Classification of elements, unit iii class 11

BY : - MR. ARVIND CHAUHAN
PGT CHEMISTRY

▪ THE METHOD OF GROUPING THE
ELEMENTS INTO DIFFERENT
CLASSES IS KNOWN AS PERIODIC
CLASSIFICATION OF ELEMENTS.
INTRODUCTION

We know by now that the elements are the basic units of all types of
matter. In 1800, only 31 elements were known. By 1865, the number
of identified elements had more than doubled to 63. At present 114
elements are known. Of them, the recently discovered elements are
man-made. Efforts to synthesize new elements are continuing. With
such a large number of elements it is very difficult to study
individually the chemistry of all these elements and their innumerable
compounds individually. To ease out this problem , scientists searched
for a systematic way to organize their knowledge by classifying the
elements. Not only that it would rationalize known chemical facts
about elements, but even predict new ones for undertaking further
study.
WHY DOWE NEED TO CLASSFY THE ELEMENTS

GENESIS OF
PERIODIC
CLASSIFICATION
OF ELEMENTS

IN 1817, Dobereiner found that the
elements could be arranged in group
of three called triads in such a way
that the middle element had an
atomic weight almost the average of
the other two.
Element At.
Weight
Element At.
weight
Elemen
t
At.
weight
Li 7 Ca 40 Cl 35.5
Na 23 Sr 88 Br 80
K 39 Ba 137 I 127
DOBEREINER’S TRIAD

John A. R. Newland (1865) found that if the
elements were arranged in increasing order of
there atomic weights, the properties of every
eight elements were similar to those of first one.
Newland called it law of octave.
But law of octave seemed to be true only for
elements up to calcium.
Although his idea was not widely accepted at that
time, he, for his work, later awarded davy medal
in 1887 by Royal Society, London.
NEWLAND’S LAW OF OCTET

Element Li Be B C N O F
At.
weigh
t
7 9 11 12 14 16 19
Element Na Mg Al Si P S Cl
At.
Weight
23 24 27 29 31 32 35.5
Element K Ca
At.
weight
39 40

▪ MENDELEEV’S PERIODIC LAW STATES THAT:-
“Properties of elements are periodic
functions of their atomic weight".
THE PERIODIC LAW, 1869 ,
DMITRI MENDELEEV

Classification of elements, unit iii class 11

Mendeleev’s periodic table is based on atomic weight.
Mendeleev’s periodic law states that – Properties of elements are
periodic functions of their atomic weight.
This table consist of seven horizontal and columns. Zero
group is not included in this table. eight vertical
Mendeleev arranged 65 elements in table.
It was clear, that there were a number of gaps of “missing”
elements in the table. Mendeleev, made several predictions
about properties of some missing elements.
Mendeleev predicted four elements: ekaboron (Eb),
ekaaluminium (El), ekamanganese (Em), and ekasilicon
(Es)

Property Eka- aluminium
(predicted)
Gallium (found) Eka-silicon
(predicted)
Germanium
(found)
At. Weight 68 70 72 72.6
Density (g/Cm3) 5.9 5.94 5.5 5.36
Melting point
(k)
Low 302.93 High 1231
Formula of
oxide
E2O3 Ga2O3 EO2 GeO2
Formula of
choride
ECl3 GaCl3 ECl4 GeCl4
MENDELEEV’S PREDICTION FOR
ELEMENTS: EKA-ALUMINIUM, EKA-SILICON

Position of hydrogen in the periodic table wasnot correctly
defined.
In certain pair of elements the increasing order of atomic
masses was not obeyed.
Isotopes have not been given separate places inthe periodic
table.
Some similar elements were separated and dissimilar elements
were grouped together.
Mendeleev could not explain the cause ofperiodicity among the
elements.
DEFECTS OF MENDELEEV’S PERIODIC TABLAE

IN 1942 MOSELEY GIVE THIS LAW BASED ON ATOMIC NUMBER.
Modern periodic law states that- The physical
and chemical properties of the elements
are the periodic function of their atomic
numbers.
MODERN PERIODIC LAW

MODERN PERIODIC TABLE IS ALSO KNOWN AS LONG FORM
OF PERIODIC TABLE OR BHOR’S TABLE
➢ The long form of periodic table consists of horizontal rows called periods
and vertical columns calledgroups.
➢ There are seven periods and each period state with a different principal
quantum number.
➢ There are 18 groups in the long form of periodic table.
➢ There are two series at the bottom of the periodic table.
➢ These consist of fourteen elements after lanthanum and fourteen Elements
which follow actinium. The first row is called lanthanides and second row
is called actinides.

Period (n) Orbital energy filled No. of elements
First 1s 2
Second 2s2p 8
Third 3s3p 8
Fourth 4s3d4p 18
Fifth 5s4d5p 18
Sixth 6s4f5d6p 32
Seventh 7s5f6d7p incomplete

The naming of new elements had been traditionally the privilege of the discoverer
(or discoverers) and the suggested name was ratified by the IUPAC. In recent years
this has led to some controversy. The new elements with very high atomic numbers
are so unstable that only minute quantities, sometimes only a few atoms of them are
obtained. Their synthesis and characterization, therefore, require highly
sophisticated costly equipment and laboratory. Such work is carried out with
competitive spirit only in some laboratories in the world. Scientists, before
collecting the reliable data on the new element, at times get tempted to claim for its
discovery. For example, both American and Soviet scientists claimed credit for
discovering element 104 . The Americans named it rutherfordium whereas soviets
named it kurchatovium. To avoid such problems, the IUPAC has made
recommendation that until a new element’s discovery is proved, and its name is
officially recognized, a systematic nomenclature be derived directly from the atomic
number of the element using the numerical roots for 0 and numbers 1-9. these are
shown in the table in the next slide.
NOMENCLATURE OF ELEMENTS WITH ATOMIC NO. MORE THAN 100

Digit Name Abbreviation
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

IUPAC NOMENCLATURE FOR ELEMENTS WITH ATOMIC NUMBER GREATER THAN 100

ELECTRONIC
CONFIGURATION OF
ELEMENTS

The distribution of electrons into
orbitals of an atom is called its
electronic configuration.

➢ The period indicates the value of n for the outermost or valence shell. In
other words, successive period in the Periodic Table is associated with the
filling of the next higher principal energy level (n = 1, n = 2, etc.).
➢ Number of elements in each period is twice the number of atomic orbitals
available in the energy level that is being filled.
➢ The first period (n = 1) starts with the filling of the lowest level (1s) and
therefore has two elements – hydrogen (ls1) and helium (ls2)
➢ The second period (n = 2) starts with lithium and the third electron enters the
2s orbital. The next element, beryllium has four electrons and has the
electronic configuration 1s22s2. Starting from the next element boron, the 2p
orbitals are filled with electrons when the L shell is completed at neon
(2s22p6). Thus there are 8 elements in the second period.
ELECTRONIC CONFIGURATION IN PERIOD

➢ The third period (n = 3) begins at sodium, and the added electron enters in 3s
orbital. Successive filling of 3s and 3p orbitals gives rise to the third period of 8
elements from sodium to argon.
➢ The fourth period (n = 4) starts at potassium, and the added electrons fill up the 4s
orbital. Now you may note that before the 4p orbital is filled, filling up of 3d orbitals
becomes energetically favorable and we come across the so called 3d transition series
of elements. This starts from scandium (Z = 21) which has the electronic configuration
3d1 4s2. The 3d orbitals are filled at zinc (Z=30) with electronic configuration 3d104s2 .
The fourth period ends at krypton with the filling up of the 4p orbitals. Altogether we
have 18 elements in this fourth period.
➢ The fifth period (n = 5) beginning with rubidium is similar to the fourth period and
contains the 4d transition series starting at yttrium (Z = 39). This period ends at xenon
with the filling up of the 5p orbitals.
➢ The sixth period (n = 6) contains 32 elements and successive electrons enter 6s, 4f, 5d
and 6p orbitals, in the order – filling up of the 4f orbitals begins with cerium (Z = 58)
and ends at lutetium (Z = 71) to give the 4f-inner transition series which is called the
lanthanoid series

➢ The seventh period (n = 7) is similar to the sixth period with
the successive filling up of the 7s, 5f, 6d and 7p orbitals
and includes most of the man- made radioactive elements.
This period will end at the element with atomic number 118
which would belong to the noble gas family. Filling up of the
5f orbitals after actinium (Z = 89) gives the 5f-inner
transition series known as the actinoid series.
➢ The 4f-and 5f-inner transition series of elements are placed
separately in the Periodic Table to maintain its structure
and to preserve the principle of classification by
keeping elements with similar properties in a single
column.

ATOMIC No. SYMBOL ELECTRONIC COIGUNFRATION
3 Li 1s22s1 (or) [He]2s1
11 Na 1s22s22p63s1 (or) [Ne]3s1
19 K 1s22s22p63s23p64s1 (or) [Ar]4s1
37 Rb 1s22s22p6sup3s23p63d104s24p65s1 (or) [Kr]5s1
55 Cs 1s22s22p63s23p63d104s24p64d105s25p66s1 (or) [Xe]6s1
87 Fr [Rn]7s1
Elements in the same vertical column or group have similar valence shell electronic
configurations, the same number of electrons in the outer orbitals, and similar
properties. For example, the Group 1 elements (alkali metals) all have ns1 valence
shell electronic configuration as shown in the table.
GROUPSWISE ELECTRONIC CONFIGURATION

➢ The elements of Group 1 (alkali metals)
and Group 2 (alkaline earth metals) which
have ns1 and ns2 outermost electronic
configuration belong to the s-Block
Elements.
➢ They are all reactive metals with low
ionization enthalpies.

➢ The p-Block Elements comprise those
belonging to Group 13 to 18 and these
together with the s-Block Elements are
called the representative Elements or
Main Group Elements.
➢ Their general outer shell configuration
is ns2 np1-6 where n=1-7.
➢ There are 36 p-block elements in the
periodic table.

➢ d-block elements are also called transitionelements.
➢ The general outer electronic configuration ofthese
elements is (n-1)d1-10ns0-2 .
➢ They are allmetals.
➢ Zn, Cd and Hg which have the electronic configuration,
(n-1)d10ns2 do not show most of the properties of
transition elements.

➢ f-block elements are also called inner transition elements.
➢ Their general outer shell electronic configurationis
(n-2)f1-14 (n-1)d0-1ns2.
➢ They are allmetals.
➢ The 14 elements of the first series from Ce(Z = 58) to Lu(Z = 71) are
called lanthanides. They are also called rare earth metals.
➢ The 14 elements of the first series from Th(Z = 90) to Lr (Z = 103)
are called actinides or actinoids.
➢ The elements after uranium are calledTransuranium Elements

Metals
➢ Metals comprise more than 78% of
known elements.
➢ All s-block, d-block and f-block elements
are metals.
➢ Metallic character increases down a group
and decreases along a period as we move
from left to right.

Non-metals
➢Non-metals are a few, i.e., 18.
➢These are H, C, N, O, P,S, Se, F, Cl, Br, I, He, Ne, Ar, Kr,
Xe, Rn and Uuo.
➢The nonmetallic character increases as one goes from left to
right across the Periodic Table.
Metalloids
➢The Elements which lie along the borderline between metals
and non-metals are called semimetals or metalloids.
➢These are : Si, Ge, As, Sb, Te, Po and At.

❖There are many observable patterns in the physical
and chemical properties of elements as we descend in
a group or move across a period in the Periodic Table.
For example, within a period, chemical reactivity
tends to be high in Group 1 metals, lower in elements
towards the middle of the table, and increases to a
maximum in the Group 17 non-metals. Likewise
within a group of representative metals (say alkali
metals) reactivity increases on moving down the
group, whereas within a group of non-metals (say
halogens), reactivity decreases down the group.

❖There are numerous physical properties of elements such as
melting and boiling points, heats of fusion and vaporization,
energy of atomization, etc. which show periodic variations.
However, we shall discuss the periodic trends with respect to
atomic and ionic radii, ionization enthalpy, electron gain
enthalpy and electronegativity.

➢ The atomic size can be find out by knowing
the distance between the atoms in the
combined state.
ATOMIC RADIUS : It is defined as “the distance from the centre
of nucleus upto the outermost electron cloud”.

➢ To estimate the size of an atom of a
non-metallic element is to measure
the distance between two atoms when
they are bound together by a single
bond in a covalent molecule and
from this value, the “Covalent
Radius” of the element can be
calculated.
➢ For example, the bond distance in the
chlorine molecule (Cl2) is 198 pm
and half this distance (99 pm), is taken
as the atomic radius of chlorine.
COVALENT RADIUS : It is defined as “ the half of internuclear
distance between the nuclei of two same atoms bonded by a
single bond.

❑ For metals, we define the term “MetallicRadius” .
❑ For example, the distance between two adjacent copper atoms in solid
copper is 256 pm; hence the metallic radius of copper is assigned a
value of 128 pm.
❑ Atomic Radius to refer to both covalent or metallic radius depending on
whether the element is a non- metal or a metal. Atomic radii can be
measured by X- ray or other spectroscopic methods.

Vander waal radius : it is defined as “ the half of
internuclear distance between the nuclei of two nearest non-
bonded atoms of same element in solid state.”

VARIATION OF ATOMIC RADIUS WITH
ATOMIC NUMBER ACROSS THE SECOND PERIOD

VARIATION OF ATOMIC RADIUS WITH
ATOMIC NUMBER FOR ALKALI METALS AND HALOGENS

❖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.

❖A cation is smaller than its parent
atom because ithas fewer electrons
while its nuclear charge remains
the same. Also no. Of shells
decrease.
❖An anion is bigger in size than its
parent atom because it has more
electrons and same nuclear charge.
For example, the ionic radius of fluoride ion (F-) is 136 pm whereas the
atomic radius of fluorine is only 64 pm. On the other hand, the atomic radius
of sodium is 186 pm compared to the ionic radius of 95 pm for Na+.

❑It represents the energy required to remove an electron from
an isolated gaseous atom (X) in its ground state.
❑It is denoted by (Δi H).
❑The successive ionization enthalpies follow the sequence:
Δi H3 > Δi H2 > Δi H1 .
❑It is expressed in units of kJ mol-1 .
IONIZATION ENERGY/ENTHALPY

❑ Ionization enthalpy generally increases as we go across a period and
decreases as we descend in a group.
❑ The effective nuclear charge experienced by a valence electron in an
atom will be less than the actual charge on the nucleus because of
“shielding” or “screening” of the valence electron from the nucleus by
the intervening core electrons.
❑Shielding is effective when the orbitals in the inner shells are
completely filled.
❑Across a period, increasing nuclear charge outweighs the shielding.
Consequently, the outermost electrons are held more and more tightly
and the ionization enthalpy increases across a period.

❑ When we consider the same principal quantum level, an s-electron is attracted
to the nucleus more than a p- electron. In beryllium, the electron removed
during the ionization is an s-electron whereas the electron removed during
ionization of boron 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
enthalpy than beryllium.

VARIATION OF FIRST IONIZATION ENTHALPIES (ΔI H)
WITH ATOMIC NUMBER FOR ELEMENTS WITH Z = 1 TO 60

(a)FIRST IONIZATION ENTHALPIES (ΔI H) OF ELEMENTS OF THE SECOND
PERIOD AS A FUNCTION OF ATOMIC NUMBER (Z)
(b) (ΔI H) OF ALKALI METALS AS A FUNCTION OFZ.
(a) (b)

ELECTRON GAIN ENTHALPIES* / (KJ MOL–1) OF
SOME MAIN GROUP ELEMENTS

❑As a general rule, electron gain enthalpy becomes more negative with
increase in the atomic number across a period. The effective nuclear
charge increases from left to right across a period and consequently it will
be easier to add an electron to a smaller atom since the added electron on
an average would be closer to the positively charged nucleus.
❑Electron gain enthalpy to become less negative as we go down a group
because the size of the atom increases and the added electron would be
farther from the nucleus.

❑A qualitative measure of the ability of an atom in a chemical
compound to attract shared electrons to itself is called
electronegativity.
❑It is not a measureable quantity.
❑However, a number of numerical scales of electronegativity of elements
viz., Pauling scale, Mulliken-Jaffe scale, Allred- Rochow scale have been
developed. The one which is the most widely used is the Pauling scale.
Linus Pauling, an American scientist, in 1922 assigned arbitrarily a value
of 4.0 to fluorine, the element considered to have the greatest ability to
attract electrons.
❑Though it is not a measurable quantity, it does provide a means of
predicting the nature of force that holds a pair of atoms together.

ELECTRONEGATIVITY VALUES (ON PAULING SCALE)
ACROSS THE PERIODS

ELECTRONEGATIVITY VALUES (ON PAULING SCALE)
DOWN A FAMILY

❑ Electronegativity is inversely related to the metallic
properties of elements.
❑ Thus, the increase in electronegativities across a period is
accompanied by an increase in nonmetallic properties (or
decrease in metallic properties) of elements.
❑ Similarly, the decrease in electronegativity down agroup is
accompanied by a decrease in nonmetallic properties (or
increase in metallic properties) of elements.

THE PERIODIC TRENDS OF ELEMENTS
IN THE PERIODIC TABLE

Most of the trends in chemical properties
of elements, such as diagonal
relationships, inert pair effect, effects of
lanthanide contraction etc. will be dealt
with along the discussion of each group
in later units. In this section we shall
study the periodicity of the valence state
shown by elements and the anomalous
properties of the second period
elements (from lithium to fluorine).

❑ THE VALENCE OF REPRESENTATIVE ELEMENTS IS USUALLY
(THOUGH NOT NECESSARILY) EQUAL TO THE NUMBER OF
ELECTRONS IN THE OUTERMOST ORBITALS AND / OR EQUAL
TO EIGHT MINUS THE NUMBER OF OUTERMOST ELECTRONS
AS SHOWN IN THE TABLE.
PERIODICITY VALENCY AND OXIDATION STATE

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.
OXIDATION STATE

PERIODIC TRENDS IN VALENCE OF ELEMENTS AS
SHOWN BY THE FORMULAS OF THEIR COMPOUNDS

❑ The first element of each of the groups 1 (lithium) and 2 (beryllium)
and groups 13-17 (boron to fluorine) differs in many respects from the
other members of their respective group. For example, lithium unlike
other alkali metals, and beryllium unlike other alkaline earth metals,
form compounds with pronounced covalent character; the other
members of these groups predominantly form ionic compounds.
❑ The behaviour of lithium and beryllium is more similar with the second
element of the following group i.e., magnesium and aluminium,
respectively. This sort of similarity is commonly referred to as diagonal
relationship in the periodic properties.
ANOMALOUS BEHAVIOUR OS SECOND PERIOD ELEMENTS

❑ The anomalous behaviour is attributed due to their :
Small size,
large charge/ radius ratio
high electronegativity of the elements
❑The first member of p-block elements displays greater ability to form
pπ – pπmultiple bonds to itself (e.g., C = C, C ≡ C, N = N, N ≡ Ν) and
to other second period elements (e.g., C = O, C = N, C ≡ N, N = O)
compared to subsequent members of the same group.

❑The periodicity is related to electronic configuration.
❑ Electron gain enthalpies become more negative across a period. In other words, the
ionization enthalpy of the extreme left element in a period is the least and the
electron gain enthalpy of the element on the extreme right is the highest negative
(note : noble gases having completely filled shells have rather positive electron
gain enthalpy values). This results into high chemical reactivity at the two extremes
and the lowest in the centre. Thus, the maximum chemical reactivity at the extreme
left (among alkali metals) is exhibited by the loss of an electron leading to the
formation of a cation and at the extreme right (among halogens) shown by the gain
of an electron forming an anion.

❑The chemical reactivity of an element can be best shown by its reactions
with oxygen and halogens.
❑The normal oxide formed by the element on extreme left is the most basic
(e.g., Na2O), whereas that formed by the element on extreme right is the
most acidic (e.g., Cl2O7). Oxides of elements in the centre are amphoteric
(e.g., Al2O3, As2O3) or neutral (e.g., CO, NO, N2O). Amphoteric oxides
behave as acidic with bases and as basic with acids, whereas neutral
oxides have no acidic or basic properties.

❑ Among transition metals (3d series), the change in atomic radii
is much smaller as compared to those of representative
elements across the period. The change in atomic radii is still
smaller among inner- transition metals (4f series). The
ionization enthalpies are intermediate between those of s- and
p-blocks.As a consequence, they are less electropositive than
group 1 and 2 metals.

Classification of elements, unit iii class 11

More Related Content

What's hot (20)

Similar to Classification of elements, unit iii class 11 (20)

Recently uploaded (20)

Classification of elements, unit iii class 11