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periodic table (Autosaved) .pdf



Original filename: periodic table (Autosaved).pdf
Title: periodic classification
Author: GAURAV

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PERIODIC CLASSIFICATION

[BY GAURAV SONI]

Lavoisier’s classification:

Mendeleev’s periodic table:

1.

1.

2.

He was the first scientist who classified the known elements into two
class called as metals and non-metals.
Failed because of metalloids.

2.
Dobereiner’s classification:

3.

4.

Failed because could not arrange all elements in triad.

Newland’s law of octave:
1.
2.
3.

4.

He arranged the known elements in increasing order of atomic mass and
explained that every eight element has properties similar to 1st one.
Like musical notes of octave
Elements:
Sa
Re
Ga
Ma
Pa
Dha
Ni
H
Li
Be
B
C
N
O
F
Na
Mg
Al
Si
P
S
Cl
K
Ca
Failed because this patern was not followed after Ca.

3.
4.

5.

Merits of Mendeleev PT
1.

(History)

2.

He arranged elements with similar properties in a group of 3 elements
named as traid.
He observed that atomic mass of middle element of each triad was
almost equal to arithmetic mean of 1st and last element atomic mass of
the triad.
TRIADS
Li(7)
Na(23) K(39)
Mean= 23
Ca(40)
Sr(88) Ba(137) 88.5
Cl(35.5) Br(80) I(127)
81.25
S(32)
Se(79) Te(127) 79.15

Periodic classification

1.

2.
3.

3.

Lother meyer plotted a curve between atomic volume and atomic mass.
He observed that elements with similar properties were occupying
similar position on the curve.
 Alkali metal occupies peak position
 Alkaline earth metal occupied descending position of the curve
 Halogen occupied ascending position of the curve.
Failed because it was difficult to remember elements of similar
properties on basis of their position on the curve.

Incorrect atomic mass number of elements are corrected by Mendeleev.
Ex, atomic mass of Be atom earlier was 13.5. later on it was corrected to 9.
Atomic mass = valency x equivalent weight [ 2 x 4.5 = 9 ]
By knowing properties of one element of one subgroup, properties of
other elements of same subgroup can be easily predicted.
Tabular form provide systematic pattern to study properties of atoms of
elements.

Demerits of Mendeleev PT:
1.
2.
3.
4.

Lother meyer’s curve:
1.
2.

He arranged the known elements in increasing order of their atomic
masses and he explained that after a regular interval elements with similar
properties are repeated.
His law=> physical and chemical properties of atom of elements are
periodic function of their atomic mass.
It contains 8 groups(vertical columns) and 7 periods(horizontal rows).
First 7 group are further divided into two subgroups A and B.
A contains normal elements => S and P block
B contains transition elements => d block
Later on additional group zero of inert gases was added. Hence modified
form of Mendeleev table contain 9 groups.

5.

He placed H atom at 2 places along with alkali metal and along with
halogen.
He didn’t give position for isotopes.
He didn’t explained cause of periodicity.
Increasing order of elements atomic mass was not found at some place.
Ca(58.9) and Ni(58.6)
Te(127.6) and I(127)
Three elements were placed collectively in eight group.

S BLOCK:

D Block:

1. Last e enters in S subshell of valence shell.
2. General configuration: [noble] ns1-2
3. To detect group number-> calculate no. of electron in last s
subshell. Period number -> value of last shell no.
4. Group 1st = alkali metals and 2nd = alkaline earth metal [ also
called as representative elements]
5. All s block elements are soft metals and have low melting
and boiling point due to weak metallic bond strength.
6. Forms ionic compound.
7. They are kept in organic solvents because they are more
reactive, except lithium.
Na -> kerosene oil
Li -> paraffin wrapper.
8. All these metals imparts color to flame except Be and Mg.
9. These metals act as strong reducing agent due to low
ionization energy. (highly electropositive)
10. Constant oxidation state: alkali metal = +1 and alkaline earth
metal = +2.

1. Last e enters in d subshell of penultimate shell (2nd last shell)
2. Configuration: (n-1)d1-10 ns1-2 [group number = (n-1)de + nse]
3. These elements are also called as transition elements(partial filled d
orbital) except Zn, Cd and Hg.
4. All d block elements are hard metals and have high MP and BP than s
block elements.
5. They have strong tendency to form complex salt. Ex K4[Fe(CN)6] ->
complex salt. (potassium Ferro cyanide)
6. These metals and their compounds can act as catalyst in different
chemical reaction. Ex Ni, Pd, Pt etc in hydrogenation.
7. These metals have variable oxidation state (more than one) due to
less energy gap in ns and (n-1)d electron.
8. D block elements classified into 4 series
3d -> 21Sc – 30Zn.,
4d -> 39Y – 48Cd
5d -> 57La, 72Hf – 80Hg,
6d -> 89Ac, 104Rf – 112Uub

BLOCKS
P BLOCK:

F BLOCK:

1.
2.
3.
4.

1.
2.
3.
4.
5.
6.
7.

Last e enters in p subshell of valence shell
General configuration: ns2np6
Group no. -> ns e + np e + 10
These elements are non-metals mostly or metalloids. Some of
them are metal also. They are also called as representative
elements(except noble gas)
5. P block elements have relative higher metallic strength than s
block elements. Hence they have more MP and BP than S block.
6. They act as a strong oxidizing agent due to high electron affinity.

Last electron in f subshell of anti-penultimate shell(third last shell)
Configuration: (n-2)f14 (n-1)d0or1 ns2
F block elements are metal and mostly radioactive in nature.
They also have complex formation tendency
Elements of lanthanide series are called as rare earth elements.
Divided into 2 series: 4f and 5f
Also called as inner transition elements. (group number = 3)

Three types:

SHIELDING EFFECT (screening effect):

1. Covalent radius: exist b/w non-metal bonded (equal to half of the
value of single covalent BL.
If bond b/w same types. Like H2, CR = d/2 where d = ra + ra
If bond b/w atoms of diff type: like AB,CR = d/2 where d = ra + rb – 0.09
(ΔEN) [r is in Ȧ]
2. Metallic radius: exist b/w metals bonded (equal to half of the value of
metallic bond length.)
3. Vander walls radius: exist b/w non bonded atoms (half of the shortest
inter nuclear distance b/w nuclei of non-bonded atom of noble gas or
non-bonded of adjacent molecule (non-metals) in solid state.
4. Order: VW>MR>CR

1. In case of multielectron atom, valence shell e are shielded by inner e from
nuclear charge of nucleus. This effect is called as S.E.
2. In this way valence e do not feel total charge of nucleus due to S.E. of
inner e. hence actual(net) charge felt by valence e is called as Zeff.
3. Zeff = Z - ∂ (∂ is screening constant of inner e)
4. Order of SE 4s> 4p > 4d> 4f (for same value of n)
5. Shielding of e of inner e also depends upon distance from nucleus. Ex 3d
subshell has more shielding than 4s.
6. Al and Ga metal have almost equal atomic size however they belong to
different period. (due to poor shielding of 3d e in Ga there is contraction
in size by more increase in nuclear charge)

1. Radius decreases due to increase in nuclear charge by one unit and
respective e enter to same shell.
2. In each period corresponding noble gas has its largest radius (because
of VW radius) and corresponding halogen has smallest radius
(covalent radius)
DOWN THE GROUP (DTG)
1. DTG atomic radius increases due to increase in number of shell
(principle QN) however there is increase in nuclear charge also by 2, 8,
8, 18, 18 and 32 units. Here increase in shell number dominates over
nuclear charge.
2. Size difference b/w two consecutive group elements decreases.

Atomic radius

ACROSS THE PERIOD (ATP)

Radius trend in d-Block elements:
1. In transition series atomic radius first decreases then becomes constant
and then increases.
2. Radius order in 3d, 4d and 5d series

3. On moving 3d to 4d radius increases. But from 4d to 5d radius is almost
constant due to lanthanide contraction (greater than expected decrease
in radius due to poor shielding of 4f e.)

SIZE OF ISOELECTRONIC SPECIES:
1. It depends on z/e ratio which is inversely proportional to the size.
2. Al3+ < Mg+2 < Na+ < F- < O-2 < N-3
3. Ar> P3- > S2- > Cl- > K+ > Ca+2
Exceptional size:
1. Li < Al
2. Li+ (73 pm) > Al+3 (53 pm) < Mg+2 (71pm)

Radius of ions:
1. Size of cation is always less than parent atom. Ex. Na> Na+
2. Size of anion is always greater than parent atom. Ex. I- > I > I+
3. Smallest monoatomic anion = FLargest anion = ISmallest cation = H+
Largest cation = Cs+
Smallest atom = H
Largest atom = Rn
Lightest element = H
Lightest metal = Li

Definition: amount of energy required to remove 1 e from isolated
gaseous neutral atom in its ground state. (Endothermic phenomena)

Factors influencing IE

M(g) + energy  M (g) + e , energy = IE1 of M = ΔH (+ve)
+

M+(g) + energy  M+2 (g) + e , energy = IE2 of M = IE1 of M+
M+2(g) + energy  M+3 (g) + e , energy = IE3 of M = IE1 of M2+ = IE2 of M+
IE1 < IE2 < IE3 for a particular element always. Successive IE of a same
element is always increasing irrespective of its stable configuration due to
continuous increase in nuclear charge.

1. In period IE increases due to increase in effective nuclear charge.
2. Highest IE in PT  He, lowest IE  Cs
3. Exception:

4. IE of 2nd group elements are higher than 3rd group because of stable
configuration of group 2nd wrt to group 3rd. [IE2 > IE13]
5. Same for group 15 and group 16. [IE15 > IE16]
6. In 3d, 4d and 5d series IE of Zn, Cd and Hg are abnormally high. [stable
configuration]

DOWN THE GROUP (DTG):
1. In DTG IE decreases due to increase in atomic size.
2. Exception:
In group 13th and 14th (p block elements) IE decreases upto 5th period
due to increase in atomic size. But 6th period elements has high IE than
respective 5th period elements due to poor shielding of 4f subshell.
Ex: IE, In < Tl And Sn < Pb
3. In 3d, 4d and 5d series: IE decreases from 3d to 4d but increases from
4d to 5d due to poor shielding of 4f e. [Lanthanide contraction] [ex.
Cu>Ag<Au]

IONISATION ENERGY

ACROSS THE PERIOD (ATP):

Example of IE1 and IE2 :
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.

IE1 (Mg) > IE1 (Na)
IE2 (Mg) < IE2 (Na)
IE1 (Cu) > IE1 (Zn)
IE2 (Cu) < IE2 (Zn)
IE1 (Mn) > IE1(Cr)
IE2 (Mn) > IE2(Cr)
IE2 (O) > IE2 (F) > IE2 (N)> IE2 (C)
IE3 (Mg) > IE3 (Al)
IE1 of F>P>S>B
IE Of Mg+2>Na+>F->O2->N3Cu>Ag<Au
Zn>Cd<Hg

Good Example
M(g)  M+ (g) + e– ; ΔH = 100 eV.
M(g)  M2+ (g) + 2e– ; ΔH = 250 eV.
Which is/are correct statement(s) ?
(A) IE1 of M(g) is 100 eV
(B) IE1 of M+ (g) is 150 eV.
(C) IE2of M(g) is 250 eV.
(D) IE2of M(g) is 150 eV.
Ans. (A,B,D)

Metallic property of s and p block elements

REDUCING PROPERTY OF METAL:

1. Decreases with increase in IE.
2. Hence left side elements are metal and right side are non-metal.

1. Reducing property depends upon the tendency of metal to
release valence e in aq solution easily.
2. Inversely proportional to IE.
3. Reducing order of ist group: Li<Na<K<Rb<Cr (expected)
Actual Na<K<Rb<Cs<Li
4. Why Li  due to very high H.E.

INERT PAIR EFFECT:
1. In general elements are stable in their higher oxidation state but in p block
elements from group 13 to 15th on moving DTG, lower OS of elements
gradually becomes stable and higher becomes unstable.
2. This is due to inertness of ns electron towards bonding.
3. This phenomena of inertness of ns e towards bonding due to poor shielding
effect of 4f subshell is called as inert pair effect.

ANS: BC, ABC
5. Pb+4 + 2e  Pb+2(more stable) [Strong Oxidizing Agent (SOA)]
Bi+5 + 2e  Bi+3(more stable) [SOA]
Tl+3 + 2e  Tl+ (more stable) [SOA]
Sn+2  Sn+4 (more stable) [SRA]

Application of IE

4. Correct stability order of cation is/ are
A. Pb+2 < Pb+4
A. Ga+<In+<Tl+
+2
+2
B. Sn <Pb
B. Ge+2<Sn+2<Pb+2
C. Sn+2<Sn+4
C. As+3<Sb+3<Bi+3
+4
+4
D. Sn <Pb
D. Ga+3 < In+3 < Tl+3

Relative stability of s and p block metal cations in different in OS.
1. Metal cation is stable in its higher OS when successive IE gap is
slightly higher than previous one.
2. Ex: Mg+2 is more stable than Mg+
3. Metal cation is stable in its lower OS when successive IE gap is
considerably higher than previous one.
4. Ex: Na+ is more stable than Na+2
5. Example: Successive energy of P block element X are given: IE1,
IE2, IE3 and IE4 are 5, 15, 30, 100 eV. Comment on position of X
in PT. [13th group]

Relative stability of d block metal cation in different OS.
1. d block metal shows variable OS due to less energy gap b/w ns
and (n-1)d e. relative stability is as follows:
IEn – IE(n-1) < 11 eV then higher OS is stable
11< IEn – IE(n-1) <20 eV Both OS are stable
IEn – IE(n-1) > 20 eV then lower OS is stable
2. Example: 1st and 2nd IE of 3 d block metals A, B and C are given.
Comment on their stability.

ACROSS THE PERIOD (ATP):

1. Amount of net energy released when 1 e is added to isolated
gaseous neutral atom in its ground state is called as electron
affinity of atom of element.
2. M(g) + e  M-(g) + energy , ΔH = -EA1
3. Electron affinity and electron gain enthalpy (ΔH) both are
equal but opposite in sign.
4. EA of M+ = IE of M
5. EA2 : when 1 more e is added to the valence shell of isolated
monovalent gaseous anion then more energy is required to
overcome repulsion b/w upcoming e and anion. In this case
net amount of energy is absorbed and is named as second
electron affinity.
6. First EGA is usually –ve but successive EGA is always +ve.
First EA is usually +ve but successive EA is always -ve.
7. M(g) + e  M- (g)

ΔH = -EA1 = -ve

M-(g) + e  M2- (g) , ΔH = -EA2 = +ve
8. Highest EGA in PT  Cl (-348 KJ/mole)
9. F + e  F- [ΔH = -328 KJ/mole]
10. O + e  O- [- EA1 = ΔH = - 141KJ/mole]
O- + e  O2- [- EA2 = ΔH = +780 KJ/mole]

ELECTRON AFFINITY
OR
ELECTRON GAIN ENTHALY

Definition:

1. EA increases as radius decreases and nuclear charge increases but with
some irregularities.
2. EA of 2nd and 15th group elements are exceptionally low due to their
full filled and half-filled stable configuration.
3. When e is added to these stable configuration then higher amount of
energy is required therefore such elements have almost zero or less
negative EA.
4. Be, N, Zn, Mn, Mg => almost zero EA

DOWN THE GROUP(DTG):
1. EA gradually decreases due to increase in atomic size but with some
irregularities.
2. EA of second period (non-metal) are exceptionally low than that of 3rd
period elements due to
1. Small size of 2nd period elements
2. High charge density of 2nd period elements.
Example C<Si>Ge>Sn>Pb
N<P>As>Sb>Bi
O<S>Se>Te>Po
F<Cl>Br>I

Definition:

Application of EN:

It is defined as tendency of atom of elements to attract shared paired
of e of single covalent bond towards itself. (not an energy, it’s a
tendency)

1. Nature of chemical bond:
a) No chemical bond is 100% ionic or covalent.
b) % I.C. = 16|XA – XB| + 3.5|XA-XB|2 [hennay and smith eqn]
If |XA – XB|= 2.1 [bond 50% ionic and 50% covalent]
|XA – XB|>2.1 predominate ionic bond
|XA – XB| < 2.1 predominate covalent bond
c) Acc to mulliken scale this difference is 1.7
2. Nature of compound in water having O-H linkage.
X-|- O—H  (basic in nature) [|Xx - Xo |> |Xo - XH| ]
X – O -|-H  (acidic in nature) [|Xo – XH |> |Xx – Xo| ]
EX: ClOH  acidic , TlOH  basic
3. Relative reduction of bond strength in water (polar solvent):
F = (1/4ᴫ ἕ) (Q1 Q2 / r2 )
F is inversely proportional to ἕ (permittivity constant of medium)
Air  1 , water  82, alcohol  32
Ex: HCl never act as a acid in air, it ionizes in aq solution to H+ and Cl, it act as stong acid in water.

ACROSS THE PERIOD(ATP):

DOWN THE GROUP(DTG):
1. It decreases as radius increases down the group.

Different scale to determine EN:
1. Mulliken’s scale:
EN(chi) = (EA + IE)/2
Above scale suffers because EA of every elements are not known.
2. Pauling’s scale:
a) It is widely accepted because it is based on B.E data of single
covalent bond. If BDE is taken in kcal/mol
b) |XA – XB|= 0.208√Δ
Where Δ = (A-B)BDE - √((A-A)BDE X (B-B)BDE)
c) He calculate EN value relative to H atom by taking H EN value
to be 1.
d) We found EN value of F = 4
e) F(4) > O(3.5) > N(3)=Cl(3) > Br(2.8) = S(2.8) > C(2.5) = I(2.5) >
H(2.1) < P(2.1)
f) ENPS = 2.8 ENMS

Electronegativity

1. It increases as nuclear charge increases and radius decreases.
2. Most EN atom of PT is fluorine.

Factors affecting EN
1. Charge on the atom of element:
As Positive charge increases Zeff increases hence EN increases.
Ex: Fe+2 < Fe+4 , Pb+2<Pb+4, Sn+2<Sn+4
2. Nature of surrounding atoms:
Ex: C in CH4 is less EN than C in CF4 .
3. Percentage S character of hybrid orbitals of same elements.
EN: SP> SP2>SP3


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