classification-of-elements-periodicity-in-properties-for-class-xi.ppt

CLASSIFICATION OF ELEMENTS AND
PERIODICITY IN PROPERTIES OF
ELEMENTS
BY-
A.P.S. BHADOURIYA
M.Sc. , B.Ed., NET
PGT-CHEMISTRY
K.V. BARABANKI

1. Why do we need
classification.?
2. Dobereniner’s triads
3. Newlands law of octave
4. Lother Meyer volume curve
5. Mendeleev’s periodic table
6. Modern periodic table
7. IUPAC nomenculature for
elements Z >100
8. Periodic Trends in Physical
Properties
Shielding effect & Effective
Nuclear Charge
Atomic Radius
Ionic Radius
Ionization Enthalpy
Electron Gain Enthalpy
Electronegativity
9. Periodic Trends in Chemical
Properties
 Periodicity of Valence or
Oxidation States
 Anomalous Properties of
Second Period Elements
 Chemical Reactivity
SESSIONOBJECTIVES

Lavoisier (1789) classified elements into
metals, non-metals, gases and earths.
During the nineteenth century, chemists
began to categorize the elements according
to similarities in their physical and chemical
properties. The end result of these studies
was our modern periodic table.

DOBEREINER’S TRIADS [ JOHN DOBEREINER (1817)]
In 1829, he classified some elements into groups of
three, which he called triads.
The elements in a triad had similar chemical
properties and orderly physical properties.
Model of triads
S.N
o
Triad
Atomic masses of
elements of triad
Arithmetic mean of
atomic masses of first
and third element
1 Cl,Br,I 35.5, 80, 127 35.5 + 127
2
= 81.25
3 Ca,Sr,Ba 40,87.5,137 40+137
2
= 88.5
2 Li,Na,K 7, 23, 39 7 + 39
2 = 23

In 1866, he suggested that elements be arranged in
“octaves” because he noticed (after arranging the
elements in order of increasing atomic mass) that
certain properties repeated every 8th element.
NEWLAND’S LAW OF OCTAVES [JOHN NEWLAND (1866)]

NEWLAND’S LAW OF OCTAVES [JOHN NEWLAND (1863)]
Element
Atomic mass
Element
Atomic mass
Element
Atomic mass
I II III IV V VI VII
Li Be B C N O F
7 9 11 12 14 16 19
Na Mg Al Si p S Cl
23 24 27 28 31 32 35.5
K Ca
39 40
Newland was first to publish the list of elements in
increasing order of atomic masses.

LOTHER-MEYER’S ATOMIC VOLUME CURVE
[LOTHER MEYER (1869)]

DMITRI MENDELEEV 1834 - 1907
In 1869 he published a table of the elements
organized by increasing atomic mass.

The physical and chemical
properties of elements are
periodic function of their atomic
masses.
MENDELEEV’S
PERIODIC LAW

Groups
•8 vertical rows.
•7 groups were subdivided in A and B.
•8th group has 9 elements in the group of
3 each.
Periods
•7 horizontal rows.
Only 63 elements were known.
MENDELEEV’S PERIODIC TABLE

MERITS OF MENDELEEV’S PERIODIC TABLE
Prediction of new
elements
(Ge, Ga, Sc)
1
Systematic study
of elements
2
Correction of
atomic mass
(Be, Au, Pt)
3

Mendeleev
•stated that if the atomic weight of an element
caused it to be placed in the wrong group, then
the weight must be wrong.
(He corrected the atomic masses of Be, In, and
U)
•was so confident in his table that he used it to
predict the physical properties of three elements
that were yet unknown.
After the discovery of these unknown elements
between 1874 and 1885, and the fact that
Mendeleev’s predictions for Sc, Ga, and Ge were
amazingly close to the actual values, his table
was generally accepted.

DEFECTS OF MENDELEEV’S PERIODIC TABLE
Position of
hydrogen.
Anomalous pairs.
(Ar and K, Co and
Ni, Te and I)
Position of
isotopes
e.g. 1H1,
1H2,
1H3

Chemically
dissimilar
elements are
grouped together.
(Cu-IA and Na-IB)
Chemically similar
elements are
placed in different
groups.
[Cu (I) and Hg (II)].
DEFECTS OF MENDELEEV’S PERIODIC TABLE

Mendeleev’s periodic table was published in 1905 when
no one had an idea of the structure of an atom.
DO YOU KNOW?
Mendeleev’s name has been immortalized by naming
the element with atomic number 101, as Mendelevium.
This name was proposed by American scientist Glenn
T. Seaborg, the discoverer of this element, “in
recognition of the pioneering role of the great Russian
Chemist who was the first to use the periodic system
of elements to predict the chemical properties of
undiscovered elements, a principle which has been the
key to the discovery of nearly all the transuranium
elements

English physicist, Henry Moseley observed regularities
in the characteristic X-ray spectra. A plot of f against
atomic number (Z ) of the elements gave a straight line
and not the plot of f vs atomic mass
He thereby showed that the atomic number is a more
fundamental property of an element than its atomic
mass.
MODERN PERIODIC LAW AND THE MODERN
PERIODIC TABLE

Mendeleev’s Periodic Law was,
therefore, accordingly modified.
This is known as the Modern Periodic
Law and can be stated as :
The physical and chemical properties of
the elements are periodic functions of
their atomic numbers.

HENRY MOSELEY
In 1913, through his work with X-rays, he determined the actual
nuclear charge (atomic number) of the elements*. He rearranged
the elements in order of increasing atomic number.
*“There is in the atom a fundamental
quantity which increases by regular
steps as we pass from each element to
the next. This quantity can only be the
charge on the central positive nucleus.”
His research was halted when the British government sent him to
serve as a foot soldier in WWI. He was killed in the fighting in
Gallipoli by a sniper’s bullet, at the age of 28. Because of this loss,
the British government later restricted its scientists to noncombatant
duties during WWII.

FEATURES OF LONG FORM OF PERIODIC
TABLE
•Contains elements arranged in
increasing order of atomic numbers.
•Explains the position of an element in
relation to other elements.
•Consists of groups and periods.

FEATURES OF LONG FORM OF PERIODIC TABLE
Groups Vertical column
Total 18. Numbered 1-18 or
IA to VII A, IB to VII B, VIII and zero.
Periods Horizontal column
Total 7 numbered from 1 to 7.
Elements in a group have similar but not
identical electronic configuration and properties
Contains 2,8,8,18,18,32 and 28 elements
respectively.

ELECTRONIC CONFIGURATIONS AND TYPES OF
ELEMENTS:
On the basis of the nature of sub-shell in which
last electron of atom enters, elements are
divided into 4 blocks
 s-Block Element
p-Block Element
d-Block Element
f- Block Element
s-,p-,d-,f- Block Elements

• Electronic configuration:
• Groups:
• All are metal, low ionisation energy and low
melting and boiling points, electropositive
elements.
• compounds are mostly ionic & colourless.
IA (alkali metals )and
IIA(alkaline earth metals
ns1 or ns2
In these elements last electron enters the s-orbital
s-Block Elements

• Groups:
• Non-metals, electronegative.
• Form covalent compounds.
ns2,np1 -6
III A to VII A and zero group (group 13-18).
In these elements last electron enters the p-orbital
p-Block Elements

• Groups:
• Variable valency high melting and boiling point.
• Coloured compounds and catalytic property.
(n-1)d1-10 ns1or2
I B to VII B and VIII groups (Gr- 3-12).
In these elements last electron enters the d-orbital,
Also known as transition metals.
d-Block Elements

• Have high melting and boiling point.
(n-2)f1-14(n-1)d0-1ns2
•Present below the periodic
table in two rows
•Lanthanides-elements after
lanthanum(Gr.-3, Pd.-6)
•Actinides-elements after
actinium. (Gr.-3, Pd.-7)
In these elements last electron enters the f-orbital,
Also known as Inner-Transition Elements
f-Block Elements

Representative elements
Transition elements
s and p block elements .
d-block elements. Valence shell and penultimate
Shell both are incomplete.
Inner Transition elements
f-block elements. Valence shell, penultimate shell
antipenultimate shell are incomplete.

Metals
•Present on left hand side of periodic
table.
•Solid,malleable,ductile and conductors .
Non-metals
•Present on right hand side of periodic
table.
•Solid or liquid or gas.
Metalloids
•Present on zig-zag between metals and non-metals.
e.g. B,Si,Ge,As,Sb and Te.

• Based on a more fundamental basis
- the atomic number
• Position of an element is related to the electronic
configuration of its atom.
• Due to separation of elements into groups, dissimilar
elements (e.g. alkali metals I A and coinage metals I B)
do not fall together.
MERITS OF LONG FORM OF PERIODIC TABLE

DEFECTS OF LONG FORM OF
PERIODIC TABLE
The problem of the position of hydrogen in
the table has not been solved completely
Configuration of Helium(1s2 ) is different
from inert gases (ns2,np6) but are placed in
the same group.
It is unable to include lanthanides and
actinides in its main body.

e.g. atomic number 115
Will be named as
un+un+pent+ium
=ununpentium
and symbol is Uup
Name
=digits name + ium
NOMENCLATURE OF THE ELEMENTS
WITH ATOMIC NUMBER >100
Digit Name Abbrevia
tion
0 nil n
1 un u
2 bi b
3 tri t
4 quad q
5 pent p
6 hex h
7 sept s
8 oct o
9 enn e

Periodic Properties
Periodic Trends in Physical Properties
Shielding effect & Effective Nuclear Charge
Atomic Radius
Ionic Radius
Ionization Enthalpy
Electron Gain Enthalpy
Electronegativity

Periodic Trends in Chemical Properties
 Periodicity of Valence or Oxidation States
 Anomalous Properties of Second Period
Elements
Periodic Properties

Shielding effect & Effective Nuclear Charge
The decrease in nuclear charge ( nuclear
force of attraction) on outermost shell
electrons due to repulsion caused by inner
shell electron is known as shielding effect of
inner shell or intervening electrons on outer
shell electron.

Due to shielding effect the nuclear charge is lowered
on outermost shell electrons, the net nuclear
charge acting on outermost shell electrons is known
as Effective Nuclear Charge. It is denoted by Z* or
Zeff.
 Z* or Zeff. = Z - σ
 where Z = nuclear charge( = atomic No.) &
 σ = shielding constant or screening constant , it is a
measure of shielding effect

Determination of ENC (Z*)
If the electron resides in s or p orbital
1. Electrons in principal shell higher than the e- in
question contribute 0 to σ .
2. Each electron in the same principal shell contribute
0.35 to σ (0.30 if it is 1S shell).
3. Electrons in (n-1) shell each contribute 0.85 to σ .
4. Eelectrons in deeper shell each contribute 1.00 to σ

If the electron resides in d or f orbital
1. All e-s in higher principal shell contribute 0 to σ
2. Each e- in same shell contribute 0.35 to σ
3. All inner shells in (n-1) and lower contribute
1.00 to σ

e.g. Calculate the Z* for the 2p electron Fluorine
(Z = 9) 1s2, 2s 2p5.
Soln. Screening constant for one of the outer electron
 6 (six) (two 2s e- and four 2p e-) = 6 X 0.35 = 2.10
 2 (two)1s e- = 2 X 0.85 = 1.70
 σ = 1.70+2.10 = 3.80
 Z* = 9 - 3.80 = 5.20

Trend of ENC in Periodic Table
In a Period - Effective nuclear charge Z*
increases increases rapidly along a
period(0.65 per next group)
e.g.
Li Be B C N O F Ne
1.3 1.95 2.6 3.3 3.9 4.6 5.2 5.9

Trend of ENC in Periodic Table
 In a Group - Effective nuclear charge Z* increases
slowly along a group.
e.g.
Gr-1 H Li Na K Rb Cs
Z* 1.0 1.3 2.2 2.2 2.2 2.2

PERIODIC TREND OF ATOMIC RADIUS
In A Period-
 atomic radius decreases with increase in atomic number
(in a period left to right)
BECAUSE in a period left to right-
 1. n (number of shells) remain constant.
 2. Z increases (by one unit)
 3. Z* increases (by 0.65 unit)
 4. Electrons are pulled close to the nucleus by the increased
Z*

 In a group-
Atomic radius increases moving down the group
 Because, along a group top to bottom
1. n increases
2. Z increases
3. No dramatic increase in Z* - almost remains
constant

IONIC RADII
 All anions are larger than their parent atoms.
because the addition of one or more electrons would result
in increased repulsion among the electrons and a decrease
in ENC.
 The cations are smaller than their parent atoms
because it has fewer electrons while its nuclear charge
remains the same & hence ENC is greater in cation than its
parent atom

ISOELECTRONIC SPECIES
 Atoms and ions which contain the same number of
electrons, are called as isoelectronic species.
For example, F–, Na+ and Mg2+ have the same number of
electrons(=10).
 The size of isoelectronic species decreases with increase in
nuclear charge. e.g.-
o2->F- >Ne>Na+>Mg2+>Al3+
---------SIZE DECREASING------

NOTE:
Metallic radii in the third row d-block are similar to
the second row d-block, but not larger as one
would expect given their larger number of
electrons.
This is due to Lanthanide Contraction as f-orbitals
have poor shielding properties.

Ionisation Energy (IE) or
Ionisation Enthalpy (ΔiH )
 Ionization: removing an electron from an atom or ion
 Ionization energy: energy required to remove an electron
from an isolated, gaseous atom or ion is called as Ionization
energy or ionisation enthalpy.
 If the atom is neutral the above defined ionisation energy
is called as first ionisation enthalpy.
 Energy required to remove an electron from an isolated,
monovalent cation is called as second Ionization energy.
 The ionization enthalpy is expressed in units of kJ /mol

X(g) + energy → X+(g) + e–.
1st ionisation enthalpy
X+(g) + energy → X++(g) + e–.
2nd ionisation enthalpy

The second ionization enthalpy will be higher than
the first ionization enthalpy because it is more
difficult to remove an electron from a positively
charged ion than from a neutral atom because a
cation has greater ENC than a neutral atom.
In the same way the third ionization enthalpy will
be higher than the second and so on.

(a) Size of the atom - IE decreases as the size of the
atom increases
(b) Nuclear Charge - IE increases with increase in
nuclear charge
(c) The type of electron - Shielding effect, Penetration
effect
(e)Electronic configuration: e.g. noble gases passes
very high value of IE due to stable octet
configuration
Factors affecting Ionisation Enthalpy (ΔiH )

On moving down a group
1. nuclear charge increases
2. Z* due to screening is almost constant
3. number of shells increases, hence atomic size
increases.
4. there is a increase in the number of inner electrons
which shield the valence electrons from the nucleus
Thus IE decreases down the group
Periodic Trend of Ionisation Enthalpy (ΔiH )

On moving across a period(L--->R)
1. the atomic size decreases
2. Effective nuclear charge increases
Thus IE increases along a period
However there are some exceptions also e.g.
 IE of Be is higher than that of B.
 IE of N is higher than that of O.
Periodic Trend of Ionisation Enthalpy

Explain why- (a). IE of Be is higher than that of B.
Ans. - In beryllium(1s2,2s2 ), the electron removed during the
ionization is an s-electron whereas the electron removed
during ionization of boron(1s2,2s2,2p1) is a p-electron. The
penetration of a 2s-electron to the nucleus is more than
that of a 2p-electron; hence the 2p electron of boron is more
shielded from the nucleus by the inner core of electrons
than the 2s electrons of beryllium.
Therefore, it is easier to remove the 2p-electron from boron
compared to the removal of a 2s- electron from beryllium.
Thus, boron has a smaller first ionization.

(b) Why IE of N is higher than that of O.
Ans. The first ionization enthalpy of oxygen compared to
nitrogen is smaller. This arises because in the nitrogen
atom(1s2,2s2,2p3) three 2p-electrons reside in different
atomic orbitals (Hund’s rule) whereas in the oxygen atom
(1s2,2s2,2p4), two of the four 2p-electrons must occupy the
same 2p-orbital resulting in an increased electron-electron
repulsion. Consequently, it is easier to remove the fourth
2p-electron from oxygen than it is, to remove one of the
three 2p-electrons from nitrogen.

Electron Gain Enthalpy (ΔegH)
 When an electron is added to a neutral gaseous atom (X) to
convert it into a negative ion, the enthalpy change
accompanying the process is defined as the Electron
GainEnthalpy (ΔegH) or Electron Affinity.
 Electron gain enthalpy provides a measure of the ease
with which an atom adds an electron to form anion as
represented by equation –
 X(g) + e --- X- (g)+energy
(electrongainenthalpy)

 Depending on the element, the process of
adding an electron to the atom can be either
endothermic or exothermic.
 For many elements energy is released when an
electron is added to the atom and the electron gain
enthalpy is negative.

 ENC- With increase in ENC, the force of attraction exerted by
the nucleus on the electrons increases. Consequently, the
atom has a greater tendency to attract additional electron
i.e., its EGE increases i.e. become more negative.
 ATOMIC SIZE-
With decrease in size ENC increases & hence EGE
increases.
 ELECTRONIC CONFIGURATION-
The value of EGE depends effectively upon electronic
configuration of elements, elements with stable electronic
configuration posses lower (less -ve) value of EGE, e.g.-
Factors Affecting E G E (ΔegH)

A. Noble gases have practically zero or +ve EGEs. This
is because they have no tendency to gain an
additional electron as they already have the stable
ns2np6 configuration
B. Halogens have high electron affinities. This is due
to their strong tendency to gain an additional
electron to change into the stable ns2np6
configuration.
Factors Affecting E G E (ΔegH)

IN A PERIOD-
The EGE increases i.e. become more negative as we move
across a period because the atomic size decreases and hence
the force of attraction exerted by the nucleus on the
electrons increases. Consequently, the atom has a greater
tendency to attract additional electron i.e., its electron
affinity increases
IN A GROUP-
The EGE decreases (-)vely because the atomic size increases
and therefore, the effective nuclear attraction decreases and
thus electron affinity decreases
PERIODIC TREND OF EGE (ΔegH)

 Explain why –
(a). electron gain enthalpy of O is less than that of the S.
(b). electron gain enthalpy of F is less than that of the Cl.
 Ans:- The electron gain enthalpy of O or F is less than that
of the succeeding element. This is because when an electron
is added to O or F, the added electron goes to the smaller n
= 2 quantum level and suffers significant repulsion from the
other electrons present in this level. For the n = 3 quantum
level (S or Cl ), the added electron occupies a larger region
of space and the electron-electron repulsion is much less.

 The tendency of an element in a molecule
to attract the shared pair of electrons
towards itself is known as electronegativity.
It is measured on Pauling scale in which F
(most EN element)is attributed to a value of
4 .
Electronegativity

Periodic trend of EN
In a Group- on moving down the group,
 Z increases but Z* almost remains constant
 number of shells (n) increases
 atomic radius increases
 force of attraction between added electron and
nucleus decreases
 Therefore EN decreases moving down the group

In a Period- On moving across a period left to right
 Z and Z* increases
 number of shells remains constant
 atomic radius decreases
 force of attraction between shared electron and
nucleus increases
Hence EN increases along a period
Periodic trend of EN

Periodicity of Valence or Oxidation States
 Anomalous Properties of Second Period
Elements
Periodic Trends in Chemical Properties

The valence of representative elements is usually (though
not necessarily) equal to the number of electrons in the
outer most orbitals and / or equal to eight minus the
number of outermost Electrons(w.r.t. H)
Some periodic trends observed in the valence of elements
(hydrides and oxides) are shown in Table
Periodicity of Valence or Oxidation States
Group 1 2 13 14 15 16 17
Number of
valence
electron
1 2 3 3 5 6 7
Valence 1 2 3 4 3,5 2,6 1,7

The oxidation state of an element in a particular
compound can be defined as the charge acquired by
its atom on the basis of electronegative
consideration from other atoms in the molecule.
Each group has a common (+)ve or (-)ve oxidation state
And it show gradual change in oxidation state in a
period
Periodicity of Valence or Oxidation States

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