same ionic structure as their ionic chlorides.

Group II and III hydrides are largely covalent. BeH2 and B2H6 form electron deficient

chain structure and dimeric molecule respectively.

Group IV elements especially carbon form a great varieties of covalent hydrides. The

great tendency to form chains and rings form in carbon is known as catenation. Hydrides

of carbon are fairly stable but silicon hydrides are spontaneously inflammable in air to

form more stable SiO2(s).

CH4(g) + 2 O2(g) Í CO2(g) + 2 H2O(g)

SiH4(g) + 2 O2(g) Í SiO2(s) + 2 H2O(g)

Phosphorus also forms covalent hydride called phosphine, PH3. However, it is unstable

due to the formation of weak P-H sigma bond from the overlap of sp3 orbital of the

third quantum shell of phosphorus with an s orbital from the first quantum shell of

hydrogen. (�Hf of PH3 is +5.4 kJ mol-1

) Pure phosphine ignites spontaneously at

temperature above 1500C.

PH3(g) + 2 O2(g) Í H3PO4(g)

The hydrides of nitrogen, oxygen, sulphur and halogens i.e. NH3, H2O, H2S, HF. HCl are

polar covalent molecules because hydrogen atom is much less electronegative than the

atom to which it is bonded.

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All these hydrides have similar structures.

However they have different bond angles

within these molecules. This is due to the

different number of lone pairs and bond

pairs in these molecules and also the

distribution of the bonding electrons around

the central atoms.

Action of water on hydrides

The ionic hydrides in Group I metals produce H- ions that are very strong bases. They

take away the protons from water molecules and the solution becomes alkaline.

Group II and III hydrides also hydrolyse readily in aqueous solution: However, they

produce a less alkaline solution because their hydroxides are less soluble due to their

higher lattice energies.

AlH3(s) + 3 H2O(l) Í Al(OH)3(aq) + 3 H2(g)

Owing to the non-polar nature of C-H bonds in hydrides of carbon, they do not dissolve

nor react with the polar water molecules.

However, silane SiH4 reacts readily with water because its more polar Si-H bonds in the

molecule. Also, the availability of 3d orbital to take up the lone pair of electrons from

water in the transition state lowers the activation energy for the reaction.

Ammonia is a weak base in water because its lone pair takes up a proton from water.

NHH

H

PHH

H

OHH

SHH

F H

Cl H

HO

HH

-

c�

c�

c�H2 + OH

-

OHH

c�H

SiH

HH

c�H

SiH

HOH

� H2

NHH

H

HO

Hc�

c�c�

NHH

H

H

+

+ OH-

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Phosphine shows only feeble basic character because phosphorus is a much larger atom

with a diffuse lone pair of electrons around the molecule.

Water is a neutral solution and it undergoes slight self-ionization.

Hydrogen sulphide is a weak acid in solution. It donates its proton by the ionization of

the weak S-H bonds in the molecule.

H2O(l)

H2S(aq) + H2O(l) ¾ H3O+(aq) + HS

-(aq) ¾

H3O+(aq) + S

2-(aq)

The hydrogen halides are strong acids in solution.

HF(aq) + H2O(l) Í H3O+(aq) + F

-(aq)

HCl(aq) + H2O(l) Í H3O+(aq) + Cl

-(aq)

OHH

c�c�

c�

c�

c�

c�.

OH

H H3O + OH+ -