classification of elements and periodicity in …new.excellencia.co.in › college › web ›...

59 EXCELLENCIA JUNIOR COLLEGESS H A M I R P E T | M A D H A P U R | S U C H I T R A | E C I L

CLASSIFICACLASSIFICACLASSIFICACLASSIFICACLASSIFICATION OF ELEMENTS TION OF ELEMENTS TION OF ELEMENTS TION OF ELEMENTS TION OF ELEMENTS AND PERIODICITY IN PRAND PERIODICITY IN PRAND PERIODICITY IN PRAND PERIODICITY IN PRAND PERIODICITY IN PROPEROPEROPEROPEROPERTIESTIESTIESTIESTIES

CLASSIFICACLASSIFICACLASSIFICACLASSIFICACLASSIFICATION OF ELEMENTSTION OF ELEMENTSTION OF ELEMENTSTION OF ELEMENTSTION OF ELEMENTS

AND PERIODICITY IN PRAND PERIODICITY IN PRAND PERIODICITY IN PRAND PERIODICITY IN PRAND PERIODICITY IN PROPEROPEROPEROPEROPERTIESTIESTIESTIESTIESSYNOPSIS

Fundamentals

Ø At present around 114 elements are known.

Ø Out of these, recently discovered elements are not

natural but synthetic.,

Ø Elements coming after 92 atomic number are known

as “Trans Uranic Elements” or

“SyntheticElements” and they are

“Radioactive”.

Dobereiner Law of Triads

Ø Doberenier between 1815-1829 gave his law of

triads .

Ø A triad is a certain group of 3 elements with similar

properties.

Ø According to him in the triads the atomic weight

of the middle element was approximately the

arithmetic mean of the other two.

Also the properties of the middle element were in

between those of other two members.

eg-1:

Element Li Na K

Atomic wt. 7 23 39

Mean of atomic masses =7 39

232

+=

eg-2:

Element Cl Br I

Atomic wt. 35.5 80 127

Mean of atomic masses =

35.5 12781.25

2

+=

eg-3:

Element Ca Sr Ba

Atomic wt. 40 88 137

Mean of atomic masses =40 137

88.52

+=

Ø But in some triad all the three elements possessednearly equal atomic masses, hence the law wasrejected ,

eg: (Fe, Co, Ni) ; (Os, Ir, Pt) etc

According to him the properties of elements havesome relationship with their atomic masses.

De - Chancourtois Classification:

(Telluric Helix)

Ø In 1862 De-Chancourtois arranged the known

elements in order of increasing atomic weights and

made cylindrical table of elements to display the

periodic reoccurance of properties.

Newland Octaves:

Ø Newland in 1865 presented the law of Octaves “If

the known elements are arranged in the increasing

order of their atomic weights ,then the 8th element

had properties similar to those of first element” as

the eight note of octaves.

eg : Li Be B C N O F

Na Mg Al Si P S Cl

K Ca

Note: This law is true only for the elements up to

calcium.

Lother Meyer

Ø Lother Mayer (Germany) and Mendeleev

(Russia) quite independently evovled identically and

showed the connection between the periodicity of

properties and atomic masses of elements.

Ø Lother Meyer plotted the physical properties such

as atomic volume,melting point and boiling point

against atomic weight and obtained a periodically

repeated pattern.

Ø Lother Meyer calculated the atomic volumes of

known elements as the ratio of molecular weight

and density.

The findings of Lother Meyer curves are :

Ø Alkali metals having the largest atomic volumes

occupy the maxima of the curve.

Ø The alkaline earth metals (Mg ,Ca ,Sr,Ba) occupy

the mid point positions on the descending portions

of curve.

Ø Halogens occupy position on ascending portions

of the curve before inert gases.

Ø The transition elements occupy minima of the curve.

EXCELLENCIA JUNIOR COLLEGESS H A M I R P E T | M A D H A P U R | S U C H I T R A | E C I L

60


Mendeleev’s Classification of

Elements Periodic LawØ The physical and chemical properties of the

elements are periodic functions of their atomic

weights.

Ø Mendeleev’s periodic table is also known as short

form of periodic table.

Ø While arranging the elements in the periodic table,

he not only followed the increasing order of atomic

weights but also considered their properties.

Ø In original Mendeleev periodic table only 63

elements were known.

Ø The elements which are most widely distributed

in nature have small atomic weights and posses

sharply defined properties.

Ø Mendeleev observed that elements with similar

properties have

i) Almost same atomic weight.

eg : Fe(56), Co(59), Ni(59)

ii) Atomic weights increasing constantly

eg : K(39), Rb(85), Cs(133)

Ø Vertical columns are called groups and there are

nine groups ( 0 to 8th) and horizontal rows are called

periods and there are seven periods.

Ø The first three periods are short periods and

remaining are long periods. Each long period has 2

rows of elements or 2 series of elements

Ø Leaving 0 and VIII ,each group is subdivided into

subgroups known as A and B group.

Ø Group VIII of the Mendeleev’s table consists of

three triads known as transition triads and

they are

i) Iron, Cobalt and Nickel

ii) Ruthenium, Rhodium and Palladium

iii) Osmium, Iridium and Platinium

Ø Zero group elements were not known at the time

of Mendeleev and later introduced by R a m s a y

and Rayleigh.

Ø Mendeleev has a fore sight to leave some gaps in

the periodic table for 3 - elements and these

elements are discovered later and included in the

table. Those three elements are

1) Eka boron presently known as Scandium

2) Eka silicon presently known as Germanium

3) Eka aluminium presently known as Gallium

Ø Mendeleev corrected the atomic weights of

Beryllium, Indium and Osmium by using

corrected valency of elements

Atomic Wt. = Equivalent Wt. x valency. .

Merits:

Ø He gave an elaborate and

comprehensive system of classification ,based on

broad range of physical and chemical properties.

Ø He broadly left some gaps in discovered elements.It

led to discovery of some new elements

eg: Ge, Sc, Ga etc.

Demerits

Ø Some elements with higher atomic weight were

placed before lower atomic weight elements in order

to maintain similar chemical nature of elements and

are called inverted pairs or anamolous pairs.

Anamalous pairs of Mendeleev’s periodic table

are

a) Ar-K b) Co-Ni c) Te-I and d) Th - Pa

Ø Position of hydrogen was not made clear.

Ø Position of lanthanides are uncertain.

Ø No place for noble or inert gases .

Ø Absence of similarity in sub-groups

eg: alkali metals (IA) and coinage metals IB

(Cu,Ag,Au)

Ø Isotopes are not included

Ø Cause of periodicity is not known

Atomic NumberØ Moseley discovered the atomic numbers from

X-ray spectra of elements by bombarding the

elements with cathode rays and the elements emitted

respective X-rays of characteristic frequency.

Ø Atomic number ‘Z’ can be related to frequency of

the X-rays emitted by using ( )v a Z b= − where

a and b are constants for an element. As atomic

number increases the frequency of characteristics

X-rays increases.

Ø A plot of

v

against Z gives a straight line.

Ø Atomic number has provided a better basis for the

periodic arrangement of the elements.

Plot of υ and atomic number (z)



Modern periodic lawØ Physical and chemical properties of the elements

are periodic functions of their atomic numbers and

electronic configuration.

Long Form of Periodic TableØ Neil’s Bohr constructed the long form of periodic

table.

Ø Modern periodic table or the long form of periodic

table is based on the electronic configurations of

the elements.

Ø There are 18 groups and 7 periods in the periodic

table.

Periods (Horizontal Rows)Ø In periods, elements are arranged in the increasing

order of their atomic numbers.

Ø The electron by which an element differs from its

previous element is called “differentiating

electron”.

Ø In each period, the differentiating electron enters

into the “s” orbital in the first element and “p” orbital

in the last element.

Ø In periods, elements are arranged according to the

“(n+l)” values order (Aufbau-Rule).

Ø Long form of the periodic table is a Graphical

Representation of the Aufbau-Rule.

Ø Generally every period starts with an Alkali Metal

and ends with Noble gas.

Ø Period number corresponds to the highest principal

quantum number (n) of the elements in the period.

eg : First period contains 2 elements, the

subsequent periods consists of 8, 8, 18, 18 & 32

elements.

RememberØ In first period, 1s orbital is filled (shortest period)

Ø In second period, 2s, 2p orbital are filled (I short

period)

Ø In third period, 3s, 3p orbitals are filled (II short

period)

Ø In 4th period, 4s, 3d, 4p orbitals are filled (I long

period)

Ø Elements with 3d configuration from Sc to Zn

(Z= 21 to 30) are placed in 4th period. It is also

called first transistional series or 3d series.

Ø In 5th period, 5s, 4d, 5p orbitals are filled (II long

period)

Ø Elements with 4d configuration [from Y(Z=39) to

Cd(Z=48)] placed in 5th period (2nd Transition

series).

Ø In 6th period, 6s, 4f, 5d, 6p orbitals are filled

(longest period)

Ø Elements with 5d configuration from La(Z=57) and

Hf (Z=72) to Hg (Z=80) are placed in 6th period.

(3rd transition series).

Ø Fourteen 4f series elements (Lanthanoids) belongs

to 6th period and III B group. Ce (Z=58) to Lu

(Z=71).

Ø In 7th period, 7s, 5f, 6d, 7p orbitals are filled

(incomplete period)

Ø Fourteen 5f series elements (Actinoids) belongs to

7th period & III B group. Th (Z=90) to Lr (Z=103).

Ø 6d-series is incomplete series.

Ø If 7th period is also completed, then the final

element of this period would be with an atomic

number 118 (Uuo).

Groups (Vertical Columns)Ø Long form of the periodic table comprises of 18-

vertical columns which are divided into main groups

and subgroups as - IA to VIIA, O groups and IIIB,

IVB, VB, VIB, VIIB, VIIIB, IB and IIB groups.

Ø VIIIB groups includes three vertical columns of

Fe Co Ni

Ru Rh Pd

Os Ir Pt

Ø We adopt the 1-18 numbering scheme

recommended by IUPAC in 1988.

Ø Main group division is based on the number of

electrons present in outer most orbit like H, Li, Na,

K, Rb, Cs and Fr have 1 electron in their outer

most orbit, so they are placed in IA group. Be, Mg,

Ca, Sr, Ba and Ra have 2 electrons in their outer

most orbit, so they are placed in IIA group.

IUPAC Nomenclature for Elements with

Z>100Ø Nomenclature of elements CNIC (commission on

nomenclature of inorganic chemistry) appointed by

IUPAC in 1994, approved a nomenclature

scheme as well as also gave official names for

elements after Z > 100 (upto atomic number 104

to 109 discovered by that time).

Ø This nomenclature is to be followed for naming the

elements until their names are officially recognised.

Ø The names are derived by using roots for the three

digits in the atomic number of the element and

adding “ium” at the end. The roots for the numbers

are.


62


Nomenclature of Elements with Atomic

Number Above 100

Class ification of elements on the basis of

their Electronic configuration

Ø Elements are classified into four blocks basing on

the orbital into which the differentiating electron

enters.

a) s-Block Elements b) p-Block Elements

c) d-Block Elements d) f-Block Elements

s - Block Elements :Ø Differentiating electrons enter into s- orbital of

valency shell.

Ø s-sublevel can accomadate 2-electrons, hence s-

block elements are arranged in two groups, IA, IIA

(or) 1, 2 groups

Ø General electronic configuration is 1 2ns

− .

Ø H, Li, Na, K, Rb, Cs, Fr elements (alkali metals)

have 1 electron in their outer shell with “

1ns

”

general outer shell configuration, they belongs to

IA.

Ø Be, Mg, Ca, Sr, Ba and Ra (Alkaline Earth

elements) have 2-electrons in their outer shell, with

“

2ns

” general outer shell configuration, they

belongs to IIA.

Ø Most of these are active metals and form ionic

substances, except lithium and beryllium.

Ø These are powerful reducing agents.

Ø They have low M.P’s and B.P’s.

Ø They impart characteristic colours in the flame

p - Block ElementsØ Differentiating electron enters into p- orbital of

valency shell

Ø The general outer shell configuration of p-block

elements.

2 1 6ns np

Ø p-block elements are arranged in 6-groups they are

from IIIA to VII A and O-group (or) 13 to 18

groups

Ø B,Al,Ga,In and Tl are called IIIA group (boron

family) these elements have 3-electrons in outershell

with “

2 1ns np

” general outer shell configuration.

Ø C,Si,Ge,Sn and Pb are called IVA group

(Carbon Family) these elements have 4-electrons

in outer shell,with “

2 2ns np

” as general outer shell

configuration.

Ø N,P,As,Sb and Bi are called VA group(Nitrogen

Family - pnicogens). These elements have 5-

electrons in outer shell, with “

2 3ns np

” as general

outer shell configuration.

Ø O, S, Se, Te, and Po are called VIA group

(Chalcogens) these elements have 6-electrons in

outer shell, with “

2 4ns np

” as general outer shell

configuration.



Ø F, Cl, Br, I and At are called VIIA group

(Halogens) these elements have 7-electrons in outer

shell, configuration with “ns2np5” as general outer

shell.

Ø He, Ne, Ar, Kr, Xe and Rn - Inert gases (O–group),

Except He ( 21S ), remaining inert gases have 8-

electrons in outer shell with “

2 6ns np

” as general

outer shell configuration.

Ø p-block contains all non-metals and metalloids

and some metals.

Ø Most of the compounds of p-block elements are

covalent.

i) Most of these are oxidising agents

j) All gaseous elements except H and He are p-

block elements.

Remember

Ø Keeping its chemical inertness, Helium is placed

along with other inert gases in 0 - group.

Ø Hence He is a p-block element with out

p-electrons.

Ø The first p-block element is Boron [(He) 2s2 2p1]

Ø The only group with all gaseous elements is “0-

group”.

d-Block ElementsØ If the differentiating electron enters into the d-orbital

of penultimate shell, the elements are called “d-

block elements”.

Ø The general electronic configuration of d-block

elements is

( ) 1 10 1 21n d ns− −

−

(n = outer shell).

Ø d-Block elements are placed between s-block and

p-block and they are also called transition elements.

Ø d-Block elements are further classified into following

transition series on the basis of which (n-1)d subshell

is being filled.

1) for 1st Transition series( 3d series) electronic

configuration is

3d1-10 4s1-2 [Sc (Z=21) to Zn (Z=30)]

2) for 2nd Transition series ( 4d - series) electronic

configuration is

4d1-10 5s1-2 [Y(Z=39) to Cd (Z=48)].

3) for 3rd Transition series (5d - series) electronic

configuration is

5d1-10 6s1-2. [La (Z=57), Hf (Z=72) to Hg (Z=80)]

4) 4th Transition series( 6d - series) is an incomplete

series.

5) Most of these are less active metals.

6) These elements form ionic and co-ordinate

covalent compounds.

7) They are all solids, except Hg which is a liquid at

room temperature.

8) They form cations with high charge.

9) They form alloys and interstitial compounds.

10) They mostly form coloured ions and also show

paramagnetism.

Remember

Ø After completion of 6s, the differentiating electron

suppose to enter into 4f, but in the case of

Lanthanum the differentiating electron is entering into

5d, instead of 4f ( 2 0 16 4 5La s f d− ). Therefore

“La” belongs to d-block (IIIB, VI period).

Ø Similarly in case of Actinium, the differentiating

electron is entering into 6d, instead of 5f

( 2 17 5 6oAc s f d− ). Therefore Ac also belongs to

d-block (IIIB, VII period).

f-block Elements:Ø If differentiating electrons enter into f-subshell of

anti penultimate i.e., (n-2) shell, the elements of this

class are called f-block elements.

Ø The general electronic configuration

( ) ( ) ( )0 11 14 22 1

orn f n d ns

−− −

(n = outer shell).

Ø These f-block elements are placed at the bottom

of the periodic table in two rows, they are 4f series

and 5f series. The properties of 4f-series elements

are similar to Lanthanum they are known as

Lanthanides (or) Lanthanons or rare earths.

Ø 4f-series - Lanthanide series - configuration

1 14 0 1 24 5 6f d s− − from Ce(58) to Lu (71) (first

inner transitional series)

Ø 4f- series elements belongs to 6th period and IIIB

Group.

Ø 5f - series elements - Actinide series - configuration

1 14 0 1 25 6 7f d s− − from Th (90) to Lr (103)(second

inner transitional series).

Ø 5f - series elements belongs to 7th period and III B

group.

Ø Most of these elements are radioactive.

Ø They have properties similar to d-block elements

Classification based on chemical

properties.Ø All the elements are divided into four types on the

basis of their chemical properties and electronic

configuration.


64


Type-1 Inert gasesØ He, Ne, Ar, Kr, Xe and Rn belongs to “0” group in

the periodic table are called Inert Gas Elements

Ø Except He (1s2), all the other elements have ns2np6

outer electronic configuration.

Ø All are chemically inert due to the presence of stable

ns2np6 (octet) configuration in their outer most shell.

He is inactive due to its completely filled ‘K’ shell.

(1s2)

Ø It is known that heavier elements (Kr, Xe) form

compounds under special controlled conditions with

Oxygen and Fluorine, So they are now called

Noble gases.

Ø All are monoatomic gases.

Ø They are also known as Rare gases (or) Aerogens.

As they are present in 1% by volume in atmosphere.

Type-II Representative elements or

normal elementsØ In these elements, the ultimate shell is incompletely

filled.

Ø Excluding “0” group, remaining s and p block

elements (IA, IIA, IIIA, IVA, VA, VIA, VIIA) are

called representative elements.

Ø Most of these elements are abundant and active.

Ø Their general outer electronic configurations

ns1-2 np1-5.

Ø Metals, non-metals and metalloids are present in

representative elements.

Ø Atoms of these elements enter in chemical

combination by losing, gaining or sharing of electrons

to attain stable nearest inert gas configuration.

Ø In case of representative elements electrons of outer

ns and np will take part in bonding.

Type - III Transition elementsØ In these elements , the ultimate shell and penultimate

shells are incompletely filled.

Ø Elements which have incompletely filled or partly

filled d-orbitals either in elementary state or in any

possible oxidation state are called transition

elements.

Ø Their properties are intermediate between s - and

p - block elements.

Ø The general electronic configuration is

( ) 1 10 0 21n d ns

− −− .

Ø II B group elements Zn (3d10 4s2), Cd (4d10 5s2)

Hg (5d10 6s2) are not transition elements (due to

the absence of partly filled d-orbitals both in atomic

and in ionic states) (Zn, Cd, Hg - are referred as

Non-typical Transition Elements) or volatile metals.

Ø In the case of Transition elements both

(n-1)d and ns electrons participate in bonding.

Ø The characteristic properties of transition elements

are

1. They are hard and heavy metals

2. Variable Oxidation states

3. Formation of coloured ions in solution due to

d-d- transition

4. Formation of metal complexes

5. Paramagnetic

6. Catalytic activity.

7. High M.P., B.P and densities.

8. Good conductors of heat and electricity

9. Alloy formation.

Ø These characteristic properties are due to

a. Small size

b. High nuclear charge

c. Unpaired electrons in d-orbitals.

Note:

1. Ni is used as a catalyst in Hydrogenation of oils.

2. Fe used as a catalyst in Haber’s process

3. Mo used as a promoter in Haber’s process.

Type-IV Inner Transition elementsØ These elements have three outermost shells

incomplete i.e., n, (n-1) and (n-2) (ultimate,

penultimate and antipenultimate shells).

Ø The f-block elements are called inner transition

elements.

Ø General configuration

( ) ( ) ( )11 14 22 1

o orn f n d ns

−− −

Ø Since the last two shells have similar configuration

these elements have similar physical and chemical

properties (eg - these elements shows common

oxidation state of +3).

Ø There are two series of inner transition elements.

4f- (Lanthanide) series - ( )11 14 24 5 6o or

f d s−

5f - (Actinide) series - ( )0 11 14 25 6 7

orf d s

−

In periodic table, lanthanides are present between

57La &

72Hf and

Actinides are present between 89

Ac & 104

Rf.

Ø Lanthanides are rare earths and actinides are mostly

synthetic.

Ø The elements from Z = 93 onwards are called

transuranic elements.



Pseudo Inert Electronic Configuration

Presence of 18 electrons in the outer most shell is

called pseudo-octet or pseudo-inert configuration.

Palladium ([ ] 0 10

36 5 4Kr s d ), a member of group 10

has such configuration.

Ø ‘La’ belongs to d-block but lanthanides are f-block.

similarly ‘Ac’ belong to d-block but actinides are

f-block

Periodicity - Periodic PropertiesØ When elements are arranged in increasing order of

atomic number, elements with similar properties

reoccur (due to similar outer electronic

configuration) at regular intervals of atomic numbers

in the periodic table. This repetition of properties is

called periodicity and such properties are called

periodic properties.

Ø Some of the properties which mainly depend on

the electronic configuration of elements such as

i) Valency ii) Effective nuclear charge

iii) Screening effect iv) Atomic radius

v) Ionic radius vi) Ionisation potential

vii) Electron affinity viii) Electronegativity

ix) Metallic nature

x)oxidation and reduction ability

xi) acidic or basic nature of the oxides, etc....

follow the general trend of periodicity. They are

called periodic properties. These properties are

especially important in s- and p-block elements.

Ø Properties like specific heat, refractive index, colour

etc., are not called periodic properties. These

properties are not related to the electronic

configuration of elements.

Ø Elements coming at intervals of 2, 8, 8, 18, 18, 32

will have similar properties and thus grouped in one

particular group.

eg-1 : Elements with atomic number 1, 3, 11, 19,

37, 55 & 87 will have similar properties.

eg-2 : Elements with atomic number 4, 12, 20, 38,

56 & 88 will have similar properties.

Note : Two successive elements in a group generally

differ by atomic number 2, 8, 8, 18, 18, 32.

Atomic RadiusØ In atoms, the electron cloud around the nucleus

extends to infinity.

Ø The distance between the centre of the nucleus and

the electron cloud of outer most energy level is called

atomic radius.

Ø Atomic radius cannot be determined directly, but

measured from the inter nuclear distance of

combined atoms, using X-ray diffraction and other

spectroscopic methods

Ø Atomic radius depends on

a) Nature of bonding

b) Number of bonds (multiplicity of bonding)

c) Oxidation state(s)

d) Co-ordination number of atom

e) bond character etc.

Ø Three types of atomic radii are considered based

on the nature of bonding they are

a) Crystal radius

b) Vander waals radius

c) Covalent radius

Ø Atomic radii expressed in angstrom, nanometers,

picometer units.

0 1 0 21A 10 nm;1A 10−= =

pm

Ø Crystal Radius (Atomic Radius) - Half of the

internuclear distance between the adjacent atoms

of a solid metallic crystal is called crystal radius

or metallic radius.

eg: The distance between two adjacent copper

atoms in solid copper is 256 pm; so metallic radius

of copper is assigned as value of 128 pm.

Ø Van der waals radius - Half of the internuclear

distance between two atoms of different molecules

which are very close to each other in solid state

due to vander waals forces is called Van der waals

radius.

Ø The distance between two adjacent chlorine atoms

of different Cl2 molecules is 360 pm, Vander waals

radius of Cl is 180 pm.

Ø Vander waals radius is 40% greater than covalent

radius.

Ø It is used for molecular substances and inert gases

in the solid state only.

Ø Covalent Radius: This term is generally used in

reference to non-metals.

Ø Covalent radius - Half of the inter nuclear distance

of the two atoms held together by a covalent bond

is called covalent radius.

Ø Note : Single bond covalent radii are additive in

nature.

eg : a) In Cl2 molecule Cl - Cl bond distance

(Internuclear distance) is 198 pm.

Covalent Radius of Cl = 99 pm.

b) In diamond C-C bond distance is 154pm.

Covalent radius of C = 77pm.


66


Ø In metals, the crystal radius (atomic radius) is slightly

more than the covalent radius.

Ø As the number of covalent bonds between two

atoms increases, the inter atomic distance between

the carbon atoms decreases

C – C > C = C > C ≡ C

(1.54A0) (1.34A0) (1.20A0)

Order of radii :Van der waal radius > crystal radius > covalent

radius.

Ø Compared to theoritical atomic radius, covalent

radius of an atom is about 20% shorter due to

overlapping of atomic orbitals.

Variation of Atomic Radius in Groups

and PeriodsØ In a period from left to right, atomic radius

decreases as the effective nuclear charge increases.

Ø Variation of atomic radius

eg-1 : In second period

Li Be B C N O F> > > > > >

eg-2 : In third period

Na > Mg > Al > Si > P > S > Cl

Ø On moving from left to right across a particular

period, the atomic radius decreases upto Halogens

and increases to Inert gases.

Ø In a given period, alkali metal is the largest and

halogen is the smallest in size.

Ø However, the radius of an inert gas is larger than

the halogen of the same period.

Ø Note : For atoms of Inert gases, only vanderwaal

radius is applicable because these are mono atomic

gases.

Ø In groups from top to bottom, the atomic radius

increases gradually due to the increase in the number

of orbits and it over weighs the effect of increased

nuclear charge.

Ø Atomic radius is least for hydrogen and is highest

for Caesium among the available elements.

Ø Variation of atomic radius

In IA group is Li < Na < K < Rb < Cs

In halogens is F < Cl < Br < I < At

Variation of Atomic Radius in

Transition Elements:Ø In case of transition elements, the decrease in size

in a period across a particular transition series is

less than in case of representative elements, this is

due to less screening effect of (n-1)d-electrons.

Ø Hence, the atomic radius decreases slightly as we

move from left to right in a transition series.

Ø From Cr to Cu the covalent radii is almost same

due to

1) Shielding effect of core electrons

2) Additional shielding effect of 3d electrons.

Ø Covalent radii of Zn is more than Cu due to

repulsions among 3d electrons.

Variation of Atomic Radius and Ionic

Radius in Lanthanides:Ø The elements in Lanthanide series are La, Ce, Pr, Nd,

Pm, Sm, Eu, Gd, Tb, Dy, Ho, Er, Tm, Yb & Lu

Ø In Lanthanides (Ce-Lu) the atomic and ionic radii

decreases steadily. This steady decrease in atomic

and ionic radii is known as “Lanthanide Contraction”

Ø The contraction is due to the fact that f-orbitals are

not capable of providing effective shielding for the

valence electrons from nuclear attraction due to

diffused shape.

Consequences of Lanthanide

ContractionØ Atomic sizes of 4d and 5d series of transition

elements become almost equal, due to which their

properties are very close.

Ø Zr and Hf, Nb and Ta, Mo and W resemble very

closely.

Ø The crystal structure and other properties of

lanthanides are very similar.

Ø Separation of lanthanides is not easy from their

mixture. Chromatographic techniques can not

separate lanthanides from their mixture.

Ø Super heavy metals of p- block exhibit inert pair

effect.

eg : Tl(III A group), Pb (IV A group),

Bi (V A group)

Variation of Atomic Radius and Ionic

Radius in Actinides :Ø The elements in actinide series are Ac, Th, Pa, U,

Np, Pu, Am, Cm, Bk, Cf, Es, Fm, Md, No and Lr

Ø The size of the trivalent ions of these elements

decreases regularly as we move from left to right.

This is because of poor shielding effect of f-electrons,

more nuclear charge and diffused shape of f-orbitals.

This is called Actinide contraction.

Ionic RadiusDefinition : It is defined as the distance between the

nucleus and the electron in the outer most shell of

an ion.



Ø When a neutral atom loses one (or) more electrons

a positive ion called cation is formed.−+

+→ eNaNa

Ø The ionic radius of cation is less than that of neutral

atom. It is because the cation has higher effective

nuclear charge.

eg: Na Na+

>

Ø Among the cations as the positive charge increases,

the ionic radius decreases.

eg: ++>

32FeFe ,

2 4 2 4,Sn Sn Pb Pb+ + + +> >

Ø When a neutral atom gains one (or) more electrons

a negative ion called anion is formed.

eg: −−→+ CleCl

Ø The radius of anion is more than that of its atom,

due to decrease in effective nuclear charge.

eg: ClCl >−

Ø Among the anions as the negative charge increases

the ionic radius increases.

eg: 2O O

− −>

Ø The decreasing order of the radii is

Anion > Atom > Cation

eg: +−>> III ; +−

>> HHH

Ø In a particular group, the ions (cations or anions)

increase in size on moving from top to bottom due

to increase in number of shells.

eg: Li Na K Rb Cs+ + + + +

< < < <

F Cl Br I− − − −

< < <

Ø

&H Cs+ +

are the smallest and largest cations

respectively.

Ø

&H I− −

are the smallest and largest anions

respectively.

Ø Smallest atom is He & largest atom is Cs.

Iso Electronic SpeciesØ The species (atoms or ions) having the same number

of electrons are known as iso - electronic species.

In iso electronic species, the size increases

with increase of negative charge and decreases

with increase of positive charge.

Decreasing order of size. C4- > N3- > O2- >

F- > Ne > Na+ > Mg2+ > Al3+ > Si4+

Ionization Energy (Ionization

Potential)Ø Ionization potential: The minimum amount of

energy required to remove the most loosely bound

electron (i.e, outer - most shell electron) from an

isolated neutral gaseous atom is called ionization

potential.

( ) ( )

1

1g gM IE M

++ →

+ e-

Ø It is an endothermic process

Ø IE is measured in eV/atom or kJ/mole or K.cal/

mole.

1 eV / atom = 23.06 K.Cal/mole = 96.45 KJ/mole

= 191.602 10 J / atom−×

Ø Energy required to remove an electron from

unipositive gaseous ion to convert it into dipositive

ion is IE2. ( ) ( )

++→+

2

2 gg MIEM + e-

Ø Energy required to remove an electron from

dipositive ion to convert it into tripositive ion is IE3.

( ) ( )++

→+3

3

2

gg MIEM + e-

Ø Ionization energy is determined by spectral studies

or discharge tube experiments.

Ø Ionization potential depends on :

1) Atomic size 2) Nuclear charge

3) Screening/shielding effect

4) Penetrating nature of orbitals

5) Electronic configuration

Ø With increase in the atomic size “IP” decreases due

to decrease in attractive force of nucleus on outer

most orbit electrons.

Ø With increase in the effective nuclear charge IP

increases.

Ø If the number of electrons in the inner shells are

more, shielding capacity of the inner electrons on

the nuclear charge will be more. Hence IP decreases.

Ø Order of screening power of orbitals

s > p > d > f

Ø As the positive charge on cation increases, IP

increases.

Ø As the -ve charge on anion increases, IP decreases.

Ø If the valency electrons are more penetrated into

inner shells, IP increases.

Ø Penetration power of different orbitals is in the order

of s > p > d > f

Ø IP of s-electrons > IP of p-electrons > IP of d-

electrons > IP of f-electrons.

Ø IP is more for atoms with exactly half filled and

completely filled orbitals.

eg: 1) IE1 of N > IE

1 of O 2) IE

1 of Be >IE

1 of B

3) IE1 of P > IE

1 of S

4) IE1 of Mg >IE

1 of Al


68


Ionisation potential curve upto element

sodium

Ø Atoms of inert gases have highest IP values due to

the presence of completely filled orbitals.

Ø In the graph showing relation between IP and

atomic number, the inert gases appear at the maxima

and alkali metals appear at the minima positions

Ionization Energy Curve

Ø The ionisation potential is the highest for helium

among all elements. The value is the least for

caesium among the available elements. When

ionisation potential values are plotted against atomic

numbers, the ionisation potential curve is obtained

as shown above.

Variation of IP in Groups & PeriodsØ In periods from left to right side IP increases, due

to decrease in atomic size and increase in effective

nuclear charge.

Ø In any period an Alkali metal atom has lowest IP

and Inert gas element has highest IP.

Ø In groups from top to bottom, IP decreases due

to the increase in the atomic size and increase in the

screening effect of inner electrons.

Ø IE order among 2nd period elements.

IE1 Li < Be > B < C < N > O < F < Ne

IE2 Li > Be < B > C < N < O > F < Ne

Ø IE order among 3rd period elements

IE1 - Na < Mg > Al < Si < P > S < Cl < Ar

IE2 - Na > Mg < Al > Si < P < S > Cl < Ar

Ø Element with Lowest IP - Cs and element with

highest IP is He.

Ø IE1 of Be greater than B due to

a) Completely filled s -orbital in Be

b) More Penetration of s-orbitals.

Ø Similarly IE of Mg is greater than Al

Ø Variation of First I.P in I A group elements

Li > Na > K > Rb > Cs

Ø I.E of coinage metals is Cu > Ag < Au.

WE1.The first ionization enthalpy ( )tH∆ values of

the third period elements, Na, Mg and Si are

respectively 496, 737 and 786 kJ mol-1. Predict

whether the first tH∆ value for Al will be more

close to 575 or 760 kJ mol-1? Justify your

answer.

Sol: It will be more close to 575 kJ mol-1. The value for

Al should be lower than that of Mg because of

effective shielding of 3p electrons from the nucleus

by 3s-electrons.

Knowledge of Successive IE

Ø Knowledge of successive IE can be used to find

the number of valence electrons

Ø For alkali metals the IE2 shows sudden jump.

Ø For alkaline earth metals, the IE3 shows sudden

jump.

Ø Theoretically, the number of IE possible for an atom

of an element is equal to its atomic number.

WE2.The successive ionization enthalpies of an

element M are 5.98, 18.82, 28.44, 119.96,

153.77, ….. eV/atom. What is the formula of

chloride of M?

Sol: Observing the I1, I

2, I

3, I

4, I

5, …. it is noticed that

there is a sudden jump form I3 and I

4.

This observation gives the idea that the element has

3 electrons in the outer most shell.

M3+ state is stable and valency is 3.

Formula of chloride of M is MCl3



WE3.The ionization enthalpy of sodium is 5.14 eV.

How many k cal of energy is required to ionize

all atoms present in one gram of gaseous Na

atoms?

Sol: 1eV atom-1 = 23 k cal mol-1

Energy required to ionize all atoms of 23 grams

(one mole) of gaseous Na atoms = 23 x 5.14 k cal

Energy required for ionization of all atoms present

in one gram of gaseous Na atoms 23 5.14

23

×=

= 5.14 k cal

Electron Affinity (EA) (or) Electrongain Enthalpy

Ø It is an atomic property which gives us an idea ofthe tendency of the element to accept the electronto form an anion.

Ø The amount of energy released when an electron isadded to a neutral isolated gaseous atom of anelement is called EA.

( ) ( ) 1EAXeX gg +→+−− (or)

( ) ( )−−

→+ gg XeX 1EAH −=∆ (Exothermic

process)Ø When an electron is added to uni-negative ion,

energy is absorbed to overcome the repulsiveforces. This energy is called second electron affinity.

EA2 has positive sign. ( ) ( )

2

g gX e X

− − −+ →

2EAH +=∆ (Endothermic process).

eg: EA1

of oxygen ( )( ) ( )1

g gO e O− −

+ → is

exothermic

But EA2 of oxygen ( )2

( ) ( )1

g gO e O

− − −+ → is

EndothermicØ EA is measured in eV/atom, Kcal/mole, KJ/moleØ EA can be calculated indirectly from Born - Haber

Cycle.Ø EA depends on size, effective nuclear charge,

shielding effect and electronic configuration of anelement.

Ø Noble gases have most stable ns2np6 configuration.Hence their EA values are positive values.

Ø For N, P - due to half filled orbitals, they have extrastability hence their EA values are close to zero (verysmall values).

Ø First electron affinity (E1) is negative for all elements

except for Be, Mg, N atoms and zero-groupelements.

( )1Be 66 KJ mol−+

( )1Mg 67 KJ mol−+

( )1N 31KJ mol−+

Variation of EA in Groups & PeriodsØ In groups, EA decreases from top to bottom as

the atomic size increases.Ø In a period from left to right side EA increases

due to decrease in size of atoms and increase in thenuclear charge.

Ø EA1 of third period element is greater than

corresponding second period element of each group(or) second element have high EA than first elementin a groupeg: 1) In VII A group EA of Cl > EA of F 2) VIA group EA of S > EA of O 3) VA group EA of P > EA of N 4) IV A group EA of Si > EA of C

Ø EA of F (-333 K.J mole-1) < EA of Cl (-348K.Jmole-1). This is due toa) Smaller size of F-atomb) Strong inter electronic repulsions

Ø Note : EA of a neutral atom = IE of its uninegativeion.EA of X = IE of X-

Ø Note : IE of a neutral atom = EA of its unipositiveion.IE of X = EA of X+

Ø Among all the elements chlorine has the maximumEA.

Ø The metal which has higher EA is Gold.

Electron Gain Enthalpies (kJ/mol) ofSome main Group elements


70


Ø Among halogens the order of EA values isCl > F > Br > I > At

Ø Decreasing order of E.A of various chalcogens is S

> Se > Te > O

WE4.Process (A): 2( ) ( )2 2g g

F e F− −

+ →

Process (B): 2( ) ( )2 2g g

Cl e Cl− −

+ →

Which of these processes is easy? Why?

Sol: 2( ) ( )2 2g g

F e F− −

+ → is easy.

Though electron gain enthalpy of Cl(g)

to give Cl-(g)

is more than that of F(g)

to give F-(g)

, the bond

dissociation of F2(g)

is very less than that of Cl2(g)

Electro Negativity (EN)Ø It is property of an atom in a molecule.

Ø The tendency of an atom to attract the shared

electron pair towards itself in a molecule is called

EN.

Ø E.N. is a relative property and has no units.

Ø Pauling Scale : EN of elements are calculated

from the values of bond energies.

Ø Pauling calculated the EN of other elements by using

the formula

∆=− 208.0BA XX .

[ ∆ is in K.Cal./mole.]

In SI units,

−AX

,

[ ∆ is in KJ/mole.]

where XA and X

B are the EN’s of A & B.

∆

is a measure of the polarity of A-B bond.

∆

= Experimental BE - Theoritical BE

Ø

∆

= Actual BE -1/2 [EA-A

+ EB-B

] , BE = Bond

Energy

Ø Hydrogen (whose EN is 2.1) is used to calculate

EN of other elements.

Ø The reference element taken by pauling for the

determination of E.N. Values of other elements is

hydrogen.

Ø Highest E.N. value is for fluorine (4.0).

Ø As the oxidation number of an atom increases, the

attraction for the electrons increases and E.N also

increases

Ø EN concept is not applicable for Inert gas elements.

WE5.Bond energies of H2, Cl

2 and HCl are

respectively 104, 58 and 100 kcal mol–1.

Calculate Pauling’s electronegativivy of

chlorine.

Sol: Average of bond energies of H2 and Cl

2 is the

calculated bond energy of

HCl

= 81 k cal mol–1

Experimental bond energy of HCl =100 k cal mol–1

∆ = Bond (resonance) stabilization energy

= 100 – 81 = 19 k cal mol–1

1 2X X− = ∆

0.208 19=

= 0.208 × 4.358 = 0.90

Since Pauling’s electro negativity of hydrogen is 2.1,

that of chlorine = 2.1 + 0.9 = 3.0

Mulliken ScaleØ According to Mulliken scale,

EN is the average of IE and EA. 2

EAIEEN

+=

Ø6.5

eVinEAeVinIEEN

+=

Ø( ) ( )/ /

540

IE in kj mole EA in kj moleEN

+=

Ø( ) ( )/ /

129

IE in kcals mole EA in kcals moleEN

+=

Ø Mulliken EN values are approximately 2.8 times

greater than Pauling EN values.

Ø Mulliken scale is applicable only to univalent

elements.

Ø Elements with same EN in pauling’s scale are

N = Cl = 3.0 C = S = I = 2.5

H = P = 2.1 Cs = Fr = 0.7

WE6.The ionization enthalpy of sulphur is 1014

kJ mol-1. If its electronegativity is 2.4, what is

its electron gain enthalpy?

Sol: In the common scale, electronegativity (E.N.) is

given in terms of ionization enthalpy (I1) and electron

gain enthalpy (E1) as

1 1.540

I EE N

+=

Substituting the values,

11014

2.4540

E+=

Electron affinity = E1 = (540 × 2.4) – 1014 = 282

Electron gain enthalpy of sulphur= -282kJmol–1



WE7.If the electronegativity value of fluorine in

Pauling scale is 4.0, then value in Mulliken

scale will be?

Sol: Electronegativity in Mulliken scale is 2.8 times

greater than Pauling scale values.

So the value of Electronegativity = 2.8 × 4 = 11.2

Variation of EN in Groups & PeriodsØ In groups from top to bottom EN decreases.

Ø eg: In I A group Li > Na > K > Rb > Cs

In halogens F > Cl > Br > I > At

Ø In periods from left to right EN increases.

Ø eg: In II period

Li < Be < B < C < N < O < F

Ø In a period, Halogen has high EN value.

Alkali metal has low EN value.

Ø Highest EN element is F(4.0)

Ø Next to F, oxygen has high EN (3.5)

Ø Least EN element is Cesium (0.7)

Ø Noble gas elements have zero EN due to octet

configuration.

Ø EN values are used to know the nature of

chemical bond.i) If EN difference is less than 1.7, the bond is

covalent in nature.

ii) equals to 1.7, the bond is 50% ionic in nature.

iii) more than 1.7, the bond is ionic in nature.

Ø E.N. values are useful in writing the formula of a

compound.

Ø E.N. values are useful in predicting the nature of

the element (metal / non-metal).

ValencyØ Valency of an element is the number of H-atoms

(or) double the number of oxygen atoms that can

combine with one atom of that element.

Ø The valency of an element is not always constant.

Ø Exhibition of more than one valency by one element

is known as variable valency.

Ø The maximum valency of a representative element

is equal to the number of electrons present in the

outermost orbit of an atom.

Ø Highest valency ever known is 8 shown by Os, Ru

and Xe

In 4OsO , the valency of Os is 8

In RuO4 the valency of Ru is 8

In

4XeO

, the valency of Xe is 8

Covalency: The number of covalent bonds formed

by an element

eg: 1) In NH3, covalency of ‘N’ is 3

2) In N2O

5, covalency of ‘N’ is 4

Ionic covalency (or) Electro valency: No.of

electrons transferred (either gain of electrons (or)

loss of electrons).

eg: 1) In NaF (Na+F-), Ionic valency of ‘Na’ is

1 and that of ‘F’ is 1.

2) In AlF3 (Al+3, F-); Ionic valency of ‘Al’ is

3 and that of ‘F’ is 1.

Ø The minimum valency exhibited by an element is

zero.

Ø Periodic trends in valence of Elements as

shown by the Formulas of their compound

WE8.Using the periodic table, predict the formula

of compound formed between an element X of

group 13 and another element Y of group 16.

Sol: The Valency of X (group 13) = 3

The valency of Y (group 16) = 2

The compound has 2 atoms of X and 3 of Y.

Hence, the formula = X2Y

3.

Oxidation StateØ The possible charge with which an atom appears in

a compound is called its oxidation state.

Ø s-block elements, oxidation state is equal to group

number. For alkali metals “ +1 ”.

For alkaline earth metals “ +2 ”

Ø Oxidation state may be positive or negative or zero

or fraction.

Ø p-block elements show multi valency, their

oxidation state change by two numbers.

Ø In III A groupthe stable oxidation state of Thallium

is +1. It is due to inert pair effect.


72


Ø In IVA group +2 is more stable than +4 for Lead

due to inert pair effect.

Ø In VA group, +3 is more stable than +5 for Bismuth

due to inert pair effect.

Ø Group IV elements show +4 and +2 oxidation

states.

Ø Group V elements show +5 and +3 oxidation states.

Ø The general oxidation state of group VI is -2.

Ø Generally oxygen shows -2 oxidation state in its

compounds but when it combines with fluorine it

shows +2 (in2OF ) and +1

(inO F

.

Ø The most electronegative element. Fluorine shows

-1 oxidation state only (in its compounds)

Ø The common oxidation state of d-block elements

is +2. All transition elements show variable

valencies.

Ø Ruthenium, Osmium and Xenon exhibit maximum

oxidation state +8.

Ø In d-block elements , +1 oxidation state is shown

by Cr, Cu, Ag, Au, Hg.

Ø The common oxidation state of f-block elements is

+3 due to their outer electron configuration

ns2(n - 1)d1.

Ø Maximum oxidation state of an element never

exceeds its group number.

WE9.What is the valency and oxidation number

of nitrogen in nitrogen pentoxide?

Sol: Based on the oxide theory, valency of N in N2O

5 is

5 (But the actual valency of N in N2O

5 is the number

of bonds formed by N = 4).

Oxidation number of N in N2O

5 = +5

WE10.Are the oxidation state and covalency of Al

in [AlCl(H2O)

5]2+ same?

Sol: No. The oxidation state of Al is +3 and the

covalency is 6.

Electro Positive Nature (EP)Ø The tendency of an element to lose an electron is

called electro positivity.

Ø It is the converse of electro negativity.

Ø As electropositivity increases, metallic character

increases.

Ø The smaller the ionisation energy or ionisation

potential the greater is the electro positivity.

Ø As electropositive nature increases, capacity to form

ionic bond increases.

Variation of EP in groups & periodsØ Electropositive nature increases down the group,

as the size of the atom increases.

Ø Electro positivity decreases across a period.

Ø In any period the strong electropositive element is

alkali metal.

Ø Most electro positive element is Cs in periodic table.

Ø The ions of strong electro positive metal do not

undergo hydrolysis.

Metallic and Non-Metallic NatureØ If an element has low electro negativity and high

EP, then it will have high metallic nature.

Ø The groups IA and IIA elements have strong metallic

nature.

Ø Group VIA and VIIA elements have strong non-

metallic nature.

Ø On moving from top to bottom in a group

a) non metallic nature decreases

b) metallic nature increases

Ø On moving from left to right in a period

a) metallic nature decreases

b) non metallic nature increases

Ø Order of metallic nature

Alkali metals > Alkaline earth metals > d-block >

p-block.

eg: 1) The order of increasing metallic character of

Si, Be, Mg, Na, P is: P < Si < Be < Mg < Na.

2) Order of Metallic nature of B, Al, Mg and K is:

K > Mg > Al > B

3) Order of nonmetallic nature of B, C, N, F and Si

is: F > N > C > B > Si

4) The metallic nature of elements in the carbon

family is: Carbon and silicon are non-metals.

Germanium is a metalloid. Tin and lead are metals.

Ø Metals are solids at room temperature except

mercury (Hg).Ga,Cs also have very low melting

points 303K and 302K respectively. so they exists

as liquids at room temperature.

Ø Non-metals are usually solids or gases at room

temperature with low melting and boiling point

(boron and carbon are exceptions).

Ø Some elements in periodic table shows both metallic

and non-metallic nature. They are called metalloids

or semi metals

eg;Silicon.Germanium,Arsenic,Antimony, Tellurium

Acidic and Basic Nature of Oxides:Ø Based on the nature, oxides are clasified into 4 types

1) Basic Oxides or Metal Oxides

2) Acidic Oxides of Non–Metal Oxides

3) Amphoteric Oxides

4) Neutral Oxides



Ø Metal oxides are basic. eg: Na2O, BaO, MgO,

CaO (Basic anhydrides)

Ø IA, IIA group metal oxides are strong bases.

Ø Non metal oxides are acidic. eg: SO2, P

2O

5, CO

2,

P2O

3, NO

2 (Acidic anhydrides)

Ø Oxides of halogens are highly acidic.

Ø Oxides of metalloids are amphoteric.

eg: As2O

3, Sb

2 O

3, TeO

2, GeO

2

Ø Some non-metallic oxides are neutral. They don’t

form acids or bases in water.

eg: CO, N2O, NO etc.,

Ø Some metallic Oxides are amphoteric.

eg : ZnO, Al2O

3, SnO

2 etc.,

Ø Acidic oxides dissolve in water to form acidic

solutions.

eg : 3 2 2 4SO + H O H SO→

Ø Basic oxides dissolve in water to form basic

solutions, known as hydroxides.

eg : 2 2

Na O+ H O 2NaOH→

Ø In groups from top to bottom

a) acidic nature of oxides decreases

b) basic nature of oxides increases

Ø In periods from left to right

a) basic nature of oxides decreases

b) acidic nature of oxides increases

Diagonal RelationshipØ In the periodic table the first element of a group has

similar properties with the second element of the

next group. This is called diagonal relationship.

Ø The diagonal relationship disappears after IVA

group.

Ø The diagonal relationship is due to

i) Similar sizes of atoms or ions

ii) Same electronegativities of the participating

elements

iii) Same polarising power.

Ø Valency is different for diagonally related pair of

elements.

Polarising power of cation α

( )2

ionic ch arg e of cation

ionic radius of cation

Ø The elements present under diagonal relationship

have very close properties.

1) BeO amphoteric, Al2O

3 amphoteric

2) Be2C or Al

4C

3 produce methane on hydrolysis.

Anomalous Properties of Second

period elementsThe first element of each of group in ‘s’ and ‘p’

block except noble gases differ in many aspects

from the other members of their respective group.

eg :1) lithium,beryllium forms covalent compounds

rest of the group members forms ionic compunds.

2) In IIIA group the maximum covalency of boron

is 4 but remaining elements shows maximum

covalency of 6.

3) The first member of p-block elements displays

greater ability to form Pπ– P

π multiple bonds itself

(eg:

C C=

,

C C≡

,

N N=

,

N N≡

) and to other

second period elements (eg:

C O=

,

C N=

,

C N≡

,

N O=

) compared to subsequent

members of the same group.

The reasons for the above anomalous behaviour is

due to their :

(a) Small size

(b) Large (charge/radius) ratio

(c) High electronegativity

(d) Absence of vacant orbitals.

Oxidation - Reduction Ability

Ø Electropositive elements have lower reduction

potenital (RP).

Ø They form stable cations in gaseous state as well as

in aqueous state.

Ø Atoms of these elements are potential suppliers of

electrons.

M

→

Mn+ + ne–

Ø The tendency of an element to supply one or more

electrons is called reduction ability. It is also the

tendency of an element to oxidise itself.

Ø Alkali metals are strong reducing agents, because

the size of metal atoms is more, ionisation potential

is less and each of the atoms have only one electron

in the valency shell.

Ø Alkaline earth metals are also good reducing agents,

but the reduction ability is less than the

corresponding alkali metal.

Variation of Reduction Ability in

Groups & PeriodsØ In a period, reduction ability gradually decreases.

The trend in the reduction ability of third period

element is: Na > Mg > Al Si.


74


Ø In a group reduction ability generally increases.

Caesium is the best reductant among the available

elements in its elementary state because both

sublimation enthalpy and ionisation enthalpy of

caesium are less.

Ø Electronegative elements are non-metals. They

usually have higher electron gain enthalpy and

reduction potentials.

Ø They form stable anions in gaseous state as well as

in aqueous state. Atoms of these elements are

potential acceptors of electrons.

A + ne– → An-

Ø The tendency of an element to gain one or more

electrons is called oxidation ability. It is also the

tendency of an element to reduce itself.

Ø Halogens are strong oxidising agents, because the

size of atoms is less and have only one vacancy in

the valency shell of each atom.

Variation of Oxidation Ability in

Groups & Periods

Ø In a period oxidation ability gradually increases. The

trend in the oxidation ability of third period elements

is : P < S < Cl.

Ø In a group oxidation ability generally decreases. The

order of oxidation ability of halogens is :

F2 > Cl

2 > Br

2 > I

2.

Ø Fluorine is the best oxidant, because dissociation

enthalpy of difluorine is less and hydration energy

is more.

eg: 1) The order of their chemical reactivity in terms

of oxidizing property of F, Cl, O and N is:

F > Cl > O > N

2) The order of oxidising ability of sulphur and

chlorine is : Cl > S

because

Cl Cl e→ −

;

2S e S+ →

Chlorine is better oxidant than sulphur. Electron gain

enthalpy is more for chlorine. Chlorine accepts

electron easily and becomes stable chloride.

Periodic Trends and Chemical

Reactivity

Ø All chemical properties are a manifestation of the

electronic configuration of elements.

Ø The atomic radii generally decrease in a period from

left to right.

As a consequence, the ionisation enthalpies increase

and electron gain enthalpies become more negative.

Ø Since ionisation potentials are less, alkali metals are

very reactive.

Ø Similarly halogens are also very reactive due to high

electron affinity. Thus high chemical activity is

witnessed at the two exteremes and the lowest in

the centre of the periodic table.

Ø Maximum chemcial reactivity at the extereme left is

exhibited by the formation of cation.

This is referred to electropositivity and the elements

act as good reductants.

Ø Maximum chemical reactivity at the extreme right

(not noble gases) is exhibited by the formation of

anion. This is referred to non-metallic nature and

the elements act as good oxidants.

Periodic trends

classification of elements and periodicity in …new.excellencia.co.in › college › web ›...

Documents