Classification of Elements

1. Introduction

By the nineteenth century, chemists had discovered many elements and needed a systematic way to organize them so that their properties could be studied and predicted. Early attempts included Dobereiner's triads and Newlands' law of octaves, but the major step was Mendeleev's periodic table (1869), which arranged elements by increasing atomic mass and grouped those with similar properties. Mendeleev even left gaps for undiscovered elements and correctly predicted their properties.

Mendeleev's table had defects (e.g. the position of isotopes, and pairs like $\ce{Ar}$ and $\ce{K}$ that appeared out of order by mass). These were resolved by Moseley's discovery that the fundamental property of an element is its atomic number, not its atomic mass. This gave the modern periodic law.

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2. Modern Periodic Law and the Modern Periodic Table

Modern Periodic Law (Moseley): The physical and chemical properties of elements are a periodic function of their atomic numbers.

The modern (long-form) periodic table arranges the elements in order of increasing atomic number. It consists of:

• 7 horizontal rows called periods, numbered 1 to 7. The period number equals the principal quantum number ($n$) of the outermost shell being filled.
• 18 vertical columns called groups, numbered 1 to 18. Elements in the same group have the same number of valence electrons, hence similar chemical properties.

2.1 Classification into blocks

Elements are grouped into blocks according to which sub-shell receives the last electron:

• $s$-block: Groups 1 and 2 (alkali and alkaline earth metals); last electron enters an $s$ orbital.
• $p$-block: Groups 13 to 18 (includes non-metals, metalloids and noble gases); last electron enters a $p$ orbital.
• $d$-block: Groups 3 to 12, the transition elements; last electron enters a $d$ orbital.
• $f$-block: the lanthanides and actinides (inner transition elements), placed separately at the bottom; last electron enters an $f$ orbital.