1. A substance Consisting of only one kind of atoms is called “an Element” .
2. The concept of elements was introduced by Robert Boyle
3. For elements the concept of symbols was developed by Berzelius.
4. At present around 118 elements are known
5. Among these 94 are natural, and the remaining are man made.
6. Elements coming after 92 atomic number are known as “Trans Uranic Elements” or “Synthetic Elements” and they are “radio active”
7. Total 11 – elements are present in gaseous state H2, O2, N2, F2,Cl2 and Inert Gases (Ozone is also a gas).
8. At 25°c – Hg and Br2 are liquids, at 30°c – Cs, Fr and Ga are liquids (The recently discovered element EKa mercury also said to be in liquid state
9. The first artificially made element is Technicium (Tc)
Historical Development of the periodic table:
Dobereiner’s Triads :
i) He made groups of three elements having similar chemical properties called TRIAD
ii) In Dobereiner triad, atomic weight of middle element is equal to the average atomic weight of first and third element.

Telluric Helix (1862):
Chancourtois arranged the then known elements in order of increasing atomic weights and made a cylindrical table of elements to display the periodic recurrence of properties.

Newland’s Law of Octaves
Newlands proposed a sequence of elements on the basis of their relative atomic weights. If elements are arranged in order of their increasing atomic weights, the eighth element resembled the first, the ninth element resembled the second and so on. Newlands compared his arrangement with repeating musical octave.

i) The properties of Li are similar to 8th element Na,
Be are similar to Mg and so on.
ii) At that time inert gases were not known
Draw back : This rule is valid only upto Ca.

Lother Meyer
The findings of Meyer are :
a) Alkali metals having the largest atomic volumes occupy the maxima of the curve.
b) The alkaline earth metals (Mg, Ca, Sr, Ba) occupy the mid point positions on the descending portions of curve.
c) Halogens occupy position on ascending portions of the curve before inert gases.
d) The transition elements occupy minima of the curve.

Mendeleev’s Classification of elements:
a) Periodic Law:
The physical and chemical properties of the elements are periodic functions of their atomic weights.
b) Mendeleev’s periodic table is also known as short form of periodic table.
c) Mendeleev observed that elements with similar properties have
i) Almost have same atomic weight
Eg: Fe(56), Co (59), Ni (59)
ii) Atomic weights increasing constantly Eg: K(39), Rb (85), Cs(133)
d) Elements are arranged in 12 horizontal rows known as series which are grouped into 6 periods.
e) 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
f) Vertical columns are called groups and there are nine groups (0 to 8th)
g) Leaving 0 and VIII, each group is subdivided into subgroups known as A and B group.
h) Group VIII of the Mendeleef 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
i) Zero group elements were later introduced by Ramsay and Rayleigh
j) 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
i) Eka boron presently known as Scandium
ii) Eka silicon presently known as Germanium
iii) Eka aluminium presently known as Gallium
k) Mendeleev corrected the atomic weight of Berylium, Indium, Osmium, Au and Pt by using corrected valency of elements.
Atomic Wt. = Equivalent Wt. X valency.

i) Some elements with higher atomic weight were placed before low atomic weight elements inorder 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 d) Th-Pa
ii) Position of hydrogen was not made clear.

Modern periodic law and the present form of the periodic table
i) 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.
ii) Atomic number ‘Z’ can be related to frequency of the X-rays emitted by using
=a(Z–b) where a, b are constants for an element
iii) A plot of against Z gives a straight line.
iv) Atomic number has provided a better basis for the periodic arrangement of the elements.

Modern periodic law. Mosley (1921):
Physical and chemical properties of the elements are periodic functions of their atomic numbers and electronic configuration.

Long Form of Periodic Table:
Proposed by Niels Bohr
When elements are arranged in the increasing order of their atomic numbers, the physical and chemical properties are repeated (periodicity) according to their electronic configuration at regular intervals.
Modern periodic table or the long form of periodic table is based on the electronic configurations of the elements.
There are 18 groups and 7 periods in the periodic table.

Periods: (Horizontal Rows)
In periods elements are arranged in the increasing order of their atomic numbers.
The electrons by which an elements differs from its previous element is called as “ differentiating electron”.
In each period, the differentiating electron enters the “S” orbital in the first element and “p” orbital in the last element.
In periods elements are arranged according to the “(n+I)” values order (Aufbau-Rule).
Long form of the periodic table is a graphical representation of the Aufbau- Rule.
Generally every period start with an Alkali metal and ends with Noble gas.

2nd period elements are also known as Bridge elements
Elements of the third period are known as Typical elements (Na, Mg, AI, Si, P, S .and CI).
3rd orbit contains 3s,3p and 3d. But according to (n+l) values order (energy order) 3d comes after 4s,hence accordingly elements with 3d configuration are placed after 4s only in 4th period. (3d series- 1 st transition period Sc(Z=21) to Zn(Z=30))
Elements with 4d configuration from Y(Z=39) to Cd(Z=48) placed in 5th period (2nd Transition series).
Elements with 5d configuration from La (Z=57) and Hf (Z=72) to Hg (Z=80) are placed in 6th period. (3rd transition series)
14 – 4f series elements belongs to 6th period. Ce(Z=58) to Lu (Z=71).
14 – 5f series elements belongs to 7th period. Th(Z=90) to Lr (Z=1 03).
6d – series is incomplete series.
If 7th period is also completed, then the final element of this period would be with an atomic number 118.
Most of the radioactive elements are actinoids.

Groups (vertical columns):
Long form of the periodic table comprises of 18 vertical columns which are divided in to main groups and subgroup as – IA to VIlA, O groups and IIIB, IVB, VB, VIB, VIIB, VIII, IB and lIB groups.
VIII group includes there vertical columns of Fe, Co and Ni (Total 9 elements are present in this group.)
We adopt the 1-18 numbering scheme recomended by IUPAC in 1988.
According to IUPAC, the groups are numbered from 1 to 18 replacing of the older notation of groups IA, IIA, IIIB…………. VIIB, VIII, IB, IIB, IIIA ………..VIIA and O
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 one electron in their outer most orbit,so they are placed in IA group.
Be, Mg, Ca, Sr,Ba and Ra have 2 electron in their outer most orbit, so they are placed in IIA guoup.
Nomenclature of elements with Atomic number > 100

Notation for IUPAC Nomenclature of elements


1. Classification of the elements into following four blocks is decided by the differentiating electron (last filling electron). Block is that in which orbital differentiating electron goes.
a) s – Block Elements        b) P – Block Elements        c) d – Block Elements        d) f – Block Elements
s – Block Elements
1. If shells upto (n -1) are completely filled and the last electron or differentiating electron enters into s-orbital of the outer most orbit (with nth shell) the elements this class are called -s- block element.
2. s- sublevel can accomadate 2- electrons, hence s-block elements are arranged to two groups, group – 1 and group 2
3. General electron configuration is ns1-2,
4. H, Li, Na, K, Rb, Cs, Fr = elements have one electron in their outer shell with nsl general outer shell configuration, they belongs to group -1.

5. Be, Mg, Ca, Sr, Ba and Ra (Alkaline Earth elements) have 2- electrons in their outer shell, with “ns2” general outer shell configuration, they belongs to group-2.
6. Total number of s- block elements are 13.

p- block – elements:
1. If shells upto ( n-1 ) are completely filled and differentiating electron enters into p- orbital of the nth orbit, elements of this class are called “p-block” elements.
2. The general outer shell configuration of p block elements is ns2np1 – ns2np62 (or) ns2np1-6
3. P – block elements are arranged in 6-groups they are from 13th to 18th groups.
a) B, AI, Ga, In and TI – Boron family – 13th group, these elements have 3 electrons in outer shell, with “ns2np1” general outer shell configuration.
b) C, Si, Ge, Sn and Pb – Carbon Family – 14th group, these elements have 4- electrons in outer shell, with “ns2np2” general outer shell configuration.
c) N, P, As, Sb and Bi – Nitrogen Family – 15th group (Pnicogens). These elements have 5- electrons in outer shell, with “ns2np3” general outer shell configuration.
d) O,S, Se, Te and Po – Chalcogens 16th group, these elements have 6 electrons in outer shell, with “ns2np4”, general outer shell configuration.
e) F, CI, Br, I and At – Halogens – 17th group, these elements have 7 electrons in outer shell, with “ns2np5”, general outer shell configuration.
f) He, Ne, Ar, Kr, Xe and Rn – Inert gases – 18th group, Except He(1s2), remaining inert gases have 8 electrons in outer shell with “ns2np6” general outer shell Configuration.

Remember :
1) Infact Helium belongs to s-block, but keeping its chemical inertness, Helium is placed along with
other inert gases in 18th- group.
2) Hence He is a p-block element without p electrons.
3) The first p-block element is Boron [(He)2s22p1]
4) The only group with all gaseous elements is

d-block elements:
If differentiating electrons enters into (n-1)d orbital (i.e-d- orbitals of penultimate shell), the elements of this class are called “d block elements”.
The general electronic configuration of d block elements is (n-1 )d1-10ns1 or 2 (n= outer shell).
d- Block elements are placed between s block and p-block and they are also called transition elements.
d-Block elements are further classified into following transtion series on the basis of which (n-1)d is being filled.
They are classified into 10 groups from group 3rd to 12th
Period           Series                   Configuration                     Elements
4                       3d                    4s2 3d1 to 4s22 3d10        from Sc to Zn
5                       4d                    5s2 4d1 to 5s2 4d10        from Y to Cd
6                       5d                    6s2 5d1 to 6s2 5d10        from La to Hg
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 (La – 6s24f05d1). Therefore “La” belongs to d- block.
Similarly in case of Actinium, the differentiating electron is entering into 6d, instead of 5f (Ac-7s25f06d1). therefore Ac also belongs to d-block.

f-block Elements:
If differentiating electron enters into f-orbital of Anti penultimate i.e.(n-2) shell, the elements of this class are called f-block elements.
The general electronic configuration (n-2)f1-14 (n-1)d0 or 1 ns2 (n = outer shell).
These f- block elements are placed at the bottom of the periodic table in two rows, they are 4 f series and 5 f series.
4f – series – Lanthanoide series – configuration 4f1-14 5d0-1 6s2 from Ce (58) to Lu (71)
4 f- series elements belongs to 6th period and 3rd group.
5f – series elements – Actinoide series configuration 5f1-14 6d0-1 7s2 from Th (90) to Lr (103).)
5 f – series elements belongs to 7th period and 3rd group.

Classification based on chemical properties:
All the elements are divided in to four types on the basis of their chemical properties and electronic configuration.

Type–I Inert gas elements:
1. He, Ne, Ar, Kr, Xe and Rn belonging to “18th” group in the periodic table are called Inert Gas Elements.
2. Except He (1s2), all the other elements have ns2 np6 outer electronic configuration.
3. All are chemically inert due to the presence of stable ns2np6 (octet) configurations in their outer most shell.
4. It is known that heavier elements (Kr, Xe) form true compounds under special controlled conditions with Oxygen and Fluorine. So they are now called Noble gases.
5. All are monoatomic gases.
6. They are also known as Rare gases (or) Aerogens.

Type-II Representative elements or normal elements:
1. These are the elements whose atomic electron shells, except outer shell are completely filled.
2. Excluding “18th” group, remaining s- and p- block elements are called representative elements.
3. Their general outer electronic configurations are ns1-2np1-5.
4. Metals, nonmetals and metalloids are present in representative elements.
5. Atoms of these elements enter chemical combination by losing, gaining or sharing of electrons to attain stable nearest inert gas configuration.
6. In case of representative elements electrons of outer ns and np will take part in bonding.

Type-III Transition elements:
1. These are the elements whose outer most and penultimate orbit is incompletely filled.
2. Elements which have incompletely filled or partly filled d- orbitals either in their atomic state or reliable ionic states are called as transition elements.
3. Their properties are intermediate to s- and p- block elements.
4. The general electronic configuration is ( n-1 )d1-9 ns1-2.
5. 12th group elements Zn (3d104s2), Cd (4d105s2) Hg (5d106s2) are not transition elements (due to the absence of partly filled d- orbitals both in Atomic and ionic states) (Zn, Cd, Hg. – are referred as Non typical Transition Elements).
6. In the case of Transition elements both (n-1) d and ns electrons participate in bonding. These elements show characteristic special properties. Like
1. Variable Oxidation states
2. Formation of coloured ions due to d–d transition
3. Formation of metal complexes
4. Paramagnetism
5. Catalytic activity
6. High m.p., B,p. and densities
7. Interstitial compound formation
8. Alloy formation
9. They are hard and heavy metals
7. These characteristic properties are due to
a. Small size
b. High nuclear charge
c. unpaired electrons in d-orbitals
NOTE: Ni used as a catalyst for Hydrogenation of oils.
Fe used as-a catalyst in Haber’s process
Mo used as a promoter in Haber’s process.

Type-IV Inner-transition elements:
1. These elements have three outer most shells incomplete i.e. – n, (n-1) and (n-2).
2. The f- block elements are called inner transition elements.
3. General configuration (n-2)f1-14 (n-1)d0 or 1 ns2.
4. Since the last two shells have similar configuration these elements have similar physical and chemical properties (ex – the elements show a common oxidation state of +3)
5. There are two series of inner transition elements.

4f – series – Lanthanoide series

4f1-14 5d0-1 6S2 configuration.
5f – series – Actinoide series.
5f1-14 6d0-1 7s2
6. Lanthanoides are rare earths, and all most all Actinoides are radioactive.
7. In periodic table, lanthanoides are present between 57La & 72Hf and
Actinoides are present between 89Ac & 104Rf.

Determination of period, block and group of an element:
Period No:
The period number of the element can be predicted from the principle quantum number (n) of the valence shell.
Eg: Electronic configuration of iodine is 1s2 2s2 2p6 3s2 3p6 4s23d104p65s2 4d10 5p5. Therefore the period number of iodine is ‘5’ as valence shell configuration is 5s2 5p5.
Block No:
The type of orbital which receives last electron is known as block
Eg: An element ‘X’ has its electronic configuration 1s2 2s2 2p6 3s2 3p6 4s23d8. As the last electron enters in d-orbital, there fore it is a d-block element.
Group No:
It is predicted from the number of electrons in the valence shell and penultimate shell as follows.
For s-block elements, group number is equal to the number of electrons in the valence shell after noble core
Eg: An element ‘Y’ having electronic configuration 1s2 2s2 2p6 3s2 3p6 4s2 or [Ar]4s2 has two electrons in valence shell. It is a s-block element and it belongs to group 2.
For p-block element, the group number is equal to 10+number of electrons in valence shell.
Eg: An element ‘z’ with electronic configuration as 1s2 2s2 2p6 3s2 3p6 4s24p3 has five electrons in its valence shell.
:. Its group number is 10+5=15.
It belongs to group VA of periodic table.
For d-block elements, group number is equal to the number of electrons in (n–1)d subshell and valency shell.
Eg: An element ‘A’ having electronic configuration as 1s2 2s2 2p6 3s2 3p6 4s13d10.
So, its group number will be 10+1=11.
In group IA, atomic number of H is 1. Atomic number of other elements will be as follows.
H-1, Li: 1+2=3,           Na: 3+8=11,            K:11+8=19,
Rb: 19+18=37,           Cs: 37+18=55;        Fr: 55+32=87

Periodic trends in properties of elements :
1. When elements are arranged in increasing order of atomic number, elements with similar properties (due to similar outer electronic configuration) at regular intervals in the periodic table.
2. Elements coming at intervals of 2,8,8,18,18,32 will have similar properties and thus grouped in one particular group.
Ex: elements with atomic number 1,3, 11, 19, 37,55 & 87. elements with atomic number 2,4,12,20,38,56 & 88 will have similar properties.
NOTE: Two adjacent elements in a group generally differ by atomic number 2,8,8,18,18,32 etc.
Atomic radius:
1. In atoms, the electron cloud around the nuclues extends to infinity.
2. The distance between the centre of the nuclues and the electron of outer most energy level is called atomic radius.
3. Atomic radius connot be determined directly, but measured from the inter nuclear distance of combined atoms, using X-ray diffraction techniques.
4. Atomic radius depends on
a. Type of bonding
b. Number of bonds (multiplicity of bonding)
c. Oxidation states etc.
5. 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
6. Atomic radius expressed in Angstrom, nanometers, picometer units. 1A0 = 10-1nm ; 1A0=102Pm
7. Crystal Radius – Half of the internuclear distance between the adjacent atoms of a solid metallic crystal is called crystal radius or metallic radius.
Ex: Distance between two sodium atoms is 372Pm. crystal radius of Na = 372/2 = 186Pm.
8. Vander waals radius – Half of the internuclear distance between two non bonded atoms of different molecules which are very close to each other due to vander waals forces is called vander waals radius.
9. The distance between two adjacent chlorine atoms of different CI2 molecules is 360 Pm, Vander waals radius of CI is 180 Pm.
10. Vander waals radius is 40% more than covalent radius.
11. It is used for more under substances in the solid state only
12. Covalent radius : This term is generally used in reference to non – metals.
13. Covalent radius – One half of the internuclear distance of the two atoms held together by a covalent bond is called covalent radius.
14. Ex – Inter nuclear distance between two chlorine atoms in CI2 molecule is 198Pm (Bond Length). Covalent radius of CI is 99Pm
15. Note:- single bond covalent radii are additive in nature.
Ex: a)in Cl2; CI – CI bond distance (Inter-nuclear distance) is 198Pm.
Covalent Radius of CI = 99Pm
b) In Diamond C-C bond distance is 154 Pm
covalent radius of C = 77Pm
In C-CI bond distance observed in a veriety of compounds is about the same as that obtained by adding covalent radii of
C (= 77 Pm) and CI (= 99 Pm)
i.e C-CI bond distance = r(CI) + r(C)
= 99 + 77
= 176Pm
16. The values of single bond covalent radii, is more than double bond covalent radii and triple bond covalent radii
17. Vander waal radius>crystal radius>covalent radius.
18. On moving from left to right across a particular period, the atomic radius decreases up to halogen atoms (17th group) and increases in Inert gases.
19. In a given period, alkali metal is the largest and halogen is the smallest.
20. For atoms of Inert gases, only vander waal radius is applicable.
21. In groups from top to bottom, the atomic radius increases gradually due to the increase in the number of orbits.
Variation of atomic radius in transition elements :
22. 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 the screening effect of (n-1 )d-electrons.
23. Hence, the atomic radius decreases slightly as we move from left to right in a transition series.
Variation of atomic radius and ionic radius in Inner transition elements :
24. In Lanthanoides (Ce – Lu) the atomic and ionic radii decrease steadily. This steady decrease in atomic and ionic radii is known as “lanthanoide Contraction”.
25. There is also a decrease in radii in Actinoides (Th – Lr) that is called Actinoide Contraction.
26. The contraction is due to the fact that – f-orbitals
are not capable of providing effective shielding for the valence electrons from nuclear attraction.
27. Consequences of lanthanoide contraction:
a. Atomic sizes of 4d and 5d transition elements become almost equal, due to which their Properties are very close
b. Zr and Hf ; Nb and Ta ; Mo and W resemble very closely
c. The crystal structure and other properties of lanthanoides are very similar
d. Separation of lanthanoides is not easy from their mixture.
e. Super heavy metals of P–block exhibit inert pair effect.
Eg : TI, Pb, Bi
28. Ionic Radius:- The effective distance from the nucleus of an ion upto which it exerts its influence on its eIectron cloud is called ionic radius.
a. Radius of cation is less than its parental atom. Size of Na+ < Na.
b. As the charge on cation increases, its radius decreases due to increase in effective nuclear charge
Size of Mg2+ < Mg1+ < Mg.
c. As the charge on anion increases, its radius increases due to decrease in effective nuclear charge.
Size of O-2 > O-1 > O
d. If an atom is capable to form anion and cation, then order of radius – Anion> Neutral atom> cation.
Ex : I > I > I+.
29. The species (atoms or ions) having the same number of electrons are known as iso – electronic species.
30. In iso electronic series, the size decreases with increase in nuclear charge
decreasing order of size
Species: C-4 > N-3 > O-2 > F-1 > Na+ > Mg2+ > Al3+ > Si4+
No.electrons   10    10     10    10    10     10    10    10
No.protons       6      7        8     9     11     12    13    14


Ionization enthalpy (Ionization potential):
1. It is an atomic property, which is shown by the atoms of all the elements.
2. Ionization Energy: The amount of energy required to remove the most loosely bound electron (ie. outer – most shell electron) from isolated neutral gaseous atom in its lower energy state (ie. ground state) to convert it into a gaseous cation is called Ionization energy. This is also called first Ionization Energy. (IE1)
3. Ionization Potential :- The amount of potential or voltage required to remove the most loosely bound electron of an atom is ionization potential.
4. IE is measured in electron volts (ev) (or) (or) k.J/mole. leV/atom = 23.06k.Cal/mole = 96.45KJ/mole
5. Energy required to remove an electron from unipositive ion to convert it into dipositive ion is 2nd ionisation energy (IE2)
6. Energy required to remove an electron from dipositive ion to convert it into tripositive ion is IE3 – so on.
The magnitude of successive IE are in the order:
lE1 < IE2 < IE3 < IE4 < ……
M(g) + lE1 → M+(g) +e
M+(g) + IE2 → M+2(g) +e-
M2+(g) + IE3 → M3+(g) +e-
7. Ionization potential depends on:
a. With increase in the atomic size “IP” decreases due to decrease in attractive force of nucleus on outer most orbit electrons.
b. With increase in shielding effect or screening effect. IP decreases.
c. Order of screening power of orbitals. s > p > d  > f
d. As the positive charge on cation increases, IP increases.
e. As the -ve charge on anion increases, IP decreases.
f. Penetration power of different orbitals is in the order of s > p > d > f
g. IP order of electrons of different orbitals of same orbit.
IP of s-electrons > IP of p- electrons> IP of d-electrons > IP of f – electrons
h. IP more for atoms with exactly half filled and completely filled orbitals
Ex: lE1 of N > IE1 of O
lE1 of P > IE1 of S
i. Atoms of inert gases have highest IP values due to the presence of completely filled orbitals.
j. In the graph showing relation between IP and atomic number, the inert gases appear at the peaks (maxima) and alkali metals appear at the bottom (minima).
k. In any period an Alkali metal atom have lowest IP and Inert gas element has highest IP.
l. In periods from left to right side IP increases, due to decrease in atomic size and increase in effective nuclear charge.
m. 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.
n. IE order 2nd period elements
IE1 – Li < B < Be < C < O < N < F < Ne
IE2 – Be < C < B < N < F < O < Ne < Li
o. IE order 3rd Period
IE1– Na < Al < Mg < Si < S < P < Cl < Ar
IE2– Mg < Si < Al < P < S < Cl < Ar < Na
p. Element with
a) Lowest IP- Cs
b) Highest IP- He
q. IE1 of Be greater than B due to
a) Completely filled s-orbital in Be
b) More Penetration of s-orbitals
r. Knowledge of successive IE can be used to find the number of valence electrons.
s. For alkali metals the IE2 shows sudden jump.
t. For alkaline earth metals, the IE3 shows sudden jump.
u. The number of IE possible for an atom of an element is equal to its atomic number.

Electron gain enthalpy:
1. It is an atomic property which gives us an idea of
the tendency of the element to accept the electron to form an anion.
2. The amount of energy released on adding an extra electron from out side to an isolated gaseous neutral atom in its lowest energy state to convert it into gaseous anion is called Electron Affinity.
(also known as EA1)
X(g) + e = X1-(g) + EA
3. EA is measured in eV/atom, K.Cal/mole, K.J/mole
4. EA2 is slightly endothermic, because of repulsions between in uninegative ion and extra electron. To over come the repulsions, some energy is supplied from out side.
Ex : O(g) + Ie → O1- + EA1 (EA1 = –48eV)
O1-(g) + Ie + EA2 → O2- (EA2 = + 8.77eV)
5. EA can be calculated from Born – Haber Cycle.
6. Be group elements have completely filled orbitals and hence the addition of any extra electron from out side to these atoms is not possible. Therefore they have practically zero EA.
7. Noble gases have most stable ns2np6 configuration. Hence their EA values are practically large positive.
8. For N, P – due to half filled orbitals, they have some extra stability hence their EA values are close to zero (very small values).
9. In groups, EA decreases from top to bottom as the atomic size increases.
10. In p-block, the second element of every group has greater EA that the first one
11. In VIl-A group EA of CI > EA of F
VIA group EA of S > EA of O
VA group EA of P > EA of N
IV A group EA of Si > EA of C
12. EA of F (–328 KJ/mole) < EA of CI (–349 KJ/mole)
This is due to
a. Smaller size of F-atom
b. Strong inter electronic repulsions.
13. In a period from left to right side EA increases due to decrease in size of atoms and increase in the nuclear charge.
14. Note:- EA of a neutral atom=IE of its
uninegative ion.
15. Note:- IE of a neutral atom = EA of its uni positive ion.
16. Among all the elements chlorine has the maximum EA.
17. The metal which has higher EA is gold.

Electro negativity:
1. It is the property of an atom in a molecule.
2. The tendency of an atom to attract the shared electron pair towards itself in a molecule is called EN.
3. Pauling Scale : EN of elements are calculated from the values of bond energies.
The elements EN values are calculated based on hydrogen value. Hydrogen EN value is 2.1
Electro negativity values (on Pauling scale) Across the periods
This is the most widely used scale and is based upon bond energy data. According to Pauling, the difference in electro negativity of two atoms A and B is given by the relationship.

  • By giving a reference value of 2.1 to H, the electronegativities of almost all the elements have been calculated. On this scale, the maximum EN value is 4.0 for fluorine.
  • Mulliken EN values are approximately 2.8 times greater than Pauling EN values.

The Pauling and Mulliken scales are related to each other by the relation
EN on Pauling scale =

Electronegativity values (on Pauling scale) Down a family

4. According to Mulliken scale EN is the average of IE and EA.

5. EN concept is not applicable for Inert gas elements.
6. In groups from top to bottom EN decreases.
7. In periods from left to right EN increases.
8. Highest EN element is F(4.0).
9. Next to O, oxygen has EN (3.5).
10. Least EN element is Cesium (0.7).
11. EN values are used to know the nature of chemical bond.
12. If EN differance is less than 1.7 the bond is covalent in nature, equal to 1.7 the bond is 50% ionic in nature, more than 1.7 the bond is ionic in nature.

Oxidation state:
1. The possible charge with which an atom appears
in a compound is called its oxidation state.
2. In s-block elements, oxidation state is equal to group number. For alkali metal atom +1 for alkaline earth metal atom +2.
3. Oxidation state may be positive or negative or zero.
4. p-block elements show multi valency, then oxidation state change by two numbers.
5. The stable oxidation state of Thallium is + 1. it is due to inert pair effect.
6. In 14th group, +2 is more stable than +4 for Lead due to inert pair effect.
7. In 15th group, +3 is more stable than +5 for Bismuth due to inert pair effect.
8. Group 14th elements show+4 and+2 oxidation states.
9. Group 15th elements show+5 and+3 oxidation states.
10. The general oxidation state of group 16th is -2.
11. The most electronegative element Fluorine show -1 oxidation state only (in its compounds).
12. The common oxidation state of d – block elements
is +2. All transition elements show variable valencies.
13. Ruthenium, Osmium and Xenon exhibit maximum oxidation state +8 (+8 is stable only for Osmium).
14. The common oxidation state of f – block elements is +3 due to their outer electron configuration ns2(n-1)d1.
15. Maximum oxidation state of an element never exceeds its group number
16. The oxidation state of metal in carbonyls is zero.

Electro positive nature:
1. The tendency of an element to loose an electron is called electro positivity.
2. It is the converse of electro negativity.
3. As electropositivity increases, metallic character increases.
4. The smaller the ionisation energy or ionisation potential the greater is the electro positivity.
5. As electropositive nature increases, capacity to form ionic bond increases.
6. Electropositive nature increases down the group, as the size of the atom increases.
7. Electro positivity decreases across a period.
8. In any period the strong electropositive element is alkali metal.

Metallic and non – metallic nature:
1. If an element has low electro negativity, and high EP, then it will have high metallic nature.
2. The groups 1st and 2nd elements have strong metallic nature.
3. Group 16th and 17th elements have strong non-metallic nature.
4. On moving from top to bottom
a) non-metallic nature decreases
b) metallic nature increases
5. On moving from left to right in a period
a) metallic nature decreases
b) non-metallic nature increases
6. Order of metallic nature
Alkali > Alkaline earth > d – block > p – Block. .

Acidic and basic nature of oxides:
1. Metal oxides are basic. Eg:Na2O,BaO,MgO,CaO.
2. 1st, 2nd groups metal oxides are strong bases.
3. Non metal oxides are acidic
SO2, P2O5, CO2, P2O3, NO2
4. Oxides of halogens are highly acidic.
5. Oxides of metalloids are amphoteric
Eg : As2O3, Sb2O3
6. Acidic oxides dissolve in water to form acidic solutions.
7. Basic oxides dissolve in water to form basic solutions, known as hydroxides.
8. In groups from top to bottom
a) acidic nature of oxides decreases
b) basic nature of oxides increases
9. In periods from left to right
a) basic nature of oxides decreases
b) acidic nature of oxides increases

1. Valency is combinining capacity of an element
2. It is the number of hydrogen atom or double number of oxygen atoms with which the element combines
Ex : 1) HCl , valency of Cl is 1
2) H2O valency of O is 2
3) NH3 , valency of N is 3
4) Al2O3 valency of Al is 3 (An atom of Al combines with 1.5 atoms of oxygen)
3. Valency is represented by its valency electrons
4. Valency of an element is represented by its group number G or ( 8-G )
5. Valency with respect to hydrogen increases from 1 to 4 and then decreases in period
6. Valency with respect to oxygen increases from 1 to 7
7. Transitional elements show varaible valencies
8. Valency of the elements of a group is usually the same as all elements possess same number of valency electrons.

Diagonal relationship:
1. 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.

2. The diagonal relationship disappears after IV group:
3. The diagonal reletionship is due to similar sizes of atoms or ions and electronegativities of the participating elements and same polarisation power.

4. The elements present under diagonal relationship have very close properties.
1. BeO amphoteric, AI2O3 amphoteric
2. Be2C or AI4C3 produce methane on hydrolysis.
Anamalous properties of 2nd period elements
The first element of each of the groups 1st (Li) and 2nd (Be) and groups 13th – 17th (B to F) differs in many respects from the other members of their respective group
The anamalous behaviour is attributed to their small size, large charge/radius ratio and high electronegativity of the elements.
The first member of the p-block elements displays greater ability to form pπ – pπ multiple bonds to itself and to other second period elements compared to subsequent members of the same group.
Ex : C = C, C Ξ C, N Ξ N, C = O, C Ξ N, N = O

Periodic trends and chemical reactivity :

** Inert gas elements are neither metallic nor non metallic.

1. Gaseous elements are H2, N2, O2,F2,Cl2, He, Ne, Ar, Kr , Xe and Rn ( II elements)
2. The liquid elements are Br2 , Hg , Ga (27ºC) and Fr ( Radioactive at around 25ºc)
3. Most abundant element in earth’s crust is oxygen
4. Most abundant element in the atmosphere is nitrogen
5. Most abundant metal in earth’s crust is Aluminium
6. Most abundant element in the sun or in the universe is hydrogen
7. The only non metal of s-block is hydrogen
8. Most active metal is Francium ( Caesium)
9. Most active non metal is Fluorine
10. Lightest metal is Lithium
11. Lightest non metal is hydrogen
12. The first man made ( synthetic ) element is Technetium ( d-block , Eka manganese)
13. Synthetic p-block element is At, s-block element is Fr
14. Elements after uranium are called trans uranic elements. ( man made)
15. Most of the transuranic element were synthesised by G.T seaborg and his coleagues
16. Elements of Zn group (II-B) are called volatile metals ( Zn,Cd,Hg have low boiling points)
17. Coinage metals are copper, silver and gold (I-B)
18. Ag, Au, Pt are called noble metals because of their low reactivity.
19. Zr and Hf are chemically similar because of almost same atomic size – Lanthanoide contraction
20. Self protective metal is Aluminium
21. Amphoteric metals are Beryllium ,Tin, Aluminium, Zinc and Lead
22. The artificial element present in Lanthanoides is prometheum.
23. As,Sb, Si,Ge,Te,At are metalloids
24. Rogue element is Hydrogen
25. The metals which occur in native state in nature are Ag, Au
26. Heaviest element exisitng in earth’s crust is uranium.
27. The densest metal is Osmium
28. The most melleable metal is Au
29. The hardest known element is carbon in the form of Diamond.
30. The hardest metals are Tungsten and Molybdenum
31. The metals that form acidic oxides are Cr and Mn
32. Metal with highest electrical conductance is silver
33. Metal with highest melting point is Tungsten (w) ( 3380ºc)
34. The element with highest melting point is carbon (Diamond 3750ºc)
35. The metal with lowest melting point is Hg (-38ºc)
36. The metal with widest liquid range (30º – 270ºc) is Gallium
37. Gallium is used as a thermometric liquid
38. The element with lowest melting point and boiling point is Helium
39. The element with highest catenation property is carbon
40. The element which is commonly known as ‘ liquid silver’ is mercury
41. The metal obtained by heating its ore is mercury
42. An alloy of mercury is called an amalgam
43. The metals which donot form amalgam are Iron and Platinum
44. The element present in chlorophyll is Mg
45. The element present in haemoglobin is Iron
46. The metal used in photo electric cells are Cs, K
47. The element used as nuclear fuel is U235
48. Ge , Ga ,B and Si are semi conductors
49. Cobalt is present in vitamin B12
50. Antimony is used in the treatment of kalaazar
51. Non- metallic conductor is graphite
52. Most abundant element of human body is oxygen
53. Lightest radioactive element is hydrogen ( Tritium)
54. Element used in cancer therapy cobalt -60
55. Metals that get passivated with HNO3 are Fe ,CO, Ni, Al, Cr
56. Metal with maximum no of isotopes -Sn
57. The metal which has higher E.A. is gold

Additional information :

Effective nuclear charge:
Due to screening effect the valency electron experiences less attraction towards nucleus. This brings decrease in the nuclear charge (z) actually present on the nucleus. The reduced nuclear charge is called effective nuclear charge and is represented by z*. It is related to the actual nuclear charge by the following formula.
z* = (z – σ). where σ is screening constant.
z is nuclear charge
The magnitude of ‘σ’ is determined by the slater’s rules. The contribution of inner electrons to the magnitude of σ is calculated in the following ways.

For ns and np orbital electrons:
Write the electronic configuration of element in the following order and group them as
(1s) (2s,2p) (3s,3p) (3d) (4s,4p) (4d,4f) (5s,5p) (5d,5f) (6s,6p) etc
Electrons to the right of the (ns,np) group contribute nothing to the screening constant.
All other electrons in the (ns,np) group contribute to the extent of 0.35 each to the screening constant
All the electrons in the (n-1) shell contribute 0.85 each to the screening constant.
All the electrons in the (n-2) shell or lower contribute 1.0 each to the screening constant.

For d- or f-electrons:
Rules 1 to 3 remain the same but rules 4 and 5 get replaced by the rule 6
All electrons in the groups lying left to (nd, nf) group contribute 1.0 each to the screening constant.
Eg 1:- Calculation of screening constants of alkalimetals for valency electrons.
Li        2, 1                            = 2 × 0.85                  = 1.7
Na      2, 8, 1                        =8 × 0.85 + 2 × 1       = 8.8
K        2, 8, 8, 1                    =8 × 0.85 + 10 × 1     = 16.8
Rb      2, 8, 18, 8, 1              =8 × 0.85 + 28 × 1     = 34.8
cs       2, 8, 18, 18, 8, 1        =8 × 0.85 + 46 × 1     = 52.8
Eg 2:- Calculation of screening constants of members of second period for valency electrons.
Li        2, 1        = 2 × 0.85                               = 1.7
Be      2, 2        = 0.35 + (2 × 0.85)                  =2.05
B        2, 3        = (2 × 0.35) + (2 × 0.85)         = 2.40
C        2, 4        = (3 × 0.35) + (2 × 0.85)         = 2.75
N        2, 5        = (4 × 0.35) + (2 × 0.85)         = 3.10
O        2, 6        = (5 × 0.35) + (2 × 0.85)         = 3.45
F        2, 7        = (6 × 0.35) + (2 × 0.85)          = 3.80
Ne      2, 8        = (7 × 0.35) + (2 × 0.85)         = 4.15
Eg 3:- Calculation of screening constant in Zn
a) For 4s electron
The electronic configuration of zn(30) is
(1s2) (2s22p6) (3s23p6) (3d10)(4s2)
σ = (1 × 0.35) + (10 × 1) + (18 × 1.85) = 25.65
b) For 3d electron
σ = (9 × 0.35) + (18 × 1) = 21.15
Calculation of effective nuclear charge in 2nd period elements
II period          z            σ              z* = (z – σ)
Li                    3           1.7                   1.3
Be                  4           2.05                 1.95
B                    5           2.40                 2.60
C                    6           2.75                 3.25
N                    7           3.10                 3.90
O                    8           3.45                4.55
F                    9           3.80                 5.20
Ne                 10          4.15                 5.85
It is observed that magnitude of effective nuclear charge increases in a period when we move from left to right.

Calculation of effective nuclear charge in IIA group elements
IIA group               Be           Mg           Ca          Sr           Ba
z                            4             12            20           38           56
s                            2.05        9.15         17.15      35.15      53.15
z* = (z – σ)            1.95        2.85          2.85        2.85       2.85
It is observed that magnitude of effective nuclear charge remains almost the same in a group of normal elements.

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