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PERIODIC TABEL
Development of the periodic table ; In 1864 the English chemist John Newlands noticed that when the
elements were arranged in order of atomic mass, every eighth element had similar properties.
Newlands referred to this peculiar relationship as the law of octaves. However, this “law” turned out to
be inadequate for elements beyond calcium, and Newlandss work was not accepted by the scientific
community.
In 1869 the Russian chemist Dmitri Mendeleev and the German chemist Lothar Meyer independently
proposed a much more extensive classification of elements. Both these classifications emphasized the
periodicity of the properties of the elements with their atomic masses. While Mendeleev’s
classification was based largely on consideration of chemical properties of elements that of Lother
Meyer emphasized physical characteristics of the ELEMENTS. Lother Meyer presented the
classification in the form of curve while Mendeler in the form of a table.
Mendeleev’s Periodic Law : The properties of the elements are a periodic function fo their atomic
mass. i.e. when the elements are arranged in order of their increasing atomic mass, elements with
similar properties are repeated after certain regular intervals.
In 1913 a young English physicist, Henry Moseley discovered a correction between what he called
atomic number the frequency of x-rays generated by bombarding an element with high energy
electrons.
v  a  Z  b
He thereby showed that the atomic number is a more fundamental property of an element that its
atomic mass. Mendeleev’s Periodic law was therefore according modified. This is known as modern
periodic law and can be stated as:
The physical and chemical properties of the elements are periodic functions of their atomic numbers.
i.e. if the elements are arranged in order of their increasing atomic number, the elements with similar
properties and repeated after certain regular intervals.
These regular intervals being 2, 8, 8, 18, 18 and 32. These number are also known as magic number.
THE LONG FORM OF THE PERIODIC TABLE or MODERN PERIODIC TABLE or BOHRIS
TABLE, Modern periodic table is constructed on the basis of repeated electronic configuration of the
atoms. This table contains 7 horizontal rows and 18 vertical coulmn’s. Horizontal raws are known as
period and vertical column are known as group. Periodic table also divided in Blocks on the basis of
the last electron enters in the orbital.
If last electron enter in s-orbital  s-Block
If last electron enter in p-orbital  p-Block
If last electron enter in d-orbital  d-Block
If last electron enter in f-orbital  f-Block
Period : There are seven horizontal rows called periods
1st period contains  2 element (shortest period)
2nd period contains  8 element (shortest period)
3rd period contains  8 element (shortest period)
4th period contains  18 element (longest period)
5th period contains  18 element (longest period)
6th period contains  32 element (longest period)
7th period contains  24 element (incomplete period)
SOME COMMONLY USED TERMS:
1. Inert gases or Noble gases : Element of group 18 are called noble gases. The s and p subshell of the
outermost shell are completely filled in these element
2. Representative or Normal element : All the s and p block elements are known as representative
elements except zero group. (outer most shell are incomplete)
3. Transition element : All the d-block element except 12th group (last two shells are incomplete)
4. Inner transition element : All the f-block elements or 4f and 5f block element are called inner
transition elements. (belongs to 3rd group and last three shells are incomplete).
5. Typical element : Elements are 3rd and 2nd periods are known as typical elements.
Prediction of period, group and Block of a given element : Write the electronic configuration of the
element 1s2 2s2 2p6 3s2………. (A – 1) d’0ns2np2
Period : n (outermost shell no).
Block : On the basis of entering of last electron inorbita.
If electron enter in s-orbital s-Block
If electron enter in p-orbital p-Block
If electron enter in d-orbital d-Block
If electron enter in f-orbital f-Block
Group :
(i)
s-Block : no of electron present in valance she
(ii)
p-block : 10 + no of electron in valance shell
(iii)
d-block : no of electron present in ns and (n – 1) d-orbital
(iv)
f-block : 3rd group
4f-series : (58 – 71)  Lanthanoids
5f-series : (90 – 103)  Actinoids
s-Block element
(i) General electronic configuration : ns1–2
(ii) They are soft metal having low melting and boiling point.
(iii) they have low ionization energies and out as a strong reducing agent.
(iv) They mostly form ionic compound.
P-Block elements :
(i) General electronic configuration : ns2 np1–6.
(ii) p-block element include both metal and non metal. The metallic character increases as we go down
the group and non metallic character increases as we move along aperiod from left to right in this
block.
(iii) They form mostly covalent compound.
d-Block element
General electronic configuration  n 1 d110ns12
3d-series – Sc(21) to Zn(30)
4d series – Y(39) to Cd(48)
5d series – La(57), Hf(72) to Hg(80)
(ii) They are hard metal and good conductor of Heat and electricity
(iii) They show variable oxidation state.
(iv) They form both covalent and ionic compound
(v) Their compounds are generally coloured and paramagnetic
f-Block element
General electronic configuration : (n–2)f1-14(n–1)d00r1 ns2
4f series – Lanthanides – 14 elements from Ce(58) to Lus
5f series – Actinides – 14 elements from Th(90) to Lw(103)
They are heavy metals
Most of the elements of the Actinide series are radioactive
PERIODIC PROPERTIES
In a periodic from left to right there is a regular change in electronic configuration of elements. In a
group from top to bottom the outermost shell electronic configuration is similar. Since the chemical
properties of the elements depends upon their electronic configuration so there is a regular change in
chemical properties in a period and in a group elements have similar chemical properties.
ATOMIC RADIUS
The radius of an atom is the distance between the centre of its nucleous and electron in the last orbit.
According to the Heisenberg’s uncertainty principle the position of a moving electron can not be
accurately determined. So, the distance between the nucleous and the outermost electron is uncertain.
The distance from the centre of the nucleous to the point upto which the density of the electron charge
cloud is maximum.
Different type of radius
Covalent Radius : It is defined as half of the distance between the two nuclic of two like atoms bonded
together by a single covalent bond. The covalent radius rA of atom A in a molecule A2 may be given
as:
rA 
dAA
2
figure
vander waals radius >> covalent radius.
Ionic Radius : It is effective distance from the nucleus of an ion upto which it has its influence on its
electron cloud.
Important general terms :
Isoelectronic : The species which have same no of electron.
Ex :- C4 , N3, O2 , F , Ne, Na  , Mg , Al5
Isotopes :- The atoms which have same atomic number and different mass number.
Ex :- 40, 1 H1, i H2 , H3
Isobar : The elements which have same mass number but different atomic number.
Ex:-
18 Ar
40
, 19K40 ,
40
20C
Isotones : The element which have same number of newtron
Ex :-
30
31
14 Si , 15P
Isodiapher : Atoms having same isotopic number are called iodiapher.
Isotopic Number = no. of newtron – no of proton = n – p = difference of neutron and proton.
Ex :
235
, 90Th231
92 U
Isosters : The molecules which have same number of atoms
Ex :- (i) N2 , Co, Cao, MgS
(ii) CO2 , 5O2 , H2O
General trend of different atomic radius :
Vander waals radius > Metallic radius > Covalent radius
Factors affecting Atomic radius
(i)
Atomic radius  no of shells
(ii)
Atomic radius 
1
effective nuclear charge (Zeff )
Concept of effective nuclear charge
Electron present in lower energy orbitals, tend to reduce the effect of nuclear charge on the electrons
present in higher energy orbitals, tend to reduce the effect of nuclear charge on the electrons present in
higher energy orbitals is knowns as screening effect. The reduce charge felt by the electron is known
as the effective nuclear charge. (Zeff = Zactual – s)
S is the screening or shielding constant whose magnitude determines by the Slater’s Rule.
SLATER’S RULE
(1) The various orbitals are grouped as follows.
(1s) (2s, 2p), 3s, 3p), (sd), (4s, 4p), (4d, 4f), (ss, sp)….
(2) For an electron in a group of s.p electrons.
(i)
No contribution from any electron present in a group of orbital lying on the right side of the
group in which the electron is present.
(ii)
A contribution of 0.35 from every other electron present in r groups of orbital under
consideration. A contribution of 0.30 from the other electron in 1s orbitals if the electron for
which s is : be calculated belong to is orbital.
(iii)
A contribution of 0.85 per electron from all electron presser in (n–1)th shell
A contribution of 1.0 per electron from all the electrons pres in the (n–2)th and the next inner
shell.
(iv)
3. For an electron in a group of d and f. The rule : and 2(ii) apply as such. However the rules 2(iii) and
(2i) are replaced by the rule that the contribution per electrons from all electrons in the inner shells is
1.0.
Shielding constant of different orbital in same shell as follows:
s > p > d > f.
Ex:- What is the Zeff for a 2p electron of a nitrogen
Sol : -
7N
1s 2s , 2p 
2
2
3
s  .35 4  2 85  3.1
Zeff = Zactual S = 7 – 3.1 = 4.9
Ex : - (a) What is the zeff for a 3d electron of a
26 Fe
atom
1s 2s 2p , 3s 3p, 3d , 4s 
2
2
6
2
6
2
S  5 35 181  19.75
Zeff = 26 – 19.25 = 6.25 for 3d electron
(b) Calculate the Zeff for us electron of a
26 Fe
atom
s  1.35 14.85 101.0  22.25
Zeff = Zcactal – s = 26 – 22.25 = 3.75
General variation of Zeff :
(i)
Zeff of contain is greater than the zeff of parentatom and Zeff of anion is less than that of
paraent atom.
(ii)
Zeff of A+ > Zeff of A > Zeff of A–
(iii)
Zeff of A–3 < Zeff of A2– < Zeff of A– < Zeff of
(iv)
In general zeff in creases moving in a period from left to right to right.
Zeff of Li < Be < B < C < N < O < F
(v)
In a group moving from top of bottom zeff almost constant.
Application of slater’s Rule
(i) A us orbital is filled earlier than a 3d orbital. K : 1s2, 2s2, 2p6, 3s2, 3p6, 3d1, 19…………………..4s1
Zeff for 3d : 19 – 18 = 1, zeff for us = 19 - .85 × 8 – 10 × 1 = 2.2
Thus, a us electron would be subjected to a greater electrostatic attraction than a 3d electron. Hence 4s
orbital is filled earlier than a 3d-orbital.
(ii) The transition metal atoms lose ns electron first during ionization consider the case of 27Co.
1s 2s 2p y, 3s .3p , 3d , 4s 
2
2
6
2
6
7
2
Zeff for 4s electron = 27 – {1 ×–35 + 15 × .85 + 10 × 1.0) = 3.9
Since 4s electron felt less electrostatic attraction then the 3d electron. So removal of electron from 4s
orbital is easier than the 3d orbital.
Periodicity in Atomic Radius and Ionic Radius :
1.
For normal elements:
(a) In a period from left to right Zeff increases and no of shells are constant. So the atomi radius
decreases.
Li
be
B
C
N
O
F
Zeff.
1.3
1.95
2.6
3.25
3.9
4.55
5.2
radius
1.23Å 0.89Å 0.80Å 0.77Å 0.74Å 0.74Å 0.72Å
(b) In a group moving from top of bottom number of shells increases and Zeff remains almost same.
So, the atomic size increases.
Zeff
Li
Na
K
Rb
Cs
1.3
2.2
2.2
2.2
2.2
1.57Å
2.03Å
2.16 Å
2.55Å
Radius 1.23Å
2.
The atomic radius of inert gas (zero group) is shown largest in a period because of its Nander
vander waals radius.
3.
For Transition element : In the first transition series, the atomic size slightly decreases from Sc
to Mn because effect of effective nuclear charge is stranger than the shielding effect. The
atomic size from the Fe to Ni almost remains the same because the both the effects balance
each other. The atomic size Cu to Zn slightly increases because shielding effect is more than
the effective nuclear charge due to d10 structure of Cu and Zn.
Sc
Ti
1.44Å 1.32
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
1.22
1.18
1.17
1.17
1.16
1.15
1.17
1.25Å
However, in vertical column transition element, there is in increase in size from 1st member to IInd
member as expected but from IInd member to IIIrd member, there is very small change in size and
some times size are same. This is due to Lanthanide contraction.
Cause of Lanthanide contraction : In lanthanide the additional electron enter in (n – 2)f orbital. The
shielding effect of (n – 2)f electron is very small because the shape of f-subshell is very much diffused.
Therefore zeff increases and the size slightly decreases.
Example : Some trend
(1) Li < Na < K < Rb < Cs < Fr
(2) Li > Be > B > C > N > O > F
(3) C–4 > N–3 > O–2 > F–
(4) Ne > Li > B > C > N
(5) Ar > Ca+2 > Sc+3
(6) I– > I > I+
(7) Fe > Fe+2 > Fe+3
(8) Cr < Mo  W
Ionisation Potential or Ionisation Energy
The minimum amount of energy required to remove the most loosely bound electron from an isolated
atom in the gaseous state is known as ionisation potential or ionisation energy of the element.
E
M 
 M  e
IE  E
Energy required for the removal of first, second and third electron from the gaseous atom is called
first, second, and third ionisation energy respectively.
Mg  IE1 
 Mg  e

Mg  IE2 
 M
g  e

M
 M
g   IE3 
g  e
The order of 1st, second and third ionisation energy may be given as:
FACTOR AFFECTING THI I.P.
(i) IE or IP  Zeff (effective nuclear charge)
(ii) IE or IP 
1
size
(iii) Effect of Half filled and fully filled orbitals the atoms having half filled and fully filled orbitals
are comparatively more stable. In case of half filled orbital there is change of one position (along the
period) and in case of fully filled orbital the IP or IE is maximum along the period.
Trend of IE : (i) In general, In a group the IP or IE decreases form top to bottom due to increment of
size.
(ii) In a period, the value of IP increases from left to right (with breaks, where the atom having stable
configuration) due to increase in zeff and decrease in size.
Some trends :
(1) Li > Na > Rb > Cs > Fr
(2) He > F > N > O > C > Be > B > Li
(3) B > Al < Ga > In < Il
800 577
579
558
589
(4) Fe < Fe+ < Fe+2
(5) I+ > I > I–
(6) Ar > Cl > P > S > Si > Al
(7) N–3 < O–2 < F– < Ne
(8) Li+ > Ne+ > O+ > F+ > N+ > B+ > C+ > Be+ > C+ > Be+ (IE2)
Application of Ionisation potential :
(i)
The elements having low values of ionisation potential readily lose their valency electron
and act as electropositive elements. These elements form cations and ionic compound.
(ii)
The elements having low values of IP act as stron reducing agents.
(iii)
The elements having low values of IP are basic in character.
ELECTRON AFFINITY
The amount of energy release when an electron is added to the most shell of or mole of an isolated
gaseous atom in its lower energy start Ag  e  Ag  E1
Electron affinity is actually the measure of tightness of holding of an extra electron to an atom. The
addition of an electron to a neutral atom is an exothermic process. The first electron affinity is given a
negative sign. The addition of a second electron to a monovalent anion as to make A2– is difficult
because both have same charge and experience a lot of repultion. The second electron affinity is,
therefore, an endothermic process and will be given a positive sign.
FACTOR AFFECTING ELECTRON AFFINITY
(i)
E.A  Zeff
(ii)
E.A  1/size
(iii)
It also depend upon electronic configuration. In generator half filled or full filled subshell
show zero EA.
TREND OF ELECTRON AFFINITY
(i)
In general EA value increases in moving form left to right in period because zeff increases
and size decreases
(ii)
In a group moving from top to bottom the EA value of elements decreases because the
atomic size increases.
Exception :- IIIrd period non metal show a greter EA in comparison with IInd period element due
to high electron density present at the IInd period non metal.
IMPORTANCE OF ELECTRON AFFINITY
(i)
The element having high values of electron affinity are capable of accepting electron
easity. They form anions and electrovalent compound with electropositive element.
(ii)
The element having high value of electron affinity act as a strong oxidizing agent.
(iii)
The elements having high values of EA are acidic in nature.
SOME TRENDS :
(i)
(ii)
F>O>C>B>N
Cl > F > Br > I
–349 –328 –325 –295
(iii)
S>O
(iv)
P>N
ELECTRONEGATIVITY :
It may be defined as the tendency of an atom to attract shared pair of electrons towards itself in a
covalently bonded molecule. It is important to note that EA and EN both measure the electron
attracting power but the former refers to an isolated gaseous atom while the latter to an atom in a
covalent molecule.
Electronegativity Scales
(i)
Paciling’s scale : If XA and XB are the electronegativity of atoms A and B respectively then
XA  XB  0.208 
 = actual bond energy –
Ex :-
EAA  EBB
Calculate the electronegativity of fluorine from the following data.
EHH  104.2 Kcal / mol1 , EFF  36.6KCal / mole , EHF  154.6K
Cal/ mole1 , X  2.1
Sol.
  EHF  EHH  EFF  134.6  104.2  36.6
  72.85 kcal / mole1
Now XFXH  0.208   0.208 72.85  1.775
XF  2.1  1.775
(ii)
 XF  3.87
Mulliken’s approach (scale) : According to Mullike the EN of an element is the average
value of its IP and EA.
X
EA  IP
, IP, and EA iness
2
but value of measured by mulliken is 2.8 times larger. Hence
X
fE  EA EE  EA

, IE and EA in eV
2  2.8
5.6
Ex. :
IP and EA of F are respectively 17.418 and 3.45 Cu. Calculate XF .
Sol.
XF 
(iii)
IE  EA 17.418  3.45

 3.72
5.6
5.6
Allred and Rochow’s approach (scale)
X
0.359 Zeff
 0.744
r2
Zeff = Z – r
r = covalent radius in Å
Ex : Calculate the x of c atom using allred Rochow method. The covalent radius of c atom is 0.77Å.
Sol. Zeff = Z – r = 6 –  0.35  4  0.85  2  2.9
xC  0.359 
2.9
 0.744  2.5
2
0.77
 
(i)
EN  Zeff
(ii)
EN  1/size
(iii)
EN sp orbital > EN sp2 orbital > EN of sp3
Trends of Electronegativity
(i)
In period, EN increases from left to right due to increase in Zeff and decrease in size
(ii)
In Group electronegativity decreases as we move from top to bottom due to increases in
size
Importance of EN : Following prediction can be made from values of EN.
(i)
Nature of bond.
(a) when xA  xB  0, then bond is purely covalent
(b) xA  xB  1.7 then bond is 50% covalent and 50% ionic
(c) xA  xB  1.7 then bond is ionic
(2) Percentage Ionic character may be calculated as
IC = 16 xA  xB  3.5  xA  xB 
2
(d)