74 CHEMISTRY
The Periodic Table is arguably the most important concept in
chemistry, both in principle and in practice. It is the everyday
support for students, it suggests new avenues of research to
professionals, and it provides a succinct organization of the
whole of chemistry. It is a remarkable demonstration of the
fact that the chemical elements are not a random cluster of
entities but instead display trends and lie together in families.
An awareness of the Periodic Table is essential to anyone who
wishes to disentangle the world and see how it is built up
from the fundamental building blocks of the chemistry, the
chemical elements.
Glenn T. Seaborg
In this Unit, we will study the historical development of the
Periodic Table as it stands today and the Modern Periodic
Law. We will also learn how the periodic classification
follows as a logical consequence of the electronic
configuration of atoms. Finally, we shall examine some of
the periodic trends in the physical and chemical properties
of the elements.
3.1 WHY DO WE NEED TO CLASSIFY ELEMENTS ?
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 synthesise 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
organise 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.
UNIT 3
After studying this Unit, you will be
able to
appreciate how the concept of
grouping elements in accordance to
their properties led to the
development of Periodic Table.
understand the Periodic Law;
understand the significance of
atomic number and electronic
configuration as the basis for
periodic classification;
name the elements with
Z >100 according to IUPAC
nomenclature;
classify elements into s, p, d, f
blocks and learn their main
characteristics;
recognise the periodic trends in
physical and chemical properties of
elements;
compare the reactivity of elements
and correlate it with their
occurrence in nature;
explain the relationship between
ionization enthalpy and metallic
character;
use scientific vocabulary
appropriately to communicate ideas
related to certain important
properties of atoms e.g., atomic/
ionic radii, ionization enthalpy,
electron gain enthalpy,
electronegativity, valence of
elements.
CLASSIFICATION OF ELEMENTS AND
PERIODICITY IN PROPERTIES
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75CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPERTIES
3.2 GENESIS OF PERIODIC
CLASSIFICATION
Classification of elements into groups and
development of Periodic Law and Periodic
Table are the consequences of systematising
the knowledge gained by a number of scientists
through their observations and experiments.
The German chemist, Johann Dobereiner in
early 1800’s was the first to consider the idea
of trends among properties of elements. By
1829 he noted a similarity among the physical
and chemical properties of several groups of
three elements (Triads). In each case, he
noticed that the middle element of each of the
Triads had an atomic weight about half way
between the atomic weights of the other two
(Table 3.1). Also the properties of the middle
element were in between those of the other
two members. Since Dobereiner’s relationship,
referred to as the Law of Triads, seemed to
work only for a few elements, it was dismissed
as coincidence. The next reported attempt to
classify elements was made by a French
geologist, A.E.B. de Chancourtois in 1862. He
arranged the then known elements in order of
increasing atomic weights and made a
cylindrical table of elements to display the
periodic recurrence of properties. This also did
not attract much attention. The English
chemist, John Alexander Newlands in 1865
profounded the Law of Octaves. He arranged
the elements in increasing order of their atomic
weights and noted that every eighth element
had properties similar to the first element
(Table 3.2). The relationship was just like every
eighth note that resembles the first in octaves
of music. Newlands’s Law of Octaves seemed
to be true only for elements up to calcium.
Although his idea was not widely accepted at
that time, he, for his work, was later awarded
Davy Medal in 1887 by the Royal Society,
London.
The Periodic Law, as we know it today owes
its development to the Russian chemist, Dmitri
Mendeleev (1834-1907) and the German
chemist, Lothar Meyer (1830-1895). Working
independently, both the chemists in 1869
proposed that on arranging elements in the
increasing order of their atomic weights,
similarities appear in physical and chemical
properties at regular intervals. Lothar Meyer
plotted the physical properties such as atomic
volume, melting point and boiling point
against atomic weight and obtained a
periodically repeated pattern. Unlike
Newlands, Lothar Meyer observed a change in
length of that repeating pattern. By 1868,
Lothar Meyer had developed a table of the
Element Atomic Element Atomic Element Atomic
weight weight weight
Li 7 Ca 40 Cl 35.5
Na 23 Sr 88 Br 80
K 39 Ba 137 I 127
Table 3.1 Dobereiner’s Triads
Table 3.2 Newlands’ Octaves
Element Li Be B C N O F
At. wt. 7 9 11
12 14 16 19
Element Na Mg Al Si P S Cl
At. wt. 23 24 27 29 31 32 35.5
Element K Ca
At. wt. 39 40
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76 CHEMISTRY
elements that closely resembles the Modern
Periodic Table. However, his work was not
published until after the work of Dmitri
Mendeleev, the scientist who is generally
credited with the development of the Modern
Periodic Table.
While Dobereiner initiated the study of
periodic relationship, it was Mendeleev who
was responsible for publishing the Periodic
Law for the first time. It states as follows :
The properties of the elements are a
periodic function of their atomic
weights.
Mendeleev arranged elements in horizontal
rows and vertical columns of a table in order
of their increasing atomic weights in such a
way that the elements with similar properties
occupied the same vertical column or group.
Mendeleev’s system of classifying elements was
more elaborate than that of Lothar Meyer’s.
He fully recognized the significance of
periodicity and used broader range of physical
and chemical properties to classify the
elements. In particular, Mendeleev relied on
the similarities in the empirical formulas and
properties of the compounds formed by the
elements. He realized that some of the elements
did not fit in with his scheme of classification
if the order of atomic weight was strictly
followed. He ignored the order of atomic
weights, thinking that the atomic
measurements might be incorrect, and placed
the elements with similar properties together.
For example, iodine with lower atomic weight
than that of tellurium (Group VI) was placed
in Group VII along with fluorine, chlorine,
bromine because of similarities in properties
(Fig. 3.1). At the same time, keeping his
primary aim of arranging the elements of
similar properties in the same group, he
proposed that some of the elements were still
undiscovered and, therefore, left several gaps
in the table. For example, both gallium and
germanium were unknown at the time
Mendeleev published his Periodic Table. He left
the gap under aluminium and a gap under
silicon, and called these elements Eka-
Aluminium and Eka-Silicon. Mendeleev
predicted not only the existence of gallium and
germanium, but also described some of their
general physical properties. These elements
were discovered later. Some of the properties
predicted by Mendeleev for these elements and
those found experimentally are listed in
Table 3.3.
The boldness of Mendeleev’s quantitative
predictions and their eventual success made
him and his Periodic Table famous.
Mendeleev’s Periodic Table published in 1905
is shown in Fig. 3.1.
Property Eka-aluminium Gallium Eka-silicon Germanium
(predicted) (found) (predicted) (found)
Atomic weight 68 70 72 72.6
Density / (g/cm
3
) 5.9 5.94 5.5 5.36
Melting point /K Low 302.93 High 1231
Formula of oxide E
2
O
3
Ga
2
O
3
EO
2
GeO
2
Formula of chloride ECl
3
GaCl
3
ECl
4
GeCl
4
Table 3.3 Mendeleev’s Predictions for the Elements Eka-aluminium (Gallium) and
Eka-silicon (Germanium)
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77CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPERTIES
PERIODIC SYSTEM OF THE ELEMENTS IN GROUPS AND SERIES
Fig. 3.1 Mendeleev’s Periodic Table published earlier
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78 CHEMISTRY
3.3 MODERN PERIODIC LAW AND THE
PRESENT FORM OF THE PERIODIC
TABLE
We must bear in mind that when Mendeleev
developed his Periodic Table, chemists knew
nothing about the internal structure of atom.
However, the beginning of the 20
th
century
witnessed profound developments in theories
about sub-atomic particles. In 1913, the
English physicist, Henry Moseley observed
regularities in the characteristic X-ray spectra
of the elements. A plot of
ν
(where
ν
is
frequency of X-rays emitted) against atomic
number (Z ) gave a straight line and not the
plot of
ν
vs atomic mass. He thereby showed
that the atomic number is a more fundamental
property of an element than its atomic mass.
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.
The Periodic Law revealed important
analogies among the 94 naturally occurring
elements (neptunium and plutonium like
actinium and protoactinium are also found in
pitch blende – an ore of uranium). It stimulated
renewed interest in Inorganic Chemistry and
has carried into the present with the creation
of artificially produced short-lived elements.
You may recall that the atomic number is
equal to the nuclear charge (i.e., number of
protons) or the number of electrons in a neutral
atom. It is then easy to visualize the significance
of quantum numbers and electronic
configurations in periodicity of elements. In
fact, it is now recognized that the Periodic Law
is essentially the consequence of the periodic
variation in electronic configurations, which
indeed determine the physical and chemical
properties of elements and their compounds.
Numerous forms of Periodic Table have
been devised from time to time. Some forms
emphasise chemical reactions and valence,
whereas others stress the electronic
configuration of elements. A modern version,
the so-called “long form” of the Periodic Table
of the elements (Fig. 3.2), is the most convenient
and widely used. The horizontal rows (which
Mendeleev called series) are called periods and
the vertical columns, groups. Elements having
similar outer electronic configurations in their
atoms are arranged in vertical columns,
referred to as groups or families. According
to the recommendation of International Union
of Pure and Applied Chemistry (IUPAC), the
groups are numbered from 1 to 18 replacing
the older notation of groups IA … VIIA, VIII, IB
… VIIB and 0.
There are altogether seven periods. The
period number corresponds to the highest
principal quantum number (n) of the elements
in the period. The first period contains 2
elements. The subsequent periods consists of
8, 8, 18, 18 and 32 elements, respectively. The
seventh period is incomplete and like the sixth
period would have a theoretical maximum (on
the basis of quantum numbers) of 32 elements.
In this form of the Periodic Table, 14 elements
of both sixth and seventh periods (lanthanoids
and actinoids, respectively) are placed in
separate panels at the bottom
*
.
3.4 NOMENCLATURE OF ELEMENTS WITH
ATOMIC NUMBERS > 100
The naming of the 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 characterisation, therefore, require highly
* Glenn T. Seaborg’s work in the middle of the 20
th
century starting with the discovery of plutonium in 1940, followed by
those of all the transuranium elements from 94 to 102 led to reconfiguration of the periodic table placing the actinoids
below the lanthanoids. In 1951, Seaborg was awarded the Nobel Prize in chemistry for his work. Element 106 has been
named Seaborgium (Sg) in his honour.
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