1.1. Historical Background of Periodic Classification
Need and Search for Classification
With the discovery of more and more new elements,
it was necessary to organize these elements systematically and need arose for a
frame work in which these elements could be classified and arranged in in order
to facilitate their study and make their study simple and systematic. The
classification of elements enabled the chemists to understand and interpret the
properties of elements in a better way.
There could be many ways of arranging the
elements; firstly they could be classified by their states (solids, liquids or
gases) at a particular temperature, secondly they could be arranged as metals,
non-metals and metalloids and thirdly one might find patterns in their
reactions with oxygen or water or other chemicals. Would one consider trying to
link these properties to the relative atomic masses of the elements?
Previously scientist tried to arrange the
elements in a scientific, systematic and an organized manner on the basis of
their atomic weight (atomic masses) as it was thought that the properties of
elements depended upon their atomic masses (the thought was grounded on
Dalton’s atomic theory). But recently, the basis of classification has been
changed and elements are arranged on the basis of their atomic numbers instead
of their atomic masses.
Different attempts of Classification
Following attempts were made to classify the
known elements:
1. Al-Razi Classification
2. Origin of Classification; Dalton’s
Atomic Theory
3. Dumas Work
4. Prout’s Attempt
5. Dobereiner’s Triads
6. Newland’s Law of Octave
7. Lother Meyer’s Classification
8. Mendeleev’s Classification
9. Modern Periodic table
1. Origin of
Classification
The basis of classification of elements was grounded on the Daltons’ atomic theory put forward by an English scientist, John Dalton in 1808, according to which:
“Atoms of different elements have different atomic masses.”
Thus it was concluded that there is a regular relationship between atomic masses and properties of elements. “This relationship proves to be the corner stone for the future classification of elements”
2. Dumas Work
Dumas (1800-1884), a French chemist arranged the elements on their combining power with chlorine. For example, elements that combined with 1 chlorine atom couldl be arranged in vertical columns in increasing order of their atomic weights and so on.
Reasons for Failure
Dumas attempt of
classification did not gain success as all elements do not combine with
chlorine and few show variable valency.
3. Prout’s Attempt
Prout,
an English chemist considered the atomic weight of hydrogen as the basis of his
classification. He considered that:
“Atomic weights of all elements are simple multiple of the atomic weight
of hydrogen”
Reason for Failure
It could not explain
the fractional atomic weights of elements.
A German chemist, Johann Wolfgang Dobereiner
in 1817 noticed an interesting pattern in certain sets of three similar
elements and classified the similar elements in the groups of three elements
(in the sequence of increasing atomic mass) known as triad. He found that the
atomic mass of the middle element lay (fall) roughly half way (midway) between
the other two (i.e. the lightest and the heaviest) elements of a triad and the
elements of a triad also resemble in properties. He also noticed that the
middle elements had properties that were an average of the other two members of
a triad when arranged by the atomic weights. e.g. He found that the density of
the middle element in most triad is roughly equal to the average of the
densities of the other two elements. The density of strontium (2.6 g/cm3)
for example is close to the average of the densities of calcium (1.55 g/cm3)
and barium (3.51g/cm3).
He put forward Law or rule of Triads,
according to which;
“Central
atom of each set of triad has an atomic mass equal to the arithmetic mean of
the atomic masses of the other two elements.”
OR
Each set
of triad (group of three elements ordered by increasing atomic weights) has
similar properties and atomic weight of the middle element of a triad was
approximately equal to arithmetic mean (average) of the atomic weights of other
two elements of a triad”.
He arranged the elements in triads. The
elements of triad resemble in properties.
He first found alkaline earth metal triad of
Ca, Sr and Ba. He further noticed the same pattern for the alkali metal triad
(Li, Na, K), the halogen triad (Cl, BR, I), Chalcogen (S, Se, Te), metalloid
triad (P, As, Sb) and transitional metal triad (Mn, Cr, Fe).
Elements
|
Atomic Mass
|
Mean Atomic Mass
|
Elements
|
Atomic Mass
|
Mean Atomic Mass
|
Lithium
|
7
|
23
|
Chlorine
|
35.5
|
= 81.25
|
Sodium
|
23
|
Bromine
|
80
|
||
Potassium
|
39
|
Iodine
|
127
|
Elements
|
Atomic Mass
|
Mean Atomic Mass
|
Elements
|
Atomic Mass
|
Mean Atomic Mass
|
Calcium
|
40
|
Sulphur
|
32
|
81.25 |
|
Strontium
|
87
|
Selenium
|
79
|
||
Barium
|
137.3
|
Tellurium
|
127.5
|
Elements
|
Atomic Mass
|
Mean Atomic Mass
|
Elements
|
Atomic Mass
|
Mean Atomic Mass
|
Manganese
|
55
|
Phosphorus
|
31
|
76.37 |
|
Chromium
|
52
|
Arsenic
|
75
|
||
Iron
|
56
|
Antimony
|
121.75
|
Reason for failure
Dobereiner’s law of triad has a very limited
application and could not be extended to the classification of all the elements
as this rule was valid for only very few elements. It failed as this rule was
not applicable for all elements i.e. all elements could not arrange in triads.
4. Newland’s
Law of Octave
In 1864, an English (London) industrial
chemist John Alexander Newland arranged the 56 (60 or 62) known elements by
order of increasing atomic weights into a table along horizontal rows seven
element long with seven vertical columns and proposed has law of octave
accordingly:
“If
elements are arranged in the ascending order of their atomic weights, the
eighth (8th) element following any given element in the series has
nearly same physical and chemical properties as first one” which means that
starting from any element, the properties of every eighth element were similar
to those of first i.e. its properties are a kind of repetition of the first
(like the eight notes of an octave of music or by the analogy with the seven
intervals of the musical scale).
Merits
1. It arranges all 56 elements into tabular form.
2. It arranges all elements with identical properties into same group.
3. Newland’s
classification of elements for the first time showed the existence of periodicity i.e. recurrence of chemical and physical properties of elements at
regular intervals.
4. It also provided a great idea towards the development of modern periodic table.
Objections
1. The
Law of Octave holds up well for the first 17 elements, but it failed rather
badly beyond calcium in predicting
a consistent trend.
2.The
heavier elements could not be accommodated by this arrangement.
3. Moreover
hydrogen as not included in his table.
4. Lother Meyer’s
Classification
In 1869, a German Physicist Julius Lother Meyer (a
contemporary of Mendeleev) classified the known 56
elements on the basis of their increasing atomic weights in
graphical form in nine vertical columns or groups
from I to IX. Meyer’s work was based on physical properties
of elements like atomic volume. He put forward his periodic law, which states
that ‘physical properties of elements are periodic function of their atomic
weights.
The volume occupied by 1 gram atomic weight
or 1 gram atom or 1 gram mole (i.e. 6.02 x 1023 atoms) of any element in solid
state is called atomic volume which is a rough measure of the relative sizes of
atoms.
Lother Meyer’s Atomic Volume Curve
Meyer arranged the elements by plotting a
graph between atomic volumes of elements (on y-axis) against their increasing
atomic masses (on x-axis).
The plot gave a curve called Atomic Volume
Curve, consisted of sharp peak (crests) and broad minima (troughs).
The curve exhibits periodicity as similar elements occupy same positions on the
curve. For example, the highly reactive alkali metals (Li, Na, K, Rb, Cs)
occupy the peak of the curve thereby showing that these elements have largest
atomic volumes.
According to Meyer, the occupying of similar
elements on same positions on the curve was called periodicity. The regular
spacing of the highest points and occupying of similar elements on the same
positions on the curve confirmed the idea of periodicity, suggested by Newland.
[Meyer was the first scientist who considered valency as a period property.]
Meyer’s curve showed the following
characteristics and periodicity:
1. Chemically similar elements occupy
similar position on the curves. For example; Alkali Metals like Li, Na, K etc.
occupy the peaks of the curve indicating that they have largest atomic volumes
than those of neighbouring elements while ascending portion of the curve just
before the peak is occupied by halogens showing their smallest atomic volumes. The
crest of each wave is occupied by an alkali metal and trough by an element of
small chemical affinity.
2. Alkali
metals occupy the peaks or crests of the curves.
3. Weak metals or elements of small
chemical affinity or transition metals occupy the troughs or minima of the curve.
4. Electronegative and gaseous volatile
elements or acidic oxides forming elements are located on the ascending
portions of the curve.
5. Electropositive or transition elements
or elements with high melting points are found on the descending portions of
the curve.
6. Midway
of ascending portions of curve is occupied by halogens.
7. Midway
of descending portions of curve is occupied by alkaline earth metals.
Meyer’s curve shows the following
characteristics and periodicity:
1. Peaks/crests of the curve
|
Alkali
metals.
|
2. Troughs/minima of the curve
|
Non-metals
or weak metals or transition metals
|
3. Ascending portions of curve
|
Electronegative
& gaseous volatile acidic oxides forming elements
|
4. Descending portions of curve
|
Electropositive
or transition elements or elements with high melting points
|
5. Midway of ascending portion of
the curve
|
Halogens
|
6. Midway of descending portion of
the curve
|
Alkaline
earth metals.
|
Objections
Lother Meyer’s Periodic Classification could not receive proper attention due to following reasons:
1. Meyer’s Periodic Table was incomplete as he left no blank spaces for undiscovered elements as compared with Mendeleev’s Periodic Table (which was characterized by remarkable predictions of discoveries of certain elements).
2. no logical basis for classification based on various physical properties such as atomic volume.
3. Chemical properties of elements were completely ignored.
4. His table was non-reproducible form of periodic table.
Mendeleev’s
Classification
Most of the credit of the development of
periodic classification of elements must go to a Russian chemist Dmitri Ivanovitch (D.I.) Mendeleev who presented the most useful and most systematic
scheme for periodic classification of elements in March 1869. (Mendeleev’s was
notorious for cutting his hair only once a year). Up till 1869, only 63 elements
were known. Mendeleev arranged the elements in the sequence of their increasing
atomic weights. He arranged the elements of similar properties in vertical
columns and dissimilar elements in horizontal rows.
Mendeleev’s Periodic Law
Since similar properties occurred periodically
as a function of atomic mass, Mendeleev stated the Periodic Law as;
“The
physical and chemical properties of elements are a periodic function of their
atomic weights i.e. if the elements are arranged in ascending order of their
atomic weights, their properties repeat in a periodic manner.”
Mendeleev’s work was an extension of Newland’s octaves. The basis of his
classification was the chemical properties of
elements. Mendeleev arranged the known 63 elements in the sequence of their increasing atomic weights,
placing the elements with similar chemical properties vertically beneath each
other. In his table, similar properties occurred periodically i.e. repeated
themselves at intervals as a function of atomic weights.
Features of Mendeleev’s Periodic
Table
Following are the main features of table:
1. The elements are arranged in ascending order of their atomic masses.
2. The Mendeleev’s periodic table consisted of 8 vertical columns called groups (i.e. group I to VIII) containing similar elements and 12 horizontal rows called Series or Periods having dissimilar elements.
3. The groups are further divided into sub groups A and B. This sub division allowed him to place elements with slightly different properties in same group thereby maintaining periodicity.
4. The elements in each group have similar chemical properties but their physical properties change gradually down the group.
5. The
group number indicates the highest valency of element that it can attain.
6. Mendeleev’s
table clearly and forcefully proved the concept of periodicity.
7. Mendeleev’s
table contained vacant spaces for undiscovered (unknown or missing) elements
with predicted atomic masses 44, 68, 72
and 100. He named them eka-boron, eka-aluminium and eka silicon.
Advantages
(Merits) of Mendeleev’s Table
1. Systematic Study of Elements
2. Special Emphasis on Chemical
Similarities
3. Determination
& Correction of Doubtful Atomic Weights
4. Vacant Spaces for Undiscovered Elements
& Prediction of Properties of 3 Unknown Elements
5. Separate
Group for A set of 3 Elements
1. Systematic Study of
Elements
Mendeleev’s table helped chemist to study the
elements more easily and systematically as it had reduced or restricted the
study of elements into a study of eight groups only. (The study of chemistry of
only one element of any group, is largely enough to predict the properties of
the other elements of the same group). For instance, the study of sodium metal
helped chemist to a large extent to predict the properties of its other group
elements like K, Rb, Cs.
2. Special Emphasis on
Chemical Similarities/ Properties take precedence over atomic weights
Mendeleev disregarded atomic masses as the
only criteria for assigning places to elements because at that time, the atomic
weights of many elements were not accurately known; nor was it certain that all
the elements has been discovered. Great emphasis was laid on chemical
similarities of elements. Thus if the properties of an element suggested that
it was out of place in the sequence of atomic masses, it was placed according
to its properties rather than its mass.
He placed 3 misfit pairs of element in his
table denying his own periodic law to maintain periodicity which was supposed
to be more important than following law. Hence some elements of higher atomic
mass were placed before elements of lower atomic mass. e.g.
Argon (40)
|
precedes
|
potassium
|
(39)
|
|
Cobalt (58.9)
|
precedes
|
nickel
|
(58.7)
|
|
Tellurium(127.6)
|
precedes
|
iodine
|
(126.9)
|
|
Thorium (232)
|
precedes
|
protactinium
|
(231)
|
3. Determination & Correction
of Doubtful Atomic Weights
Mendeleev’s classification helped in
correcting the doubtful atomic weights of a number of elements which has been
assigned incorrect values and put them in proper places in the periodic table.
Mendeleev used the formula; Atomic
weight =
Equivalent weight x valency for calculating atomic weights of
elements. The equivalent weight of elements can be calculated by any of the
known methods and the valency can be obtained by consulting the periodic table.
For example,
(i) atomic
weight of Be was correct from 13.5 to 9. With this atomic weight, Be was given
a position between Li–7 and B–11.
The properties of Be justify this position in periodic table.
(ii) Similarly atomic weight of indium was
readjusted from 75.80 to 113. (Indium was supposed to have the valency 2 and
equivalent weight 38, so its atomic weight would be 2 x 38 = 76 and with this
atomic weight it would be placed between Zn-65 and Sr-87. There was no place
between Zn and Sr. Mendeleev suggested if indium were taken as trivalent, its
atomic weight would be 3 x 38 = 114 and thus would get the place between Cd –
112 and Sn -118 that justified its position.
(iii) Also,
atomic weight of Cr that had been an atomic weight of 43 was recalculated and
found to be 52 and allocated proper
place to it.
4. Vacant
Spaces for Undiscovered Elements & Prediction of Properties of 3 Unknown
Elements
In order to maintain families of chemically
similar elements, he left blank spaces in his table for undiscovered elements
after boron, aluminium and silicon which allowed his theory to be tested.
Comparing the properties of their group elements, he successfully predicted the
three unknown elements, which he named Eka-Boron, Eka-Aluminium and Eka-Silicon
(eka means first i.e. eka-silicon means literally first comes silicon and then
comes unknown element). This prediction helped in their discovery. By 1886,
chemists had discovered all the three elements and had been named as scandium (
), gallium (
) and germanium (
) respectively. Mendeleev’s predictions of their
properties proved to be remarkably accurate. e.g.
Comparative Properties of eka-Aluminium and Gallium
S. #
|
Properties
|
Predicted Properties of eka-aluminium
|
Observed Properties of Gallium
|
1.
|
Atomic weight
|
About 68
|
69.7
|
2.
|
Density
|
5.9 g/cm3
|
5.94 g/cm3
|
3.
|
Valency
|
Three (3)
|
Three (3)
|
4.
|
Melting point
|
Low
|
30°C (29.8°C)
|
5.
|
Formula of oxide
|
R2O3 or eka Al2O3
|
Ga2O3
|
6.
|
Formula of chloride
|
RCl3 or eka-AlCl3
|
GaCl3
|
7.
|
Density of oxide
|
Eka-AlCl3 is more volatile than
ZnCl2
|
GaCl3 is more volatile than ZnCl2
|
8.
|
Chemistry of hydroxide
|
Eka-Al(OH)3 dissolves in both
acids and bases i.e. it is amphoteric.
|
Ga(OH)3 dissolves in both acids
and bases i.e. it is amphoteric.
|
Comparative Properties of eka-Silicon and Germanium
S. #
|
Properties
|
Predicted Properties of
eka-silicon
|
Observed Properties of Germanium
|
1.
|
Atomic weight
|
About 72
|
72.6
|
2.
|
Density
|
5.5 g/cm3
|
5.47 g/cm3
|
3.
|
Valency
|
Four (4)
|
Four (4)
|
4.
|
Melting point
|
High (825°C)
|
938°C (947)
|
5.
|
Formula of oxide
|
RO2 or eka-SiO2
|
GeO2
|
6.
|
Formula of chloride
|
RCl4 or eka-SiCl4
|
GeCl4
|
7.
|
Density of oxide
|
4.7 g/cm3
|
4.70 g/cm3
|
8.
|
Density of chloride
|
1.9 g/cm3 (liquid)
|
1.84 g/cm3
|
9.
|
B.P. of Chloride
|
Less than 100°C
|
86°C
|
10.
|
Colour
|
Dark gray
|
Light gray (grayish white)
|
11.
|
Atomic volume
|
13 cm3
|
13.6 cm3
|
5. Separate Group for A set of 3
Elements
Mendeleev observed that a set of 3 elements
i.e. Fe, Co, Ni, and Ru, Rh, Pd and Os,
Ir, Pt had very similar properties
and could not be assigned to any particular group. He, therefore, placed these
elements at one place in Group VIII.
Defects or Demerits of Mendeleev’s Periodic Table
1. Failure to explain
atomic structure
2. No place for isotopes of elements
3. Failure to place rare earth (Lanthanides and
Actinides) in the main body of periodic table
4. Anomalous position of hydrogen
5. Group number does not represent valency
6. Neglection of variable valency
7. Anomalous or Misfit Pairs of elements
(i) Elements Cu, Ag, Au were placed with
dissimilar elements Li, Na, K, Rb, Cs.
(ii) Similar elements Cu and Hg were placed
separately.
(iii) Elements of higher atomic weight placed
earlier than elements of lighter atomic weights.
1. Failure to explain atomic structure
Mendeleev’s periodic
table failed to account for atomic structure as it was based on atomic weight
and not on atomic number. Also Mendeleev’s table was silent about electronic
configuration of elements.
2. No place for isotopes of elements
There was no room for isotopes in Mendeleev’s
table as it was not possible to accommodate the large number of isotopes in the
periodic table.
3. Failure to place rare earth
in the main body of periodic table
Lanthanides (elements
with atomic numbers 58 to 71) and actinides (elements with atomic numbers 90 to
103) had not been placed in the main body of the periodic table. Rather they
had been given a separate position at the bottom of the periodic table.
4. Anomalous position of hydrogen
The
placement of hydrogen in group I along with alkali metals was a matter of
dispute.
5. Group number does not represent valency
Group number did not
represent the valency of the elements e.g. excepting osmium, elements in group
VIII did not show a valency of 8. Also the elements in the middle of the long periods
(e.g. Mn, Cr etc.) exhibited variable valency.
6. Neglection of variable valency
Elements
with variable valencies were considered to have fixed valency.
7. Anomalous or Misfit
Pairs of elements
(i). Dissimilar elements
placed in the same group
Many elements with
dissimilar properties had been placed in the same group e.g. Alkali metals and
coinage metals were place in same group in spite of their entirely different
properties. However division of groups into sub-groups solved the issue later.
Also Mn had been placed with halogens.
(ii). Similar elements
placed in different groups
Similar pairs of elements were placed in
different groups. For instance Ba and Pb resemble in many properties but they
were kept in different groups. Moreover similar elements Cu and Hg were also
placed separately.
(iii). Misfit position of elements
of group VIII
Group VIII has 9
elements placed in three available columns. These elements did not fit in the
system.
(iv). Position of four anomalous pairs of elements
For placing elements
in the proper groups, certain elements of higher atomic masses precede those of
lower atomic masses in Mendeleev’s table. This was against Mendeleev’s Periodic
Law. These misfit pairs of elements were:
Argon (40)
|
precedes
|
potassium
|
(39)
|
|
Cobalt (58.9)
|
precedes
|
nickel
|
(58.7)
|
|
Tellurium(127.6)
|
precedes
|
iodine
|
(126.9)
|
|
Thorium (232)
|
precedes
|
protactinium
|
(231)
|
