Periodic classification of elements Grade-11

Topic: Periodic Classification of Elements
Class : XI
By
DP Paneru
M.Sc Chemistry
d.ppaneru1991@gmail.com

Atomic
Number
Element Symbol Atomic Mass
1 Hydrogen H 1.008
2 Helium He 4.0026
3 Lithium Li 6.94
4 Beryllium Be 9.0122
5 Boron B 10.81
6 Carbon C 12.011
7 Nitrogen N 14.007
8 Oxygen O 15.999
9 Fluorine F 18.998
10 Neon Ne 20.180
11 Sodium Na 22.990
12 Magnesium Mg 24.305
13 Aluminium Al 26.982
14 Silicon Si 28.085
15 Phosphorus P 30.974
Atomic
Number
Element Symbol Symbol
16 Sulfur S 32.06
17 Chlorine Cl 35.45
18 Argon Ar 39.948
19 Potassium K 39.098
20 Calcium Ca 40.078
21 Scandium Sc 44.956
22 Titanium Ti 47.867
23 Vanadium V 50.942
24 Chromium Cr 51.996
25 Manganese Mn 54.938
26 Iron Fe 55.845
27 Cobalt Co 58.933
28 Nickel Ni 58.693
29 Copper Cu 63.546
30 Zinc Zn 65.38

Content Area: General and Physical Chemistry
4. Classification of elements and Periodic Table
Modern periodic law and modern periodic table
Classification of elements into different groups, periods and blocks
IUPAC classification of elements
Nuclear charge and effective nuclear charge
Periodic trend and periodicity
Atomic radii
Ionic radii
Ionization energy
Electron affinity
Electronegativity
Metallic characters (General trend and explanation only)

Why elements are classified?
The arranging of elements into different groups on the
basis of the similarities in their properties is called
classification of elements.
The classification of similar elements into groups
makes the study of elements easier.

The elements are first classified
by the French scientist Antoine
lavoisier in 1789. The earliest
attempt to classify elements was
grouping the known elements
(about 30 elements) into two
groups called metals and non
metals.
The defect in this classification
was that it had no place for
metalloids.
Antoine Lavoisier

A German chemist Dobereiner (1817)
classified elements in the increasing
order of their atomic masses into three
groups called triads.
In each triad the atomic mass of the
middle element was approximately
equal to the average atomic mass of the
other two elements.
The defect in this classification was
that all the known elements could not
be correctly arranged into triads.Wolfgang Dobereiner

Newland (1864) classified the elements (56) in
the increasing order of their atomic masses into
groups of eight elements called octaves.
DEFECTS
 All the known elements and elements discovered later
could not be correctly arranged into octaves.
Some elements having different properties were placed
in the same rows like cobalt and nickel having different
properties were placed along with Fluorine, Chlorine and
Bromine.Newland octaves

In order to remove demerits of
previous attempt of classification of
elements, A Russian Chemist,
Dmitri Mendeleev in 1869
classified elements (63) on the basis
of increasing order of atomic mass
.
It state that the “physical and
chemical properties of elements
are periodic function of their
atomic mass or atomic weight”
Dmitri Mendeleev

Feature of Mendeleev’s
periodic table
It consists of 6 horizontal
rows called periods.
It consists of 8 vertical
rows called groups.
The groups 1 to 7 had two
sub groups called A sub
group and B sub group.
 Group 8 had 3 rows of
elements.
Elements having similar
properties were placed in the
same groups.
There are some spaces left
vacant in the table to
accommodate the elements to
be discovered in future.
There was separate column
left for inert gases.

Periodic classification of elements Grade-11

i) Elements were classified on a more fundamental basis of
their atomic masses and properties.
ii) Spaces were left vacant to accommodate the elements to be
discovered in future .
iii) It could predict the properties of the elements which helped
in the discovery of new elements.
iv) The inert gas elements discovered later could be placed in a
separate group without disturbing the table.
V) Correction of faulty atomic weights
Mendeleev corrected atomic weight of Be 9 from 13.5
Advantage of Mendeleev’s periodic Table

Demerits of Mendeleev's periodic table
1. Position of hydrogen:
The position of hydrogen in his periodic table is not properly defined. Sometimes
hydrogen loses electron like Group I elements and sometimes it gains electrons
like those of group VII. Due to this property of hydrogen, it can be placed
either in group I or group VII.
2. Separation of chemically similar elements:
Some chemically similar elements have been placed in different groups such as
copper and mercury, gold and platinum, silver and thallium, barium and lead.
3. Grouping of chemically dissimilar elements:
Certain chemically dissimilar elements have been placed in the same group. For
example: Copper, silver and gold have been placed in group I along with highly
electropositive metals like lithium, sodium and potassium.
4. Position of isotopes:
Isotopes have the same atomic number, similar chemical properties but different
atomic masses. Based on the periodic law, an isotope is an element, which must
have a separate place in the periodic table. However, there are no places for
isotopes in his periodic table.

Demerits of Mendeleev's periodic table (cont..)
5. Anomalous pair of element:
Certain elements having higher atomic weights have been placed wrongly before
those with less atomic weight. For example, Argon (Ar) having higher atomic
weight(39.9) has been placed before Potassium (K) with atomic weight (39).
6. Position of lanthanides and actinides:
The fourteen elements from atomic number 58 (cerium) to 71 (lutetium) are termed
as lanthanides (rare earth elements).
The elements from atomic number 90(thorium) to 103 (lawrencium) are termed as
actinides (trans-uranium elements).
These elements have similar properties with that of actinium. Mendeleev could not
arrange their property in the fixed position.
7. Position of iron and platinum group:
Mendeleev’s did not justify further division of group I to VII into two sub group A &
B and three division of group VIII.

In order to remove demerits of
previously discovered Mendeleev’s
classification of elements A British
physicist, Henery Moseley in 1913
classified elements on the basis of
increasing order of atomic Number
.
It state that the “physical and
chemical properties of elements
are periodic function of their
atomic number”
Henry Gwyn Jeffreys Moseley

Feature of Modern Periodic
table
1.The elements are arranged
based on their increasing
atomic number.
2. There are 7 periods and 18
groups in the table.
3. The inert gases i.e. He, Ne,
Ar, etc. are kept in the 0
group at the extreme right
side of the
table.
4. Lanthanides and actinides
are kept below separate from
main body of table.

Component of Modern periodic table
1. Periods:
Modern periodic table contains horizontal rows called periods. Each period
starts with an alkali metal and ends with an inert gas element. Elements
present in the same period have same number of shells.
It contains 7 periods and are represented by n=one, n=two and so on.
Periods of modern periodic table are of 3 types:
i. Short period:
1st period contains only two elements whereas 2nd and 3rd period contains
elements each. So, these periods are called short period.
ii. Long period:
4th and 5th period contain 18 elements each and 6th period contains 32
elements. So, they are called long period.
iii. Incomplete Long Period:
7th period contains 30 elements and some gap left for the elements, which
will be discovered in the future. So, it is called incomplete long period.

GROUPS
The Modern Periodic Table consists of 18 groups or vertical columns.
Elements present in the same group show same physical and chemical properties.
 Also the elements present in the same group have same number of electrons in the
outermost shell.
According to American convention, the groups are denoted by roman numerals
followed by either an capital alphabet "A" if the group is in the s-block or p-block, or
by "B" if the group is in the d-block .

Groups:
1. Group IA elements:
 Elements like hydrogen, lithium, sodium and potassium etc.
 They can form electropositive cations after losing their one valence electrons.
 Elements of group IA are called alkali metals. As their oxides are strongly basic.
2. Group II A:
In group II A, elements like beryllium, magnesium, calcium, etc.
Their valency is +2 and they can form bivalent cations.
 They are called alkaline earth metals because they form basic hydroxides which
are less soluble in water.
3. Group III A elements:
B, Al, Ga, In, TI etc
Boron family
Valency +3
Non-metal
4. Group IVA elements:
C, Si, Ge, Sn, Pb etc
Carbon family.
Valency ±4
5.Group VA elements:
N,P, As, Sb, Bi etc
Nitrogen family.
Valency +5/-3
7.Group VIIAelements:
F, Cl, Br, I etc
Halogen family.
Valency +7/-1
6.Group VIA elements:
O, S, Se, Te, Po etc
Oxygen family.
Valency +6/-2
5.Group VIIIA elements:
He, Ne, Ar, Kr, Xe etc
Zero group/ Inert gas.
Valency zero.

8. Transitional element:
Those elements lying in between group II A and IIIA (i.e.fromgroup I B to
group VII B and Group VIII) are called transition elements.
They are metal.
Electronic configuration is (n-1)d1-10 ns1-2
9. Inner-transitional element:
The fourteen elements from atomic number 58 (cerium) to 71 (lutetium) are
termed as lanthanides (rare earth elements).
 The elements from atomic number 90(thorium) to 103 (lawrencium) are
termed as actinides (trans-uranium elements).
 F-block elements.
 Electronic configuration is (n-2)f1-14 (n-1)d1-10 ns1-2

The following are the advantages of Modern Periodic table.
1. This is scientific classification because the atomic number is more
fundamental properties of an element than atomic mass.
2. In this law elements are further classify into four major electronic block s,
p, d, and f which further helps to study the properties of elements more
systematically.
3. It fixed the position of hydrogen in IA group because hydrogen is s-block
element.
4. It remove the anomalous pair of element eg. Ar-K, Co-Ni, Te-I by
arranging elements on basis of atomic number.
5. Since the elements are arranged in increasing atomic number, the positions
of isotopes is justified.
6. Having similar properties of lanthanide and actinides are excluded from
main body which is easier to study their properties.
7. It gives explanation for the periodicity of elements and tells the reason
why all elements in a group having similar properties, which differ
between the groups and periods.
8. In the Modern periodic table, the metals, non-metals, transition metals,
gases are separately placed in a specific location with a specific identity.

 Group VII consists of three column without any
justification.
 The location of helium amongst the p-block element is not
fully justified as its electronic configuration, it to be in s-
block.
 The exclusion of lanthanides and actinides from main body
of table is still remains to be solved.

Division of elements in s,p,d and f sub-shell

s-Block elements:
 In this block elements, last electron enters the ns orbital.
The s-block elements lie on the extreme left of the periodic table.
The elements of groups IA, IIA belong to this block.
Valence shell electronic configuration of these elements varies from ns1 to ns2.
Valence shell electronic configuration group IA elements is ns1 and group IIA
is ns2.
S-block elements are electropositive in nature.
They have +1/ +2 valency/oxidation state.
 They form basic oxides and their hydroxides are basic in nature.
They impart characteristic colour to the flame.
They have low ionization potentials.
They are solids at room temperature (Cs is liquid at about at 35oC)
They are good reducing agents.

p-Block elements:
 In this block elements, last electron enters the p-orbital.
The p-block elements lie on the extreme right of the periodic table.
The elements of are belong to this block IIIA, IVA,VA,VIA.VIIA & VIIIA.
Valence shell electronic configuration of these elements varies from ns2 nP1-6.
p-block elements are some non-metals, metalloids and inert gases.
The valency of these block elements is varies from +3,±4,-3,-2,-1 & Zero.
 They have variable oxidation states.
They form acidic oxides.
Most of p-block elements are electronegative in nature.
Generally they form covalent compounds. Halogens form salts with alkali
metals
They have high ionization potentials.
They are solids/liquids/gases at room temperature (Br is liquid).

d-Block elements:
 In this block elements, last electron enters the d-orbital.
The d-block elements lie between s-bock and p-block in the periodic table.
The elements of are belong to this block IB, IIB, IIIB,IVB,VB,VIB,VIIB,VIIIB.
Valence shell electronic configuration of these elements varies from (n-1)d1-10ns
d-block elements are transition metals.
The valency of these block elements is varies.
Elements as they show transitional properties between s and p-block elements.
Many of their compounds are colored.
They readily form complexes by acting as Lewis acid.
Most of them and their compounds show paramagnetic, ferromagnetic
behavior.
Most of them act as good catalysts.

f-Block elements:
 In this block elements, last electron enters the f-orbital.
The f-block elements lie out side below main body of table.
The elements of are belong to this block lanthanides and actinides.
Valence shell electronic configuration of these elements varies from
(n-2)f1-14 (n-1)d1-10ns1-2 .
f-block elements are inner-transition metals. (The valence electrons of these elements
transitions into the (n-2) f block which is the "anti-penultimate energy level" and lies in the f-orbital.
Hence these are called as 'inner transition elements'.)
There are two series in the f block i,e, 4f and 5f.
Properties are similar to d-block elements.
 Most lanthanides are widely used in lasers
 All the actinides are radioactive.
These are highly electro-positive (show+3,+4,+5,+6 oxidation states)
These metals tarnish in air.
They have number of isotopes.

Aufbau Principle
According to this principle: “Filling of electrons always occurs from lower energy
level to higher energy level. The orbital with lower energy level are filled up first.
Only after then the orbitals with higher energy are filled. By following this rule, we can
predict the electron configurations for atoms or ions.

So the electronic configuration of first 20
elements becomes:
Hydrogen: 1s1
Helium: 1s2
Lithium: 1s2, 2s1
Beryllium: 1s2, 2s2
Boron: 1s2, 2s2, 2p1
Carbon: 1s2, 2s2, 2p2
Nitrogen: 1s2, 2s2, 2p3
Oxygen: 1s2, 2s2, 2p4
Fluorine: 1s2, 2s2, 2p5
Neon: 1s2, 2s2, 2p6
Sodium: 1s2, 2s2, 2p6, 3s1
Magnesium: 1s2, 2s2, 2p6, 3s2
Aluminum: 1s2, 2s2, 2p6, 3s2, 3p1
Silicon: 1s2, 2s2, 2p6, 3s2, 3p2
Phosphorus: 1s2, 2s2, 2p6, 3s2, 3p3
Sulphur: 1s2, 2s2, 2p6, 3s2, 3p4
Chlorine: 1s2, 2s2, 2p6, 3s2, 3p5
Argon: 1s2, 2s2, 2p6, 3s2, 3p6
Potassium: 1s2, 2s2, 2p6, 3s2, 3p6, 4s1
Calcium: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2
Here are some exception of Aufbau’s principle. This principle does not hold true for
chromium, copper.
For chromium:
Expected electronic configuration: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d4
Actual electronic configuration: 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5
For copper:
Expected electronic configuration: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d9
Actual electronic configuration: 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d10
Note: Half filled and full filled orbitals are more stable because:
Half filled orbitals have symmetry among the orbitals.
There is a exchange of energy between the electrons.

What is meant by periodicity? Explain with example.
Ans: The recurrence of elements with similar properties after certain regular
intervals when these elements are arranged in the increasing order of their
atomic numbers is called periodicity.
3Li = 1S2,2S1
11Na = 1S2, 2S22P6, 3S1
19K = 1S2, 2S22P6, 3S23P6, 4S1
37Rb = 1S2, 2S22P6, 3S23P63d10, 4S24P6, 5S1
55Cs = 1S2, 2S22P6, 3S23P63d10, 4S24P64d10, 5S25P6,6S1
All these elements have one electron in their valence shell and have the similar
properties.
From the electronic configuration of these elements it is clear that all elements
having same number of electrons in the valence shell have similar properties.
Thus, the recurrence of similar properties is due to the recurrence of similar
electronic configurations or atomic structure. Hence, Cause of periodicity is the
recurrence of similar electronic configuration.

What are iso-electronic species with examples.
Ans: Neutral species or ions of different elements having same number
of electrons but different magnitude of nuclear charge are called iso-
electronic species or ions. Size of such ions depends upon nuclear
charge.
Example: Sulphide (S2 -), chloride(Cl-) and potassium (K+) ions are iso-
electronic ions because each has 18 electrons but have different nuclear
charges i.e. +16, +17 and +19 respectively.
Why lanthanides and actinides are placed in separate rows at the
bottom of the periodic table?
Ans: Lanthanides and actinides are placed in separate rows at the
bottom of the periodic table to avoid unnecessarily sidewise expansion
of the periodic table. As, In both series of elements, the outermost shell
and penultimate shell ( the second outermost shell of the atom) are
incompletely filled but the filling of only f - orbitals of ante-penultimate
shell (one inner to penultimate shell ) occurs.

Periodic Properties
1. Atomic size/ radius
2. Ionization energy (Ionization potential) (I.E or
I.P)
3. Electron affinity (E.A)
4. Electronegativity (E.N)

Atomic size
It refers to the distance between the centre of the nucleus of an atom to its
outermost shell of electrons.
Since the absolute value of atomic size cannot be determined, it is
expressed in terms of the operational definitions such as ionic radius,
covalent radius, Vanderwaal's radius, and metallic radius. The absolute
value of atomic radius can not be determined because-
(i) It is not possible to locate the exact position of electrons in an atom as
an orbital has no sharp boundaries.
(ii) It is not possible to isolate an individual atom for its size
determination.
(iii) In a group of atoms, the probability distribution of electrons is
influenced by the presence of neighbouring atoms.
Thus, the size of an atom may change in going from one environment to
other.

(a) Covalent radius:
It is defined as half of the distance between two successive nuclei of two
covalent bonded an atom in a molecule.
Covalent radius
If the bond length between two atoms is d, then
covalent radius = 1/2 * [internuclear distance in two covalent bonded atoms]

(b) Vander Waals' radius:
It is defined as one-half of the distance between the nuclei of two non-
bonded isolated atoms or two adjacent atoms in the solid state.
In general, Vanderwaals' radius is greater than the covalent radius of an
atom.

c) Metallic radius:
It is defined as half the distance between two successive nuclei of two
adjacent metal atoms in the metallic-closed packed crystal lattice.
(d) Ionic radius:
It is the distance of the outermost shell of an anion or cation from its
nucleus.
An atom can be changed into a cation (loss of an electron) and to an
anion (by gaining of electrons).
Note:
The size of the cation is always less than that of a parent atom (due
to increased nuclear charge).
whereas the size of the anion is always greater than the parent atom
(due to decreased effective nuclear charge).

(e) Effective nuclear charge:
It is defined as the net nuclear attraction towards the valence shell
electron or in other words, the actual nuclear charge, where the
electrostatic force of attraction is being experienced by the outer electron.
Greater the effective nuclear charge, more tightly is the hold with the
nucleus. For example - In Na –
Another example of atomic and ionic radius
Generally, when an element (non-metal) gains an electron, its
ionic radius increases due to decreased effective nuclear charge.

Variation of Atomic Radii in Periodic Table
Variation in a period
Generally, when we move from left to right in the same period, the atomic
radii decreases due to increased effective nuclear charge and thus, the
electron cloud is attached more strongly towards the nucleus which gives rise
to a contraction in atomic radius in the period.
Variation in a group
Generally, when we move from top to bottom in the same group, the atomic
radii of the element increases because, as we go down from top to bottom of
the group the nuclear charge increases, at the same time an extra shell is
added in successive members. Here effect of increase of nuclear charge is
dominated by effect of increase shell. As a result, the size of the atom
increase top to bottom in the same group.
Iso-electronic ions or species:
They are ions of the different elements which have the same number of
electrons but the different magnitude of the nuclear charge. The size of
isoelectronic ions decreases with the increase in the nuclear charge.

Q. A positive ion is smaller but negative ion is larger than their
parent atom, explain.
Ans: An anion is formed by the gain of one or more electrons by the neutral
atom. In the anion the nuclear charge is same as in the parent atom but
the number of electrons has increased. Since same nuclear charge acts on
increased number of electrons, the effective nuclear charge per electron
decreases in the anion. The electron cloud is held less tightly by the
nucleus. This cause increase in size of anion.
Q. Which of the following pairs would have larger size and why?
i. K or K+ ii. F or F - iii. Fe2+ and Fe3+
 K > K+
 F < F -
 Fe2+ > Fe3+

Q. Arrange the following in order of increasing size
a. I, I+, I - b. Na, Mg, K c. Cl -, S 2-, Mg 2+,Br-, Na +
Ans: a. I+< I< I -
b. Na< Mg< K
c. Mg 2+ <Na +< Br-<Cl -< S 2-
Q. Based on their positions in the periodic table, list the following ions in order of
increasing radius: K+, Ca2+, Al3+, Si4+.
Q. Which of the following atoms and ions is (are) isoelectronic with S2+: Si4+, Cl3+,
Ar, Si, Al3+?

Ionization energy (Ionization potential) (I.E or I.P)
It is the amount of energy required to remove the loosely bound electron
from the isolated gaseous atom to change it into a gaseous cation. It is
denoted by I.P or I.E and its unit is eV.
A(g) + I.E1 → A+ + e-
A+ + I.E2 → A++ + e-
A++ + I.E3 → A+++ + e-
The amounts of energies required to remove 2nd, 3rd ,4th… electron from
cations after removal of 1st electron are successive ionization energies.
The ionization energy order is
I.E1<I.E2<I.E3<I.E4…
In general successive ionization energies always increase as the electrons to
be removed get closer to the nucleus, there is a strong attraction between
nucleus and electron, therefore more energy is required to remove the next
electron.

Factors responsible for ionization energies:
(i) Atomic size or radius:
Ionization energy is inversely proportional to the atomic size of the element i.e.
higher the atomic size of the element, lower will be its ionization energy and
vice versa. As the distance of the outer shell electron increases from the atomic
nucleus, the attractive force decreases and the outer electrons are loosely held.
(ii) Shielding effect or number of intervening electrons or screening:
The phenomenon in which the penultimate shell (n-1) electrons act as a screen or
a shield between the nucleus and the valence cell electrons, thus reducing the
effective nuclear charge is known as the shielding effect. The penultimate shell
electrons repel the valence shell electrons to keep them loosely held with the
nucleus. Evidently, greater the shielding effect, lesser is the effective nuclear
charge and lesser is the ionization effect.

(iii) Nuclear charge:
The greater the effective nuclear charge, the more strongly the electrons are held
on by the nucleus. More energy is required to remove the electron and
consequently higher is the ionization energy and vice versa.
(iv) Penetration effect:
The sub-shells of a shell are more closer (penetrating) to the nucleus, more tightly
the electrons are held toward the nucleus and more is the ionization energy.
Ionization energy : s > p > d > f for a given shell.
(v) Electronic configuration
Stable the configuration of an atom, it is Difficult to lose an electron from its valence
shell and hence higher will be the value of ionization energy.
Stability order configuration is
Completely filled shell (K.L.M.N...) > Completely filled sub-shell(s2,p6,d10,f14)>
Half filled sub-shell (s1, p6, d5,f7) > Partially filled sub-shell

Variation of ionization energy:
In Period
IE of an element generally increase from left to right in the same period
and become maximum in a zero group because of gradually decrease in size
and increase in nuclear charge from left to right.
Although, ionization energy increases from left to right across a period,
variation along the period is not quite linear.
In Group
In general, ionization energy decreases from top to bottom down a group.
This is due to increasing atomic radii is much more pronounced than that of
nuclear charge and electron experience less of nuclear attraction. Thus IP
decreasing in group on moving top to bottom.

Variation of ionization energies in period and groups in Periodic table

H
Li
Be
B
C
N
O
F
Na
Mg
Al
Si
P
S
Cl
Variation of ionization energy

"The ionization energy of Na+ is more than that of Ne although
both have same electronic configuration ", explain why?
Ans: Although Na+ and Ne have same electronic configuration (iso-electronic
species) but the nuclear charge in Na+ is more than in Ne. Hence, the
electrons are more tightly held in Na+ and it has higher ionization energy.
A, B and C represents elements of atomic number 2, 6, 19
respectively.
a. Write their electronic structure
Ans: A= 2He => 1S2
B = 6C = 1S2, 2S22P2
C = 19K => 1S2, 2S22P6, 3S23P6, 4S1
b. Predict their blocks, period and group
Ans: A= p-block, 0-group and 1st period
B = p-block, IVA-group and 2nd period
C = s-block, IA-group and 4th period
c. Which one is a representative element?
Ans: C = 19K is a representative element.

The first ionization of noble gases is higher than that of halogen.
Explain.
Ans: The noble gas have stable electronic configuration so it is very difficult
to remove the electron from their outermost orbit. Thus, they have highest
ionization energy. Whereas halogens have unstable electronic configuration
as compared to noble gas and it is easy to remove the outermost electrons.
Hence, the first ionization energy of noble gas is greater than that of halogen.

1. Compare the size of F- and Na+ with the atomic size of Ne.
2. Write down the electronic configuration of Cu2+ and Ar.
3. Mention the necessity for the classification of the elements.
4. State the Mendeleev's/Modern periodic law.
5. What is the basis of classification of elements in the long form of periodic table?
6. What is meant by periodicity? Explain with example.
7. What are iso-electronic species with examples.
8. Why lanthanides and actinides are placed in separate rows at the bottom of the
periodic table?
9. Why lanthanides and actinides are placed in separate rows at the bottom of the
periodic table?
10. A positive ion is smaller but negative ion is larger than their parent atom, explain.
11. What are the factors which influences the ionization energy?
12. "The ionization energy of Na+ is more than that of Ne although both have same
electronic configuration ", explain why?

1. Among F, Cl, Br and I which has the smallest ionization potential
2. The first ionization of noble gases is higher than that of halogen. Explain.
3. How does atomic size vary within a horizontal row of the periodic table?
4. Show your acquaintance with modern periodic table.
7. Compare the size of F- and Na+ with the atomic size of Neon.
8. Name the four blocks of elements in the periodic table. Write their general
electronic configuration; mention at least four characteristics of each block.
9. In what respect does Modern periodic table differ from Mendeleev's periodic
table. How does it help in removing the defects of Mendeleev's periodic table?
10. In what respect does Modern periodic table differ from Mendeleev's periodic
table. How does it help in removing the defects of Mendeleev's periodic table?

Electron affinity (E.A)
It is the amount of energy released when a gaseous atom accepts the
electron to form a gaseous anion. This is an exothermic process for all
non-noble gas elements. It is denoted by E.A. and its unit is kj/mole.
A(g) + e → A- + E.A
Variation of electron affinity:
In Period
Although, electron affinity increases from left to right across a period,
variation along the period is not quite linear.
This is due to increasing effective nuclear charge, which more readily pulls
new electrons inward.
In Period
In general, electron affinity decreases from top to bottom down a group.
This is due to a increasing atomic radii and decreasing electron-electron repulsions
and thus experience less of an electron-nucleus attraction.

Factors Affecting Electron Affinity
The various factors which influence the electron affinity can be explained under
following heads:-
Electron affinity is inversely proportional to the atomic size of the element i.e.
higher the atomic size of the element, lower will be its electron affinity and vice
versa. As the distance of the outer shell electron increases from the atomic nucleus,
the attractive force decreases and the tendency to accept extra electron will
decrease.
Evidently, greater the shielding effect, lesser is the effective nuclear charge and lesser is
the electron affinity.

(iii) Nuclear charge:
Greater the nuclear charge, greater will be the attraction for the incoming electron
and as a result larger will be the value of electron affinity.
(iv) Electronic configuration
Stable the configuration of an atom, lesser will be its tendency to accept an electron
and hence lower the value of its electron affinity.
Stability order configuration is
Completely filled shell (K.L.M.N...) > Completely filled sub-shell (s2,p6,d10,f14) >
Half filled sub-shell (s1, p6, d5,f7) > Partially filled sub-shell

What are some exceptions to these general trends?
Notice that the Group IIA elements have much lower electron affinities than the
Group IA elements. This is because Group IIA elements already have a full
valence s shell and the addition of a new electron would move to the higher
energy p shell. This occupation of a new, higher-energy shell is very energetically
unfavourable and so this accounts for lower, electron affinity.
Notice also the same trend in the as compared to the more negative electron
affinities of. This is because the Group VA elements stable configuration. Half
field p-orbital is more stable than partially field. So, less tendency to accept
extra electron. Hence Group IVA elements feel less electron affinity.
Notice that the Period 3 elements actually have the greatest electron affinities,
instead of the Period 2 elements as the trend suggests. This is because the
atomic radii of the Period 2 elements are considerably smaller, and thus the
electrons experience significant electron-electron repulsions that are not
completely off-set by the electron-nucleus attractions. For the Period 3
elements, their radii are larger and thus the balance between electron-nucleus
attractions and electron-electron repulsions is maximized.

Explain why?
Electron affinity of halogen is high.
EA of halogen is high. halogen have a greater electron affinity because of their
atomic structures: first, they have monoevalent electrons than others in the
respective period, thus it is easier for the halogen to gain electrons to fulfill a
stable octet and secondly, the valence electron shell is closer to the nucleus, thus it
is harder to remove an electron and it easierto attract electrons from other
elements. Thus halogen has high EA.
Electron affinity values of Be and Mg are almost zero.
EA values of Be and Mg are almost zero. It is because the outer electronic
configuration of Be is 1s2, 2s2. Here,' s'sub-shell is completely filled and is
quite stable. Likewise the outer electronicconfiguration of Mg is (Ne)3s2 , here
also the outer sub-shell is completely filledand it is also stable so in case of both
Be and Mg there is no addition of electron.So, the EA of Be and Mg are almost
zero.

Electronegativity (E.N)
It is the tendency of an atom in a molecule to attract the shared pair of electrons.
It is denoted by F.N. and its unit is
The most electronegative element in the periodic table is Flourine (4) and the least
electronegative element is Caesium (0.7) and inert gases have zero electro
negativity value.
Variation of electronegativity
In Period
Electronegativity increases from left to right across a period, variation along the
period.
This is due to increasing effective nuclear charge.
In Period
In general, electron affinity decreases from top to bottom down a group.
This is due to a increasing atomic radii and decreasing electron-electron repulsions and thus
experience less of an electron-nucleus attraction.
Fluorine has highest E.N.

Factors responsible for electronegativity
Electronegativity is inversely proportional to the atomic size of the element i.e.
higher the atomic size of the element, lower will be its electronegativity and vice
versa. As the distance of the outer shell electron increases from the atomic
nucleus, the attractive force decreases and the tendency to attract shared pair of
electron towards it self is decrease.
The phenomenon in which the penultimate shell (n-1) electrons act as a screen or a
shield between the nucleus and the valence cell electrons, thus reducing the effective
nuclear charge is known as the shielding effect. The tendency to attract shared pair of
electron towards it self is decrease.
(iii) Charge on an atom:
A cation attract the electron pair more readily towards itself than its parent atom. An
anion attracts the electrons pair less readily than its parent atom. This is due to size of
cation is smaller than its parent atom and size of anion is greater than its parent atom.

Which has larger electro negativity and why?
a. Na, O b. F, Cl
Ans: a. O < Na Electro negativity goes on increasing with increase in atomic
number as we move along the period.
b. F >Cl
Electro negativity deceases down the group as effective nuclear charge
decreases down the group.

What do you understand by diagonal relationship?
Ans: A diagonal relationship is said to exist between certain pairs of diagonally
adjacent elements in the second and third periods of the periodic table. These
pairs (Lithium (Li) and magnesium (Mg), beryllium (Be)
and aluminium (Al), boron (B) and silicon (Si) etc.) exhibit similar properties;
for example, boron and silicon are both semiconductors, forming halides that
are hydrolyzed in water and have acidic oxides.

Periodic classification of elements Grade-11

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Periodic classification of elements Grade-11