p-BLOCK ELEMENTS
Group 15 Elements – Nitrogen Family
Members
of the family: Nitrogen is the first member of the group VA. Other members of
this group are phosphorus (P), arsenic (As), antimony (Sb) and bismuth (Bi).
Nitrogen
is a non-metal. The non-metallic character of the VA group decreases on going
down the group:
1.Write
the general outer electronic configuration of group 15 elements
General
outer electronic configuration is [NG]
ns2 np3 .
2. What are the common oxidation
states shown by these elements?
These
elements have five electrons in the valence shell. These elements can gain
three electrons to complete their octets. Hence show -3 oxidation state.In
addition to - 3 oxidation state, the elements of group 15 exhibit +3 (due to
inert pair effect) and +5 oxidation states (by losing all 5 electrons). For
e.g., phosphorus forms pentahalides such as PF5, PCl5 (+5
oxidation state) and trihalides PCl3, PF3 (+3 oxidation
state).
3. Nitrogen is chemically inert at room temperature
while P is highly reactive
N2 is
a triply bonded molecule and the bond dissociation energy of this molecule at
room temp is quite high. Its
small atomic size.
P on the other hand has only P-P
single bond which is comparatively weaker.
4. The ionization energies of these elements is quite high
The larger ionization energy is due to - greater
nuclear charge, small size and stable half filled configuration of the atoms npx1,
npy1, npz1
5. The electronegativity values of elements of
group 15 are higher than the corresponding elements of group 14.
The
elements of group 15 have smaller size and greater nuclear charge of atoms and
therefore they have higher electronegativity values.
6. Nitrogen can’t show +5 state while P can. or
PCl5 exists but NCl5 doesn’t. Why?
P has
vacant d orbital due to which it can extend its covalency to +5 and hence can
form pentahalides. N on the other hand doesn’t have vacant d orbitals in its
valence shell.
Nitrogen
can’t show catenation but phosphorus can.
Nitrogen atom is quite small, &
can satisfy its valency by forming pπ-pπ multiple bond and thus do not show catenation property.
Phosphorus and rest of the members of the family cannot be linked by multiple
bonds due to their comparatively bigger size and weaker bond dissociation
enthalpy& they undergo catenation and exists as poly–atomic molecules.
Ammonia has
tetrahedral shape but doesn’t have bond angle of 109028’.
Ammonia has 3 bp and 1 lp. Due to
bp-lp repulsion the bond angle reduces from 109028’ to 107.80
In earth crust
nitrogen occurs as chile salt petre and Indian salt petre. Write the
formula of both compound.
Chile Saltpetre – NaNO3 Indian Saltpetre – KNO3.
On moving
from As to Bi in group 15 only a small increase in covalent radius is observed.
Due to presence of completely filled d or f orbitals which have poor
shielding effect and so the outer electrons experience stronger nuclear
attraction.
Nitrogen
exists as a diatomic molecule while phosphorus exists as tetra atomic.
N has tendency to form pπ
- pπ multiple bonds due to its small size but P being a
bigger atom cannot form pπ - pπ multiple
bonds effectively and depends on two other P atoms to fulfil its
covalency.
Nitrogen has
maximum covalency of four.
Due to availability of only one s
& three p orbitals for bonding and no vacant d- orbitals in its outermost shell.
PH3
has lower b-pt. than NH3. Why?
PH3 molecules do not form
intermolecular H-bonding due to less electronegativity of P while NH3
molecules form intermolecular H-bonds as N is
more electronegative in nature and thus PH3 has lower b-pt.
than NH3.
PH3 is a
weaker Lewis base than NH3
Both PH3 and NH3 have a lone pair of
electron. But since the P atom is larger in size than the N the electron
density over P is less compared to that in N.
NH3 is water
soluble but PH3 is not.
NH3 can form strong intermolecular H
bonding with water as N is more EN than P
PCl3 fumes in
moist air.
PCl3 when comes in contact with
moisture it forms HCl
PCl3
+ H2O
→ POCl3 + HCl
PCl5
is ionic in solid state Or solid PCl5 is
ionic compound but in the liquid state and gaseous states it is covalent
Because it exists as [PCl4]+ [PCl6]-
in solid state ie, in combination of 2
ions. This is not possible in liquid or gaseous states
PCl5 is
highly unstable or PCl5 is a good chlorinating agent. Why ?
PCl5 has trigonal bipyramidal shape
with bond angles of 900 and 1200.
PCl5 gets decomposed
reality on heating to liberate Cl2, thus act as a chlorinating
agent,
H3PO2
is a better reducing agent than H3PO3.
In H3PO2, two
H atoms are bonded directly to P atom, while in H3PO3,
only one H atom is bonded directly to P
atom. More the number of H atoms bonded
directly to P atom, more better reducing agent the oxy acid is.
NH3
is a good complexing agent. Why ?
Due to its small size, presence
of lone pair of e- on N and
due to greater charge density
NO2
is coloured but its dimer N2O4 is colourless.
NO2 has an unpaired electron and it thus absorbs
light from visible region and appears coloured. N2O4 does
not have any unpaired electron and thus is colourless.
NO becomes
brown when released in air.
NO reacts with O2 to form
NO2 when released in air. NO2 so formed is brown in
colour.
Arrange the
following in decreasing order of the property indicated.
a) NH3, PH3,
ASH3, SbH3, (basic
strength).
NH3 >
PH3 > AsH3 >
SbH3.
What happen when NH3/aq.NH3
is added to
a)
White
ppt of AgCl.
AgCl(ppt) + 2 NH3 →
[ Ag (NH)2 Cl ]
Colourless
b)
aqueous
FeCl3 solution.
FeCl3 (ag)
+ 3NH4OH (ag) → Fe (OH)3 +
3NH4Cl
Brown ppt.
c)
aqueous
CuSO4 solution.
CuSO4(ag)
+ 4NH3(ag) → [ Cu(NH3)4] SO4.
(
pale blue )
deep blue solution.
Explain the chemistry of brown ring test for
nitrates.
In this test, aqueous ferrous sulphate
solution is added to the aqueous solution of the salt containing nitrates and
then concentrate H2SO4
is added along the sides of the test
tube. The brown ring formed at the junction of the solution and H2SO4
layers indicate the presence of nitrate ion in the solution. In this test Fe2+
ions reduce nitrates to nitric oxide which reacts with Fe2+ to form
a brown coloured complex.
NO3- + 3Fe2+ + 4H+ →
NO + 3Fe3+ + 2H2O.
[Fe (H2O)6]2+ +
NO → [ Fe (H2O)5 NO ] 2+ + H2O.
(brown ring
complex.)
Catenation
The
elements of group 15 also show a tendency to form bonds with itself known as
catenation. All these elements show this property but to a much smaller extent
than carbon. For e.g., hydrazine (H2N-NH2) has two N
atoms bonded together, hydrazoic acid, (N3H), has three N-atoms,
azide ion,
, has also three N atoms bonded together, while diphosphine (P2H4)
has two phosphorus atoms bonded together. The lesser tendency of elements of
group 15 to show catenation in comparison to carbon is their low (M-M) bond
dissociation energies.
|
Bond
|
C – C
|
N - N
|
P - P
|
As - As
|
|
Bond energy
|
353.3
|
163.8
|
201.6
|
147.4
|
Ammonia has Hydrogen
bonding
Nitrogen
in its compounds with hydrogen shows hydrogen bonding though to a lesser extent
than oxygen. For example, ammonia (NH3) shows intermolecular
hydrogen bonding.
Oxides of Nitrogen
They
range from N2O (oxidation state of nitrogen +1) through NO, N2O3,
NO2, N2O4 to N2O5 in
which the oxidation state of nitrogen is from +2 to +5.
Structures of the Oxides of
Nitrogen
Manufacture of Ammonia : Haber's
process
The
manufacture of ammonia by Haber's process involves the direct combination of
nitrogen and hydrogen.
Low temperature
The
temperature should be remain as low as possible, (although at unusually low
temperatures, the rate of reaction becomes slow). It has been found that the
temperature, which optimizes the yield of ammonia for the reaction, is maximum
at about 500°C.
High pressure
Since
Haber's process proceeds with a decrease in volume, it is favored by high
pressure. In actual practice, a pressure of 200 - 900 atmospheres is employed.
Catalyst
A catalyst is usually
employed to increase the speed of the reaction. Finely divided iron containing
molybdenum or alumina is used as a catalyst. Molybdenum or alumina (Al2O3)
acts as a promoter and increases the efficiency of the catalyst. A mixture of
iron oxide and potassium aluminate has been found to work more effectively.
Explain
the structure of ammonia
Ammonia is a covalent molecule as
is shown by its dot structure. The ammonia molecule is formed due to the
overlap of three sp3 hybrid orbitals and orbitals of three
hydrogens. The fourth sp3 hybrid orbital is occupied by a lone-pair.
This gives a trigonal pyramidal shape to ammonia molecule. The H-N-H bond angle
is 107.3°, which is slightly less than the tetrahedral angle of 109°28. This is
because the lone pair - bond pair repulsions tend to push the N-H bonds
slightly inwards. In liquid and solid states, ammonia is associated through
hydrogen bonds.
Why is
ammonia a strong Lewis base
Ammonia
molecule has a strong tendency to donate its lone pair of
electrons of nitrogen to other
molecules. Thus, it acts like a strong Lewis base. In aqueous solutions, NH3
ionizes in accordance with the reaction.
Explain the Industrial
Preparation of nitric acid
On a commercial
scale, nitric acid is manufactured through the Ostwald's process - the process
of catalytic oxidation of ammonia.
Step1: Oxidation
of ammonia to nitric oxide
Ammonia
is oxidized by air in the presence of Pt catalyst at 800°C to give nitric
oxide.
Step 2: Oxidation
of NO to NO2
The
nitric oxide is oxidised by air at temperature below 100°C, to give nitrogen
dioxide (NO2)
Step 3:Formation
of nitric acid
Nitrogen dioxide
is then converted to nitric acid by absorbing NO2 in water, in the
presence of air.
Draw the structure
of nitric acid
Nitric acid is a monobasic
acid and may be structurally represented as,
Explain the structure
of nitrate ion
The
nitrate ion, NO3- is isoelectronic with carbonate ion, CO32-.
It is also a planar ion. The nitrate ion can be represented by the resonance
structures shown below:
·
What is aqua-regia?
A mixture
of concentrated HCl and concentrated HNO3 (3 : 1 by volume) is
called aqua-regia. It can dissolve noble metals like gold and platinum.
NO2 exists as a dimer N2O4
NO2 has an odd electron
over it which makes it unstable. Hence it dimerises.
NO2
is coloured but its dimmer N2O4 is colourless.
NO2 has an unpaired electron and it thus absorbs
light from visible region and appears coloured. N2O4 does
not have any unpaired electron and it thus is colourless.
NO becomes brown when released in air.
NO reacts with O2 to form
NO2 when released in air. NO2 so formed is brown in
colour.
Group 16 Elements – Oxygen Family
Oxygen
Group 16 Elements- Oxygen(O), Sulphur(S),
Selenium(Se), Tellurium (Te) and polonium (Po).
General Electronic Configuration: [Noble
gas] ns2np4
Oxygen is highly electronegative in
nature. It shows oxidation state of -2.
Sulphur and the rest of the elements
are less electronegative and thus, exhibits
+2 oxidation state in their compounds. Their atoms have also have vacant
d- orbital in their valence shell. As a result they can also show +4 (due to
inert pair effect) and +6 oxidation state.
Oxygen
gas exists as diatomic molecule. So, gas is also called dioxygen. It also
occurs in three isotopic forms such as
The important
physical values of atomic and molecular oxygen are listed below.
1. H2O
is a liquid while H2S is gas.
Oxygen being more
electronegative than Sulphur can form strong intermolecular H –
bonding in H2O, it is liquid.
2. Sulphur in
vapour state exhibits paramagnetism.
Sulphur in vapour state
exists as S2 molecule. In this state it has two unpaired electrons in its antibonding π* 3py
orbital and thus is paramagnetic.
3. SF6
is not easily hydrolysed
Because S is hexacovalent (+6 state) and for hydrolysis
it has to form a co- ordinate bond with water. But since it cannot further expand its covalency beyond 6 it cannot undergo
hydrolysis.
4. SF6 is less reactive than SF4.
This is so because S atom
in SF6 is more sterically
hindered than in SF4.
Moreover SF4 contains lone pair of electrons on S atom that makes
it more reactive
5. Mention
the conditions to obtain maximum yield of H2SO4 by contact process.
i)
Low temperature of about 700 K. ii) High pressure of 2 atmospheres.
iii)
Use of V2O5 as a catalyst.
6.
9. Arrange
the hydrides of group 16 in decreasing
order of the property indicated.
a) boiling point: H2O >
H2Te > H2Se >
H2S.
Oxygen is the most EN element in gp-16 and
hence can form strong intermolecular H- bonding. Other elements being less EN
can’t form this. The bp of other hydrides depend on the van der Waal’s forces
acting between them. As the size of atom increases this van der Waal’s forces
also increases.
b)
thermal stability: H2O >
H2S > H2Se >
H2Te
As the
size difference between H and the element M increases the bond between
them becomes
Weak.
c) acidic character: H2O
< H2S <
H2Se < H2Te.
As the size
difference between H and the element M increases the bond between them becomes
Weak
and the bond breaks easily releasing H+ ions
d) reducing character:
As the size
difference between H and the element M increases the bond between them becomes
Weak
and release of H becomes easier.
10. H2S
acts only as a reducing agent but SO2 acts both as reducing agents
as well as oxidizing agent.
In H2S , S is in the oxidation state of -2 and it can only
increase its oxidation state either to +4 or +6 by losing electrons and thus
acts only as reducing agent. In SO2,
S is an oxidation state of +4 and its can undergo oxidation to +6 or
reduction to -2 state and thus act as
both reducing as well as oxidizing agent.
11. SF6
exists but SH6 and SCl6 doesn’t.
F is the most EN element and hence can
oxidise S to +6 state by promoting it’s electrons to d- orbital. H and being less EN can’t oxidise S.
12. SF6 exists but SCl6 doesn’t.
Due to small size
of S, six large Cl atoms cannot be accomodated around S atom. But small six F
atoms can be
accomodated. Moreover Cl being
less EN can’t oxidise S to +6 state.
13. SO2 has zero dipole moment.
SO2 has a trigonal planar structure
with bond angle of 1200. The individual S-O bond dipoles get cancelled
by the resultant of other two dipole moments.
14. Why is
sulphuric acid highly viscous?
Due to intermolecular H bonding
15.Draw the
structures of: H2SO4, H2SO3, H2S2O6, H2S2O7, SF4.
Group 17 Elements – Halogen Family
Halogens are present in group 17 of
the p-block. There are five elements which are named as fluorine(F),
chlorine(Cl),bromine(Br),iodine(I),and astatine(At).
General Electronic Configuration: ns2p5.
Common oxidation state is -1. Other
than this they can show +1, +3, +5 and +7 states too.
|
Element
|
Atomic
Radius (Aº)
|
Ionic
Radius (Aº)
|
Ionization
Energy (kJ mol-1)
|
Melting
point (K)
|
Boiling
point (K)
|
Electron
affinity
|
Electro
negativity
|
|
F
|
0.72
|
1.86
|
1681
|
53
|
85
|
332.6
|
4.0
|
|
Cl
|
0.99
|
1.81
|
1255
|
172
|
238
|
348.5
|
3.0
|
|
Br
|
1.14
|
1.95
|
1142
|
266
|
332
|
324.7
|
2.8
|
|
I
|
1.33
|
2.16
|
1007
|
386
|
456
|
295.5
|
2.5
|
1.
All the halogens are
coloured in nature.
Due to absorbtion of energy from the
visible light by the molecules for the excitation of outer non-bonded electrons
to higher energy levels.
|
Halogen
|
Fluorine
|
Chlorine
|
Bromine
|
Iodine
|
|
Colour
|
Light
yellow
|
Greenish
yellow
|
Reddish
brown
|
Dark
violet
|
The amount of energy required for
excitation depends upon the size of the atom. Fluorine atom is the smallest and
the force of attraction between the nucleus and the outer electrons is very
large. As a result, it requires large excitation energy and absorbs violet
1ight (high energy) and therefore, appears pale yellow. On the other hand,
iodine needs very less excitation energy and absorbs yellow light of low
energy. Thus it appears dark violet. Similarly, we can explain the greenish
yellow color of chlorine and reddish brown color of iodine.
2.
Bond dissociation enthalpy
of F2 is less than that of Cl2.
The bond dissociation enthalpy for F-F
bond is expected to be more than for Cl-Cl bond. But its actual value is less.
This is because of high inter-electronic repulsions between non-bonding
electrons present on the atoms in the fluorine molecule (F-F).
3.
Halogens can exhibit
positive oxidation states too.
Oxidation State: Except F, all other
elements have vacant d-orbitals in the valence shell to which one two or even
three electrons can be promoted from the valence s and p-orbitals.Hence, F can
show only -1 oxidation state in its compounds while the other halogens exhibits
-1,+1,+3,+5 and +7 oxidation states.
4. Halogens
are good oxidising agents.
Since the halogens have strong
electron accepting tendencies, they are powerful oxidizing agents. The
oxidizing power is also related to the standard reduction potential(Eo)
values. Greater the Eo value, more is the oxidizing nature.
5. HF
exists in liquid state while HCl is a gas
The liquid state of hydrogen fluoride
is attributed to the presence of intermolecular hydrogen bonding in the
molecules due to their highly polar nature.
6.
Arrange
the hydrides of group -17 in the decreasing order of:
a)
Thermal stability: The stability of the hydrogen halides
decreases from HF to HI.
As the
size difference between H and the element M increases the bond between them
becomes weak
b)
Reducing nature: The reducing nature of the hydrogen halides is linked to
their thermal stability.
As
the bond strength weakens the release of H becomes easier. Thus, HF is the
weakest reducing agent, whereas HI has the maximum reducing strength.
c)
Acidic character: The acidic nature of the
hydrogen halides is also linked to their thermal stability. As the bond strength weakens the release of H+
becomes easier. Thus, HF is the weakest acid whereas HI strong acid.
7. Compare
the relative acidic strengths of oxo acids of halogens
(a)
Acid strength of oxoacids of different halogens with the same oxidation state
decreases with the increase in the atomic number.
HClO > HBrO > HIO
More
is the electronegativity of halogen atom greater will be its electron
attracting tendency facilating the release of H+ ion from O-H bond.
Thus, HClO is the strongest acid because the Cl is the most electronegative
among all the halogens.
(b)
Acid strength of the oxo acids of the same halogen atom increases with the
increase in the oxidation number of the halogen atom.
HClO4 > HClO3
> HClO2 > HClO
The order is explained by considering the
relative stabilities of the anions that are formed by the release of H+
ions from the acids in solution.
8.
Halogens have very high Ionization energies
The
ionization energies of halogens are very high. This indicates that they have
very little tendency to lose electrons. However, on going down the group from
fluorine to astatine, the ionization energy decreases. This is due to gradual
increase in atomic size, which is maximum for iodine. Consequently, it has the
least ionization energy in family.
9. Cl has
exceptionally higher electron affinity than F
Fluorine has unexpectedly less electron
affinity than chlorine. Therefore, chlorine has the highest electron affinity
in this group. The lower electron affinity of fluorine as compared to chlorine
is due to very small size of the fluorine atom. As a result, there are strong
inter-electronic repulsions in the relatively small 2p subshell of fluorine and
thus the incoming electron does not feel much attraction. Therefore, its
electron affinity is small. Thus, electron affinity among halogens varies as: F
< Cl > Br > I.
Chlorine has the highest electron
affinity in the periodic table.
Interhalogen Compounds
The
compounds containing two or more halogen atoms are called inter halogen
compounds. Each halogen combines with every other halogen to form interhalogen
compounds. For e.g., ClF, ICl3, BrF5 etc.
They are
of two types:
(i)
Neutral molecules containing two or more halogen atoms. For e.g., ICl, BrF5,
IF5, IF7 etc.
(ii)
Negatively charged interhalogen anions or polyhalide ions such as
The
different (types of) interhalogens of the type AX (diatomic), AX3
(tetra atomic), AX5 (hexa atomic) and AX7 (octa atomic)
are given below:
Some
characteristics of inter halogen compounds are:
(i) They
are covalent compounds.
(ii) They
are more reactive than the constituent halogens. It is because A-X bond is
relatively weaker than X-X bond.
(iii)
They are very good oxidising agents.
(iv)
Their melting and boiling points increase with the increase in the difference
of electronegativity.
(v)
Chlorofluoro hydrocarbons are known as Freons and are used as
refrigerants. For e.g., Freon-11 is CCl3F, Freon-12 is CCl2F2,
Freon-13 is CClF3 etc.
Sub Topics
Shapes
The
molecular structure of interhalogen compounds can be explained on the basis of
VSEPR theory. The structures of a few interhalogens are shown in figure.
Introduction to Noble Gases
The
elements helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon
(Rn) are grouped together in Group 18 of the Periodic Table. All of these are
gaseous under ordinary conditions of temperature and pressure. The last number
of the group, radon is obtained from radioactive disintegration of radium. All
others are present in air in traces. They are also known as rare gases because
they are found in very small quantities in nature. They are highly non-reactive
and do not take part in chemical combinations to form compounds like other
elements and are, therefore, also called inert gases
The noble
gases are a group of chemical elements with very similar
properties: under standard conditions, they are all odorless, colorless, monatomic gases, with very low chemical
reactivity.
The six noble gases that occur naturally are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and the radioactive radon (Rn).
Chemical properties
Neon,
like all noble gases, has a full valence shell. Noble gases have eight
electrons in the outermost shell, except in the case of helium, which has two.
The
noble gases are colorless, odorless, tasteless, and nonflammable under standard
conditions. They were once labeled group 0 in the periodic table because
it was believed they had a valence of zero,
meaning their atoms cannot
combine with those of other elements to form compounds. However,
it was later discovered some do indeed form compounds, causing this label to
fall into disuse.[11]Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior:
|
2
|
2
|
|
|
10
|
2,
8
|
|
|
18
|
2,
8, 8
|
|
|
36
|
2,
8, 18, 8
|
|
|
54
|
2,
8, 18, 18, 8
|
|
|
86
|
2,
8, 18, 32, 18, 8
|
Compounds of xenon
Fluorides of xenon
The important fluorides of xenon are xenon difluoride(XeF2), Xenon tetrafluoride(Xef4) and xenon hexafluoride.
1. Xenon difluoride(XeF2)
It is prepared by heating a mixture of Xenon and fluorine in the ration 2:1 at 400 degree Celsius and 1 bar pressure in a sealed nickel tube.
Xe + F2 ---Ni----> XeF2
XeF2 undergoes hydrolysis when treated with water an d evolves oxygen.
2XeF2 + 2H2O -------> 2Xe + 4HF + O2
In XeF2, Xenon is sp3d hybridised and the molecule has linear structure as shown.
2. Xenon tetrafluoride (Xef4)
It is prepared by heating a mixture of Xe and F2 in the molecular ratio 1:5 at 400 degree Celsius and 6 atm in a sealed nickel tube.
Xe + 2 F2 ---------> XeF4
XeF4 react with water and produces explosive XeO3
6 XeF4 + 12 H2O ---------> 2 XeO3 + 4 Xe + 3O2 + 24 HF
In XeF4, Xenon is in sp3d2 hybridised state and has square planar geometry.
3. Xenon hexafluoride (XeF6)
It is prepared by heating a mixture of xenon and fluorine in the ration 1:20 at 300 degree Celsius and 60 atm in a nicked vessel.
Xe + 3 F2 --------> Xe F6
XeF6 Undergoes slow hydrolysis with atmospheric moisture producing highly explosive XeO3.
XeF6 + 3H2O --------> XeO3 + 6 HF
XeF6 molecule possesses distorted octahedral structure. The Xe atom in XeF6 is in sp3d3 hybridisation.
Oxides of xenon
1. Xenon trioxide (XeO3)
XeO3 prepared by slow hydrolysis of XeF6
XeF6 + 3H2O -------> XeO3 + ^ HF
XeO3 is soluble in water and its aqueous solution is weakly acidic.
XeO3 + H2O <_-_-_-_-_-_-_-_-> H+ + HXeO4-
XeO3 has pyramidal structure in which Xe is in sp3 hybridisation.
2. Xenon tetroxide (XeO4)
It is prepared by treating barium perxenate (Ba2XeO6) with anhydrous sulfuric acid.
Ba2XeO6 + 2H2SO4 --------> XeO4 + 2BaSO4 + 2H2O
XeO4 is highly unstable and has tetrahedral structure
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