Periodicity : Properties of compoundschemistry/Download/period_2.pdfmelting points of compounds such...
Transcript of Periodicity : Properties of compoundschemistry/Download/period_2.pdfmelting points of compounds such...
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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+ -