Chemistry · Ch 3 — Classification of Elements and Periodicity in Properties

Genesis of Periodic Classification

3.2

Genesis of Periodic Classification

3.2 Genesis of Periodic Classification

The modern Periodic Table did not appear overnight. It is the product of nearly a century of observations, guesses, corrections, and bold predictions by several scientists. Each contributor added a piece to the puzzle, and the story of how elements came to be classified is as instructive as the table itself.

Dobereiner's Triads (1829)

The German chemist Johann Dobereiner was the first to notice that elements could be grouped in sets of three — triads — that shared similar chemical and physical properties. By 1829, he had identified several such triads.

\text{Atomic weight of the middle element} \approx \frac{\text{Atomic weight of first element} + \text{Atomic weight of third element}}{2}

For example, in the triad lithium–sodium–potassium:

Element	Atomic weight
Li	7
Na	23
K	39

The atomic weight of sodium (23) is almost exactly the average of lithium (7) and potassium (39): $(7 + 39)/2 = 23$ .

Dobereiner also observed that the properties of the middle element — its reactivity, melting point, density — fell between those of the other two members. This was the first hint that atomic weight and properties were connected in a regular way.

Watch out

Dobereiner's Law of Triads worked only for a few groups (Li–Na–K, Ca–Sr–Ba, Cl–Br–I). When scientists tried to extend it to other elements, it failed. Most chemists dismissed it as coincidence, and the idea did not gain wide acceptance.

De Chancourtois' Cylindrical Table (1862)

The French geologist A.E.B. de Chancourtois took a different approach. He arranged the known elements in order of increasing atomic weight and wrapped this list around a cylinder. Elements that lined up vertically on the cylinder showed similar properties. This was the first attempt to display periodic recurrence of properties in a visual form.

Note

De Chancourtois' work was published in a geological journal and did not reach many chemists. Like Dobereiner's triads, it attracted little attention at the time.

Newlands' Law of Octaves (1865)

The English chemist John Alexander Newlands made a more systematic attempt. He arranged the elements in increasing order of atomic weight and noticed a striking pattern: every eighth element resembled the first, much like the eighth note in a musical octave.

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

Lithium (first element) and sodium (eighth element) are both highly reactive metals. Beryllium (second) and magnesium (ninth) are both alkaline earth metals. The pattern held — but only up to calcium.

Watch out

Newlands' Law of Octaves broke down for elements beyond calcium. When new elements were discovered later, they did not fit into his neat octaves. The scientific community ridiculed his idea, comparing it to arranging elements alphabetically. However, in 1887, the Royal Society of London awarded him the Davy Medal in recognition of his pioneering work.

Mendeleev and Lothar Meyer (1869)

The Periodic Law as we know it today owes its development to two chemists working independently: the Russian Dmitri Mendeleev (1834–1907) and the German Lothar Meyer (1830–1895).

Lothar Meyer plotted physical properties — atomic volume, melting point, boiling point — against atomic weight. He obtained a periodically repeated pattern, but unlike Newlands, he noticed that the length of the repeating pattern changed. By 1868, he had developed a table of elements that closely resembles the Modern Periodic Table. However, his work was not published until after Mendeleev's.

Important

Mendeleev is generally credited with the development of the Modern Periodic Table because he published his work first and, more importantly, because he made bold quantitative predictions that were later verified.

Mendeleev's Periodic Law

Mendeleev stated the Periodic Law for the first time:

\text{The properties of the elements are a periodic function of their atomic weights.}

He arranged elements in horizontal rows (which he called series) and vertical columns (groups) in order of increasing atomic weight, such that elements with similar properties occupied the same vertical column.

Key Features of Mendeleev's System

1. He used a broader range of properties. Unlike his predecessors, Mendeleev relied not just on atomic weights but also on the empirical formulas and properties of compounds formed by the elements. For example, he looked at the formulas of oxides and hydrides to decide where an element belonged.

2. He ignored atomic weight when necessary. Mendeleev realized that if he strictly followed atomic weight order, some elements would end up in groups where they did not fit chemically. He made a crucial decision: he placed elements with similar properties together, even if it meant reversing the atomic weight order.

Note

The most famous example is tellurium (atomic weight 127.6) and iodine (atomic weight 126.9). Tellurium has a higher atomic weight than iodine, but tellurium resembles the Group VI elements (oxygen, sulfur, selenium) while iodine resembles the Group VII elements (fluorine, chlorine, bromine). Mendeleev placed tellurium in Group VI and iodine in Group VII, assuming the atomic weight measurements were incorrect. He was right about the grouping — and later, Moseley showed that atomic number, not atomic weight, is the fundamental property.

3. He left gaps for undiscovered elements. Mendeleev was so confident in his system that he left blank spaces in his table where no known element fit. He predicted that these gaps would be filled by elements yet to be discovered.

Important

Mendeleev predicted the existence and properties of two such elements:

Eka-aluminium (the element under aluminium) — later discovered as gallium (1875)
Eka-silicon (the element under silicon) — later discovered as germanium (1886)

Mendeleev's Predictions vs. Experimental Findings

The table below shows how accurately Mendeleev predicted the properties of these elements:

Property	Eka-aluminium (predicted)	Gallium (found)	Eka-silicon (predicted)	Germanium (found)
Atomic 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	$E_2O_3$	$Ga_2O_3$	$EO_2$	$GeO_2$
Formula of chloride	$ECl_3$	$GaCl_3$	$ECl_4$	$GeCl_4$

Tip

The prefix "eka-" comes from Sanskrit, meaning "one" — so eka-aluminium means "one place below aluminium." This naming convention is still used for undiscovered elements in some contexts.

The success of these predictions — especially the quantitative ones — made Mendeleev and his Periodic Table famous. His 1905 version of the table is shown in Fig. 3.1 of the textbook.

The Modern Periodic Law (Moseley, 1913)

When Mendeleev developed his table, nothing was known about the internal structure of the atom. The beginning of the 20th century changed that.

The English physicist Henry Moseley studied the characteristic X-ray spectra of elements. He found that when he plotted the square root of the frequency of X-rays emitted ( $\sqrt{\nu}$ ) against the atomic number ( $Z$ ), he got a straight line. When he plotted the same against atomic mass, he did not.

\sqrt{\nu} \propto Z

This showed that atomic number — the number of protons in the nucleus — is a more fundamental property of an element than its atomic mass. Mendeleev's Periodic Law was therefore modified.

\text{The physical and chemical properties of the elements are periodic functions of their atomic numbers.}

This is the Modern Periodic Law. …

Table 3.1Dobereiner's Triads

Table 3.2Newlands' Octaves

Element	Li	Be	B	C	N	O	F
At. wt.	7	9	11	12	14	16	19

Figure 3.1Mendeleev's Periodic Table published earlier.

Fig. 3.1 — Mendeleev's Periodic Table published earlier.

Drawn by us to help you understand the concept clearly, and verified to make sure it's accurate. For exams, practice from your NCERT textbook's own diagram.

Fig. 3.1 is a reproduction of Mendeleev’s original periodic table from 1869. It is not the modern 18-column table you are used to. Instead, it is a grid of horizontal rows (which Mendeleev called series) and vertical columns (called groups, numbered I through VIII). The elements are arranged in order of increasing atomic weight from left to right across each series, and from top to bottom down the table.

The most striking feature of this figure is the gaps. Mendeleev deliberately left blank spaces in the table — for example, under aluminium (Group III) and under silicon (Group IV). He predicted that these gaps corresponded to undiscovered elements, which he named eka-aluminium and eka-silicon (the prefix eka- means “one” in Sanskrit, indicating the element one place below the known one). The figure shows these gaps as dashes or the notation ‘—’. This was a bold and revolutionary idea: the periodic table was not just a list of known facts, but a predictive tool.

At the top of each group column, you will see general formulas for oxides and hydrides — for instance, $R_2O$ for Group I, $RO$ for Group II, $R_2O_3$ for Group III, $RO_2$ for Group IV, $R_2O_5$ for Group V, $RO_3$ for Group VI, $R_2O_7$ for Group VII, and $RO_4$ for Group VIII. These formulas were Mendeleev’s way of encoding the combining capacity (what we now call valence) of the elements in each group. The fact that elements in the same column formed compounds with the same empirical formula was a powerful clue that they shared similar chemical behaviour. …

Table 3.3Mendeleev's Predictions for the Elements Eka-aluminium (Gallium) and Eka-silicon (Germanium)

Property	Eka-aluminium (predicted)	Gallium (found)	Eka-silicon (predicted)	Germanium (found)
Atomic weight	68	70	72	72.6

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Chemistry · Ch 3 — Classification of Elements and Periodicity in Properties

Genesis of Periodic Classification

3.2