Classification of Elements and Periodicity

Imagine walking into a big kirana shop with 118 different items scattered everywhere — it would be impossible to remember. That was the same problem chemists faced with 118 elements. They needed shelves and groups to organise them, so early scientists tried to find patterns.

Dobereiner's Triads (1817): He noticed groups of three elements with similar properties where the atomic mass of the middle element was almost the average of the other two. For example, Li (7), Na (23), K (39) — Na is the average of Li and K. Same for Ca (40), Sr (87.5), Ba (137) and Cl (35.5), Br (80), I (127). The idea was clever but only a few triads could be found, so it failed.

Newlands' Law of Octaves (1865): He arranged 56 known elements in increasing atomic mass and found that every 8th element had similar properties, just like the 8th note in music Sa Re Ga Ma... Li, Na and K behaved similarly. But this only worked till Calcium. After that it broke down, and noble gases discovered later had no place at all.

Mendeleev's Periodic Table (1869): With 63 elements, he gave the first real periodic law — properties of elements are a periodic function of their atomic masses. His genius was leaving gaps for undiscovered elements and predicting them correctly — eka-Aluminium became Gallium (68) and eka-Silicon became Germanium (72). He also corrected doubtful atomic masses. But there were problems — hydrogen had no fixed position, isotopes would need same place, and Argon (39.9) came before Potassium (39.1) which broke the mass order.

By 1913, Henry Moseley showed through X-ray experiments that atomic number (Z), not atomic mass, is the real identity of an element. This gave us the Modern Periodic Law — properties of elements are a periodic function of their atomic numbers. This single change fixed the isotope problem, because isotopes have same Z.

Today we use the Long Form Periodic Table with 7 periods and 18 groups. The number of elements in each period is 2, 8, 8, 18, 18, 32, 32. Period 1 has only H and He (2), Periods 2 and 3 have 8 each, Periods 4 and 5 have 18 each, and Periods 6 and 7 have 32 each — including the lanthanides and actinides kept separately at the bottom. Group 1 is alkali metals, Group 2 alkaline earth metals, Groups 13-18 are p-block, Groups 3-12 are d-block, and the bottom two rows are f-block with 14 lanthanides (58-71) and 14 actinides (90-103).

Think of it like apartment floors — each floor (period) has a fixed number of flats, and families with same outer electronic configuration live in the same vertical column (group), so they behave similarly.

The periodic table is also divided by where the last electron enters — we call these blocks, like different sections in a library.

s-block (Groups 1 & 2): Last electron in ns orbital — ns1 and ns2. These are alkali metals (+1) and alkaline earth metals (+2), highly reactive because they have very low ionisation enthalpy and love to lose electrons. Soft metals, like ghee that melts quickly.

p-block (Groups 13 to 18): Last electron in np — ns2 np1 to np6. This block has full variety — metals, non-metals, metalloids (B, Si), halogens and noble gases. d-block (Groups 3 to 12): Last electron in (n-1)d orbital, called transition elements — they show variable oxidation states, form coloured compounds, act as catalysts and are often magnetic.

f-block: Kept at the bottom, inner transition elements. Lanthanides (4f, 58-71) and Actinides (5f, 90-103). Because 4f electrons shield poorly, size shrinks slowly from La to Lu — this is lanthanide contraction, which is why Zr and Hf have almost same size.

Atomic radius is simply how big an atom is — the distance from nucleus to outermost electron. But we measure it in three ways. Covalent radius is half the distance between two same atoms bonded together, metallic radius is half the distance between two adjacent metal atoms in a metal lattice, and van der Waals radius is half the distance between two non-bonded neighbours. So naturally, van der Waals > metallic > covalent.

Across a period from left to right, nuclear charge increases but electrons are added in same shell, so the pull increases and size decreases. Down a group, a new shell is added each time, like adding a new floor to a building, so size increases. A special case is lanthanide contraction — the 4f electrons shield very poorly, so from La to Lu size hardly changes, which makes Zr and Hf almost same size despite different periods.

Ionic radius: A cation is always smaller than its parent atom because it loses a shell or the same shell contracts, while an anion is larger because added electrons increase repulsion. For isoelectronic species (same number of electrons), more positive charge means smaller size — for 10 electrons, Na+ (102 pm) > Mg2+ (72 pm) > Al3+ (54 pm).

We need two energies to understand reactivity. Ionisation Enthalpy (IE) is the energy needed to remove the outermost electron from a gaseous atom. Removing the first electron is IE1, second is IE2, and IE1 < IE2 < IE3 because once positive charge forms, it holds remaining electrons tighter.

Trend: Across a period, size decreases and effective nuclear charge (Z*) increases, so IE increases. Down a group, size increases and outer electron is far away, so IE decreases. But there are exceptions that students love. Be (899 kJ/mol) > B (801) because Be has fully filled stable 2s2 configuration which is hard to break, and N (1402) > O (1314) because N has half-filled stable 2p3 configuration.

Electron Gain Enthalpy (EG enthalpy) is energy change when an electron is added — more negative means more energy released and more favourable. Chlorine has the most negative value (-349 kJ/mol), even more than Fluorine (-328) because F is so small that incoming electron feels repulsion. Across period, EG becomes more negative left to right, down the group it becomes less negative. Noble gases have positive EG values — they don't like adding electrons at all.

Electronegativity (EN) is how strongly an atom in a bond pulls the shared electrons towards itself. On Pauling scale, Fluorine is highest at 4.0, followed by Oxygen at 3.5. Trend is same as IE — across period, EN increases because nuclear pull increases and size decreases; down a group, EN decreases because size increases and pull weakens.

Second period elements Li, Be, B behave oddly compared to their own group — called anomalous behaviour. They are very small, have high charge density, and tend to form covalent compounds instead of ionic, much like how a small child holds its toys tighter than a grown-up.

This leads to Diagonal Relationship — elements placed diagonally have similar charge/radius ratio and polarizing power. Pairs are Li-Mg, Be-Al, B-Si. Examples your teacher will always ask: Both Li and Mg form nitrides Li3N and Mg3N2, both carbonates and nitrates decompose on heating, both chlorides are covalent and soluble in organic solvents. Be and Al both dissolve in NaOH to give complexes, both have amphoteric oxides and covalent chloride-bridged structures. B and Si are both metalloids, both form hydrides like B2H6 and SiH4, both have acidic oxides B2O3, SiO2 with network structures.

All periodic trends finally tell us how reactive an element will be. Reactivity depends on IE, EG, and size — how easily it loses or gains electrons.

For s-block, reactivity increases down the group — Li < Na < K < Rb < Cs because IE decreases down, so losing electron becomes easier. That's why Cs reacts explosively with water while Li reacts calmly, just like how lighter oil catches fire faster than heavy oil. For p-block, heavy elements show inert pair effect — the ns2 electrons become reluctant to bond due to poor shielding of d and f electrons, so Tl shows +1 more stable than +3, and Pb shows +2 more stable than +4.

Overall metallic character: Across a period from left to right, metallic character decreases — metals to metalloids to non-metals to noble gases, and down a group metallic character increases. In short, left-bottom corner of periodic table is most metallic like heavy utensils in kitchen, right-top corner is most non-metallic like airy gases. This summary connects all trends together.