periodic table, modern periodic law, nomenclature of elements greater than 100,electronic configuration and types of elements,periodic trends in properties of elements.ionization enthalpy, effective nuclear charge, electronegativity, s, p d and f block elements, covalent radius, ionic radius, predicition of group, period and block, electron gain enthalpy, periodic trends and chemical reactivity

1. Notes-by
SAMPATHU ARUNKUMAR

2.  Modern periodic law and the present form of
periodic table, periodic trends in properties
of
elements ‐atomic radii, ionic radii, inert gas
radii, Ionization enthalpy, electron gain
enthalpy, electronegativity, valency.
Nomenclature of elements with atomic number
greater than 100.
CLASS:11 CBSE SHORT NOTES

3.  Modern periodic law states that physical and
chemical properties of the elements are
periodic function of the atomic numbers .
 If the elements are arranged in order of their
increasing atomic numbers, the elements
with similar properties are repeated after
certain regular intervals.

4.  Atomic mass depends upon the number of
protons and neutrons in the nucleus.
 The physical and chemical properties could
be different depending upon the number of
electrons and the electronic configuration in
any atoms.
 Cause of periodicity:
The properties of the elements are repeated
after certain regular intervals when these
elements are arranged in order of their
increasing atomic numbers.

5.  The cause of periodicity in properties is the
repetition of similar outer electronic
configuration after certain regular intervals.
 All the elements of group IA or 1 i.e alkali metals
have the similar outer electronic configuration.
 All the halogens i.e elements of group VII A or
17 have similar outer electronic configuration.
ns2np5. hence possesses similar properties.
 All elements of group 18 have similar electronic
configuration i.e ns2np6 and posses similar
properties.

6.  The table which is based upon the electronic
configuration of the elements is called the
present form or the long form of periodic table.
 The long form of periodic table consists of 18
vertical columns(groups)and 7 horizontal
rows(periods).

9. Electronic configuration of elements in period:
The total number of elements in each period is
twice the number of orbitals available in the
energy level that is being filled.
(i) In first period filling of electrons in the first
energy shell i.e(K shell), n=1. Since this energy
shell has only 1 orbital i.e 1s which can
accommodate only 2 electrons, therefore, first
period has only 2 elements.

10.  In second period filling of electrons in the
second energy shell(L-shell), n=2. This shell
has 4 orbitals(one 2s and three 2p) which can
accommodate 8 electrons, therefore second
period contains 8 elements. It starts with
Lithium(Z=3) and ends at neon(Z=10).

11.  In third period filling of electron in the third
shells. n=3. This shell has 9 orbitals(one 3s,
three3p and five 3d orbital ). The 3d orbital
have higher energy than 4s orbital. Therefore
3d orbitals are filled only after filling of 4s
orbital.
 Third period involves the filling of only 4
orbitals and thus contains 8 elements. It
starts with sodium(Z=11) and ends at argon
(Z=18).

12.  In fourth period filling of electrons in the fourth
energy level, n=4. It starts with potassium (Z=19)
and ends at calcium(Z=20).
 After filling 4s orbitals, the filling of five 3d
orbitals begins since the energy of 3d orbital is
lower than those of 4p orbitals but higher than
that of 4s orbital.
 The filling of 4d and 4f orbital does not occur in
this period since their energies are higher than
that of even 5s orbital.
 The filling of the 3d orbital starts from
scandium(Z=21) and ends at Zinc(Z=30).

13.  These 10 elements constitute the 3d
transition series.
 The filling of 4p orbital begins at
gallium(Z=31) and ends at krypton(Z=36)
which has the outer electronic configuration
as 4s23d104p6.
 In the 4th period, the filling of only 9
orbitals(one 4s, five 3d and three 4p) occurs
which can accommodate at the maximum 18
electrons.
 Therefore 4th period contain 18 electrons
from potassium to Krypton.

14.  The fifth period also contains 18 elements
since only 9 orbitals (one 5s, five 4d and three
5p) are available for filling with electrons.
 It begins with rubidium(Z=37) in which one
electron enters 5s orbital.
 The filling of 4d orbital starts at
yittrium(Z=39) and ends at cadmium(Z=48)
These ten elements constitute 4d transition
series. Filling of 5p orbitals starts at
indium(Z=49) and ends at xenon(Z=54).

15.  The sixth period corresponds to the filling of 6th
energy level i.e n=6. Only 16 orbitals( one 6s, 4f
and 5d ,6p) are available for filling with
electrons, therefore 6th period contains 32
elements.
 It begins with caesium(Z=55) in which one
electron enters the 6s orbital and ends up with
radon(Z=86) in which the filling of 6p orbital is
complete.
 After the filling of 6s orbital, the next electron
enters the 5d orbital and therefore the filling of
seven 4f orbitals begins with Cerium(Z=58) and
ends up with lutetium(Z=71). These 14 elements
constitutes the first inner transition series called
lanthanides or lanthanoids.

16.  Filling of 5d orbitals which started at
lanthanum continuous from hafnium(Z=72)
till it is filled at mercury(Z=80). These 10
elements constitutes the 5d transition series.
After the filling of 5d orbitals, the filling of 6p
orbitals starts at thallium(Z=81) and ends at
the radon(Z=86).

17.  The seventh period corresponds to filling of 7th
energy shells n=7. It also contain 32 elements
corresponds to the filling of 16 orbitals(1-7s,
5-5f, 5-6d and 3-7p)
 After the filling of 7s orbital, the next two
electrons enters the 6d orbitals and therefore
the filling of seven 5f orbitals begin with
proactinium(Z=91) and ends up with
lawrencium(Z=103).

18.  Thorium does not have any electron in the 5f
orbital, yet get it is considered to be a f block
element since its properties resemble more
the f block element than the d block
elements.
 These 14 elements from thorium(Z=90) to
lawrencium(Z=103) constitute the second(or
5f) inner transition series which is called as
actinides are actinoids.

19.  Filling of 5d orbitals which started at actinum
(Z=89) continues till it is completed at these
Uub(Z=112). These 10 elements constitute
the 6d transition series. The filling of 6d,
orbital the filling of 7p orbitals begins at
Uut(Z=113) which ends at Uut(Z=118) which
belongs to noble gas family.
 The first three periods containing 2,8,8
elements are known as short periods while
the next three periods containing 18, 18, 32
elements are called Long periods.

20.  Elements in the long form of periodic table
have been divided into four blocks s,p,d andf.
 S block elements:
 Elements in which the last electron enters the
s orbital of their respective outermost shells
are called s block elements.
 Elements of group 1(hydrogen and alkali
metals) and group 2 (alkaline earth metals)
and helium comprises s block element.
 There are 14 s-block elements in the periodic
table.

21.  General electronic configuration of S- block
elements: ns1-2 where n = 2-7
 Helium belongs to S- block but its
positioning with in the p block along with
other group 18 elements is justified because
it has a completely filled valence shell and
thus exhibits properties characteristic of
other noble gases.

22.  Characteristics of s-block elements:
 They are soft metals with low melting and
boiling points.
 They have low ionisation enthalpies and are
highly electropositive.
 They lose the valence electrons readily to
form +1 and +2 ions.
 They are very reactive metals. The metallic
character and reactivity increases as we move
down the group. Because of high reactivity
they are never found pure in nature.

23.  The compounds of s-block elements with the
exception of beryllium are predominantly
ionic.
P-block elements:
 Elements in which the last electrons enter any
one of the three p orbitals of their respective
outermost shell are called p-block elements.
 The outer shell electronic configuration vary
from ns1np6 to ns2 np6 as we move from
group 13 to group 18.

24.  Elements of group 13-18 in which p orbitals
are being progressively filled in are called p-
block elements. There are 36 p-block elements
in the periodic table.
 Group 18 are called noble gases, except
helium all the elements have closed shell ns2
np6 electronic configuration in the outermost
shell.
 All the energy levels which are occupied are
completely filled and this stable arrangement
of electrons cannot be easily altered by
addition or removal of electrons.

25.  Group 18 elements are highly stable and
hence exhibit very low chemical reactivity.
 The elements of group 17 are called halogens
while those of group 16 are called
chalcogens.
 Elements have high negative electron gain
enthalpy and hence readily add one or two
electrons to attain the stable noble gas
configuration.

26.  Characteristics of P-block elements:
 P block elements include both metals and non-
metals but the number of non-metals is much
higher than that of metals.
 Metallic character increases from top to bottom
within a group and non metallic character
increases from left to right along a period.
 The ionisation enthalpies are relatively higher
as compared to those of s block elements.
 They mostly form covalent compounds.
 Some of them show more than 1 oxidation
state in their compounds.

27.  Their oxidising character increases from left to
right in a period and reducing character
increases from top to bottom in a group.
 D-Block elements:
 Elements in which the last electron enters any
one of the five d orbitals of their respective
penultimate shells are called d-block element.
 General outer shell electronic configuration of d-
block elements(n-1) d1-10 ns0-2, where n = 4 – 7.

28. The d –block elements are also called transition
elements.
All these elements are further divided into four
series called 1st, 2nd , 3rd and 4th transition series.
The first transition series forms a part of the
fourth period of the long form of periodic table.
It contains 10 elements from scandium(Z=21) to
zinc(Z=30) in which 3d orbitals are being
progressively filled in.
The second transition series -5th period contains
10 elements from Y(Z=39) to cd(Z=48).

29. The third transitions series which form a part of
the sixth period also contains 10 elements i.e
(La=Z=57 to Hg-Z=80).
The fourth transition series which forms a part
of 7th period also contains 10 elements.
Ac(Z=89) to Cn-Z=112.
There are 40 d block elements in the periodic
table.

30.  They are hard, malleable and ductile metals
with high melting and boiling point.
 They are good conductors of heat and
electricity.
 Their ionisation enthalpy are between s and p
block elements.
 They show variable oxidation states.
 They form both ionic and covalent
compounds.
 Their compounds are generally coloured and
paramagnetic.

31.  Most of the transition metals such V, Cr, Mn,
Fe, Co ,Ni ,Cu etc and their compounds are
used as catlayst.
 Most of the transition metals form alloys.

32. f-Block elements:
 The elements in which the last electron enters
any one of the 7 f-orbitals of their respective
anti -penultimate shell are called f block
elements.
 General outer shell electronic configuration of
f block elements: (n-2) f 0-14 (n-1) d0-10 ns2
 There are two series of f block elements each
containing 14 elements. These are placed at
the bottom of the periodic table.

33.  The elements of the first series i.e.
cerium(Z=58) to lutetium(Z=71) which form
a part of the sixth period are called
as lanthanides or lanthanoids.
 since all these elements follow lanthanum in
the periodic table and also closely resemble
lanthanum in their properties.
 These are also called rare Earth elements
since they occur scarcely in the earth crust. In
lanthanides ,4f orbitals are been
progressively filled in.

34.  The elements of the second series i.e.
thorium( Z=90) to lawrencium(Z=103) which
forms are part of the incomplete 7th period
are called actinides are actinoids since all
these elements follow actinium in the
periodic table and also closely resemble
actinium in their properties. 5f orbitals are
being progressively filled in.

35.  The first 3 elements i.e. Thorium,
protoactinium and uranium occur in nature
but the remaining 11 elements i.e.
neptunium(Z=93) to lawrencium(Z=103) have
been prepared artificial through nuclear
reactions.

36.  Characteristics of f block elements:
 1)They are heavy metals
 2)They have high melting and boiling point
 3)The show variable Oxidation State
 4)Their compounds are generally coloured
 5)They have high tendency to form
complexes
 6)Most of the elements of the actinide series
are Radioactive.

37.  Prediction of period, group and block of a
given element
 The period of an element corresponds to the
principal quantum number of the valence
shell.
 The block of an element corresponds to the
type of orbital which receive the last electron.
 The group of an element is predicted from
the number of electrons in the valence shell
or/and penultimate shell as follows:

38. a)For s block elements ,group number is equal
to the number of valence electrons.
b) For p block elements ,group number is equal
to 10+number of electrons in the valence shell.
c)For d block elements ,group number is equal
to the number of electrons in a (n-1) sub shell
+ the number of electrons in valence shell.

39. Write the electronic configuration of the
element with atomic number 29. Predict the
period, group number and block to which it
belongs.
Answer : Z=29 1s2 2s2 2p6 3s2 3p6 3d10 4s1
 Elements receive the last electron in the 3d
orbital ,therefore, it belongs to d block
elements and its group number = No. of
electrons in the penultimate shell and valence
shell = 10+ 1 =11
 The period of the element = No. of principal
quantum number of the valence shell = 4th

40.  The properties which are directly or indirectly
related to their electronic configuration and
which show a regular gradation when we
move from left to right in a period or from
top to bottom in a group are called periodic
properties.

41. Atomic radius:
 The distance from the centre of the nucleus
to the outermost shell containing electrons.
 The distance from the centre of the nucleus
to the point up to which the density of the
electron cloud is maximum.
Types of atomic radii:
 1) Covalent radius
 2) Van der waals radius
 3) Metallic radius

43. Variation of Atomic radius across a period:
 The atomic size generally decreases across a
period .
 It is because within the period the outer
electrons are in the same valence shell and
the effective nuclear charge increases as the
atomic number increases resulting in the
increased attraction of electrons to the
nucleus.

44. Variation within a group:
 The atomic radii of elements increases with
increase in atomic number as we move from
top to bottom in a group.
 As we move down the group the principal
quantum Number increases. A new energy
shell is added at each succeeding element
and the valence electrons lie farther and
farther away from the nucleus. As a result the
attraction of the nucleus for the electron
decreases and hence the atomic radius
increases.

47.  The removal of an electron from an atom
results in the formation of a cation, whereas
gain of an electron leads to an anion.
 The ionic radii can be estimated by
measuring the distances between cations and
anions in ionic crystals.

48.  A cation is smaller than its parent atom
because it has fewer electrons while its nuclear
charge remains the same.
 The size of an anion will be larger than the
parent atom because the addition of one or
more electrons.
 For ex: Ionic radius of fluoride ion(F-) is 136pm
where as the atomic radius of fluorine is only
64pm.
 The atomic radius of sodium is 186pm
compared to the ionic radius of 95pm of Na+ .

49.  Ions which contain same number of electrons
we call them isoelectronic species.
 Ex:O2-,F-, Na+, Mg2+ have same number of
electrons 10.

50.  The ionic radii of element exhibit the same
trend as the atomic radii.
 As nuclear charge increases, the force of
attraction by the nucleus on the electrons
also increases. As a result ionic radii
decreases.

51.  Ionization Enthalpy:
 The minimum amount of energy required to
remove the most loosely bound electron from
an isolated gaseous atom so to convert it into
gaseous cation is called ionisation enthalpy.
 It is expressed in the unit of kj/mol.
 It is represented by ΔH

52.  The energy required to remove the most
loosely bound electrons from the isolated
gaseous atom is called its first ionisation
enthalpy and is denoted by Δ H 1 .
 M (g) + Δ i H 1 ———–> M + (g) + e – (g)
 The energies required to knock out second
and third electrons are called second and
third ionisation energies.
 M + (g) + Δ i H 2 ———–> M 2+ (g) + e – (g)
 M 2+ (g) + Δ i H 3 ———–> M 3+ (g) + e – (g)

53.  When one electron has been removed from
the neutral gaseous atom the positively
charged ions formed has 1 electrons less
than the number of protons in the nucleus.
As a result the electrostatic attraction
between the nucleus and the remaining
electrons in the cation increases i.e. effective
nuclear charge increases. The positive ion
holds its remaining electrons more firmly.
Therefore ,the energy required to remove
another electron from this positively charged
Ion or second electron from the neutral atom
must be higher than the first.

54.  Factors governing the ionization enthalpy:
 Factors governing the ionization
enthalpy
1) Nuclear charge
2)Atomic size
3)Penetration effect of the electrons
4)Screening effect of inner electrons
5)Effect of exactly half filled and completely
filled orbitals

55.  Ionization enthalpy increases as we move
along a period from left to right due to
increased nuclear charge.
Atomic size: Ionization enthalpy decreases as
the atomic size increases.
Ionisation enthalpy increases as the penetration
effect of the electrons increases.
S>P>d>f If the penetration effect of the
electron is more, it is closer to the nucleus
hence IE is high.

56.  As the shielding or screening effect of the
inner electrons increases, the ionization
enthalpy decreases.
 Due to repulsion the valence shell electrons
do not feel the full charge of the nucleus. The
actual charge felt by the valence shell
electrons is called effective nuclear charge.

57.  Be to B:The ionization enthalpy of B is lower
than that of Be. This is due to:
 1)The outermost electron in B is present in 2p
orbitals while in Be it is present in 2s orbital.
Since 2s electrons are more penetrating
towards the nucleus than 2p electrons
,therefore, lesser amount of energy is
required to Knock out a 2p electron than 2s
electron. The first ionization enthalpy of
Boron is lower than that of Be.
 2)The 2p electron of B is not strongly
attracted by the nucleus as the 2s electron of
beryllium. Consequently the first ionization
enthalpy of Boron is lower than that of
beryllium.

58.  N to O :The first ionization enthalpy of oxygen is
lower than that of Nitrogen although the nuclear
charge of oxygen is higher than that of Nitrogen.
 1)The electronic configuration of N in which the
2p- orbitals are exactly half filled is more stable
than the electronic configuration of oxygen in
which the 2p orbitals are neither half and nor
completely filled.
 It is difficult to remove an electron from N than
from O. The first ionization enthalpy of nitrogen is
higher than that of oxygen.

59. 2)The removal of an electron from oxygen gives a
stable electronic configuration with exactly half
filled 2p subshell while this is not so in case of N.
The removal of an electron from oxygen give the
most stable electronic configuration than that
obtainable from nitrogen.
Along a period:
 As we move from left to right in a period, the
ionisation enthalpy increases with increasing
atomic number.
 IE keep on decreasing as we move down a group
from one element to the other.

60.  Electron gain enthalpy of an element may be
defined as the energy released when a neutral
isolated gaseous atom accepts an extra electron
to form the gaseous negative Ion i.e. anion. It is
denoted by ΔHeg .
 Electron gain enthalpy is measured in electron
volts per atom or kJ per mole.
 The process of adding an electron to the atom
can be either exothermic or endothermic.

61.  Positive electron gain enthalpy: Metals Ex: Na
and Mg
 Negative electron gain enthalpy: Non metals
Ex: S and Cl
 Factors on which the electron gain enthalpy
depends:
 (i) Atomic Size
 (ii) Nuclear charge
 (iii) Electronic configuration

62. Atomic Size
 As the size of the atom increases, the
distance between the nucleus and the last
shell which receives the incoming electrons
increases. As a result, the force of attraction
between the nucleus and the incoming
electron decreases and hence the electron
gain enthalpy becomes less negative.

63.  Nuclear Charge
 As the nuclear charge increases ,the force of
attraction between the nucleus and the
incoming electron increases and hence the
enthalpy becomes more negative.
 Electronic Configuration
 Elements having exactly half filled or
completely filled orbitals are very stable.

64.  Energy has to be supplied to add an electron.
Hence their electron gain enthalpy have large
positive values since they do not accept the
additional electrons so easily.
 The electron gain enthalpy becomes less
negative in going from top to bottom in a
group and more negative in going from left
to right in a period.

65.  The electron gain enthalpy of fluorine is less
negative than that of chlorine.
 This is due to its small size. As a result of its
small size ,the electron electron repulsion in
the relatively compact to 2p subshell are
comparatively large and hence the incoming
electron is not accepted with the same ease
as is the case with chlorine.

66.  The electron gain enthalpy of noble gases is
positive
 The atoms of these elements have completely
filled subshell. As a result there is no room in
their valence orbitals and the additional
electron has to be placed in an orbital of next
higher shell. As a result, energy has to be
supplied to add on additional electrons.
 The electron gain enthalpy increases from left
to right.
 The electron gain enthalpy decreases from
top to bottom

68.  Electronegativity of an element is the tendency
of its atoms to attract the shared pair of
electrons towards itself in a covalent bond.

69.  It is not a measurable quantity
 There is numerical scales of electronegativity
 Pauling scale, Mulliken-Jaffe scale, Alfred-
Rochow scale have been developed.
 Pauling scale is used to measure
electronegativity

70.  The electronegativity of any given element is
not constant, it varies depending on the
element to which it is bound.
 Electronegativity generally increases across a
period from left to right in the periodic table.
 Decreases down the group

72. Metallic and non metallic character
 As the electronegativity increases, the non
metallic character increases. As the
electronegativity decreases, the metallic
character increases.
Polar and non-polar bonds
 A bond between two similar atoms is said to
be known non-polar since the shared pair of
electrons is equally attracted by the two
atoms as the electronegativity of the atoms is
the same.
For Ex: H2 , Cl2, O2, N2

73.  The more electronegative atom acquires a
partial negative charge and the less
electronegative atoms acquire a partial
positive charge. As a result ,two poles are
developed and the molecule is said to
be Polar.
 For Ex: H-F , HCl, HBr, HI

74.  Periodicity of valence or oxidation states:
The electrons present in the outermost shell of
an atom are called valence electrons and the
number of these electrons determine the valence
or the valency of the atom.
The valence of an atom equal to either the
number of valence electrons are equal to 8 minus
the number of valence electron.

75.  Consider the two oxygen containing
compounds: OF2 and Na2O.
 The order of electronegativity of the three
elements involved in these compounds is
F>O>Na.
 Fluorine shares one electron with oxygen in
the OF2 molecule.
 Fluorine shows oxidation state -1.
 Oxygen shares two electrons with fluorine
atoms and exhibits oxidation state +2.

76.  In Na2O, oxygen has more electronegative
accepts two electrons one from each of the
two sodium atoms and shows -2 oxidation
state.
 On the other hands sodium loses one
electron to oxygen and is given oxidation
state is +1.

78.  Anomalous properties of second period
elements
 The first element of group 1 and 2 and
group 13 to 17 differ in many respect from
the other members of their respective group.
 Some elements on the 2nd period shows
similarities with the elements of the 3rd
period present diagonally to each other,
though belonging to different group.
 Similarities in properties of the elements
placed diagonally to each other is
called diagonal relationship.

80.  The reason for the different chemical behaviour
of the first member of a particular group of
elements in the s and p block compared to the
other members of the same group are
 1)small size
 2)large charge/ radius ratio
 3)high electronegativity
 4)absence of d orbital
 5)the maximum covalency of first member of
each group is 4 whereas other members of the
groups can have a maximum covalency of 6.
 6)Because of smaller size and high
electronegativity ,first member of each group
show greater ability to form pπ-pπ multiple
bonds either with itself or with other members
of the second period.
• C=C, CC, N=N

81.  The total number of orbitals available in
valence shell is known as covalency.
 Periodic Trends and chemical reactivity
 The atomic and ionic radii decrease in period
from left to right.
 Ionization enthalpies increases across a
period.
 Electron gain enthalpies become more
negative across a period

82.  The chemical reactivity of alkali metals on the
extreme left is due to their ability to lose an
electron to form corresponding cation.
 Chemical reactivity of halogens on the
extreme right is due to their ability to gain an
extra electron leading to the formation of the
anion.
 Alkali metals are good reducing agents while
halogens are good oxidizing agents.

83.  Metallic character of an element which is highest
at the extreme left decreases while the non
metallic character increases on moving across the
period from left to right.
 The elements at the extreme left of the periodic
table readily combine with oxygen to form oxides
which are most basic.
 Elements on extreme right from oxides which are
most acidic.
 Oxides of the elements in the centre are either
amphoteric or neutral.
 The amphoteric oxides show both acidic and basic
properties .
 Neutral oxides have neither acidic not basic
properties.