Periodic Classification Of Elements (Mendeléev’s)
Making Order Out Of Chaos – Mendeléev’s Periodic Table
After the attempts by Döbereiner and Newlands, the task of classifying elements continued. The most successful early classification was proposed by the Russian chemist, Dmitri Mendeléev. Independently, a German chemist, Lothar Meyer, also worked on a similar classification, but Mendeléev is credited with developing and publishing the periodic table first in 1869.
Mendeléev studied the relationship between the atomic masses of the elements and their physical and chemical properties. He focused on the compounds formed by the elements with oxygen and hydrogen, as these are very reactive elements and form compounds with most other elements. He observed that elements with similar properties appeared at regular intervals when arranged in increasing order of their atomic masses.
This observation led Mendeléev to propose his Periodic Law:
"The properties of elements are a periodic function of their atomic masses."
Based on this law, Mendeléev arranged the then known elements in a table, which is now known as Mendeléev's Periodic Table.
Structure of Mendeléev’s Periodic Table:
Mendeléev arranged the elements in horizontal rows called Periods and vertical columns called Groups. He placed elements with similar properties in the same group. His table had 7 periods and 8 groups (I to VIII).
Groups I to VII were divided into two sub-groups, A and B. Group VIII consisted of three elements in each period (e.g., Fe, Co, Ni in the fourth period) and was placed separately.
Achievements Of Mendeléev’s Periodic Table
Mendeléev's periodic table was a landmark achievement in the history of chemistry. It provided a systematic and logical way to classify elements and had several significant successes:
1. Systematic Study of Elements: By arranging elements into groups and periods, Mendeléev made the study of elements much easier. Elements in the same group had similar properties, so studying the properties of one element in a group allowed chemists to predict the properties of other elements in that same group.
2. Prediction of New Elements and Their Properties: Mendeléev left gaps in his periodic table. He was confident that these gaps corresponded to elements that were yet to be discovered. Based on the properties of the elements surrounding these gaps, he boldly predicted the properties of these undiscovered elements.
For example, he predicted the existence of three elements and named them based on the element above them in the same group, adding the Sanskrit prefix 'Eka' (meaning 'one'):
- Eka-Boron (predicted to be similar to Boron) - later discovered as Scandium (Sc).
- Eka-Aluminium (predicted to be similar to Aluminium) - later discovered as Gallium (Ga). Mendeléev predicted its atomic mass (68 u), density (5.9 g/cm³), melting point (low), and formula of its oxide (EkaAl₂O₃) and chloride (EkaAlCl₃). When Gallium was discovered, its properties matched these predictions remarkably well (atomic mass 69.7 u, density 5.94 g/cm³, melting point 30.2 °C, oxide Ga₂O₃, chloride GaCl₃).
- Eka-Silicon (predicted to be similar to Silicon) - later discovered as Germanium (Ge). Mendeléev predicted its atomic mass (72 u), density (5.5 g/cm³), oxide formula (EkaSi₂O₄), and chloride formula (EkaSiCl₄). Germanium was found to have atomic mass 72.6 u, density 5.35 g/cm³, oxide GeO₂, and chloride GeCl₄. The match was again very close.
These accurate predictions gave immense credibility to Mendeléev's periodic table.
3. Correction of Atomic Masses: The periodic table helped in correcting the atomic masses of some elements. For instance, the atomic mass of Beryllium (Be) was initially determined as 13.5, placing it in a position where its properties did not match the other elements in the group. Based on its properties similar to Aluminium and Silicon, Mendeléev suggested that its atomic mass should be 9, and its valency should be 2. This corrected value (9.012 u) placed Beryllium correctly in Group II alongside Magnesium (Mg).
4. Accommodation of Noble Gases: Noble gases like Helium (He), Neon (Ne), and Argon (Ar) were discovered much later. Their discovery posed a challenge to the existing periodic table. However, since they are very unreactive (inert), they could be easily placed in a new separate group (Group 0) without disturbing the existing arrangement. This showed the robustness of Mendeléev's basic framework.
Limitations Of Mendeléev’s Classification
Despite its great successes, Mendeléev's periodic table was not without its flaws and limitations. These limitations became apparent as more elements were discovered and the understanding of atomic structure advanced:
1. Position of Hydrogen: Hydrogen is a unique element. It shows properties similar to both alkali metals (Group I) and halogens (Group VII). Like alkali metals, it forms positive ions ($$\text{H}^{+}$$) and combines with halogens, oxygen, and sulphur to form compounds with similar formulae (e.g., HCl, $$H_2O$$, $$H_2S$$ similar to NaCl, $$Na_2O$$, $$Na_2S$$). Like halogens, it exists as a diatomic molecule ($$\text{H}_2$$), combines with metals and forms covalent compounds with non-metals. Mendeléev placed Hydrogen in Group I, but its dual nature remained a challenge and could not be explained properly by its position.
2. Position of Isotopes: Isotopes are atoms of the same element having the same atomic number but different atomic masses (e.g., $$^{35}Cl$$ and $$^{37}Cl$$). According to Mendeléev's periodic law, which is based on atomic mass, isotopes should occupy different positions in the table because they have different atomic masses. However, isotopes have the same chemical properties. Placing isotopes of the same element in different positions would contradict the grouping of elements by similar properties. Mendeléev's table could not accommodate isotopes, as their existence was unknown at the time he proposed his table.
3. Anomalous Pair of Elements: In some cases, Mendeléev had to place elements in decreasing order of atomic mass to ensure that elements with similar properties were grouped together. This contradicted his own periodic law. For example:
- Argon (Ar, atomic mass 39.9 u) was placed before Potassium (K, atomic mass 39.1 u). Argon (a noble gas) has properties similar to Neon and Krypton, while Potassium (an alkali metal) is similar to Sodium and Lithium. Placing Ar before K ensured they were in their correct respective groups (Group 0 and Group I).
- Cobalt (Co, atomic mass 58.9 u) was placed before Nickel (Ni, atomic mass 58.7 u). Cobalt was placed with Rhodium and Iridium, while Nickel was placed with Palladium and Platinum, aligning with their properties.
- Tellurium (Te, atomic mass 127.6 u) was placed before Iodine (I, atomic mass 126.9 u). This placed Te in Group VI (with S, Se) and I in Group VII (with F, Cl, Br), matching their properties.
These 'inversions' or 'anomalous pairs' could not be explained by Mendeléev's law based on atomic mass.
4. Position of Lanthanides and Actinides: The 14 elements following Lanthanum (Lanthanides) and the 14 elements following Actinium (Actinides) have very similar chemical properties. In Mendeléev's table, there was no single place to accommodate these series of elements. They were placed together within Group III, which did not accurately reflect their position or the number of elements in these series.
5. Difference in Properties of Elements in Sub-groups: While elements in the main groups (A sub-groups) showed clear similarities, elements in the sub-groups A and B within the same group often had quite different properties. For instance, elements in Group IA (like Li, Na, K) are very different from elements in Group IB (like Cu, Ag, Au) in terms of reactivity and other characteristics.
Despite these limitations, Mendeléev's periodic table was a monumental step. It not only classified existing knowledge but also provided a framework for future discoveries and highlighted the concept of periodicity, which would later be explained by the electronic structure of atoms.