Why this chapter matters

The periodic table helps predict rather than merely recall chemical behaviour. Once electron configuration and effective nuclear attraction are understood, many seemingly separate questions—ionic size, metallic character, oxide nature and halogen reactivity—become one connected set of ideas.

From early classifications to the modern periodic law

Earlier attempts to organise elements included Döbereiner’s triads and Newlands’ law of octaves. Mendeleev arranged elements mainly by atomic masses and recurring properties, leaving spaces that helped predict missing elements. His table could not consistently explain isotopes or certain mass-order anomalies.

The modern periodic law states that physical and chemical properties recur periodically with atomic number. Atomic number determines nuclear charge and, for a neutral atom, electron number. Repetition of outer-shell electronic structures explains much of periodic behaviour.

Groups, periods and the s, p, d and f blocks

The long-form periodic table has 7 periods and 18 groups. A period is a horizontal row; a group is a vertical column. Elements in the same main-group column often share related valence-shell configurations and hence similar chemistry. The s block mainly includes groups 1–2 (with helium’s electron configuration an instructive exception to its group placement); the p block generally spans groups 13–18; d-block elements occupy groups 3–12 in the central table, with transition-metal definition requiring care; f-block series are usually shown separately to save space.

The period number of a typical main-group element matches the highest occupied principal quantum number. Helium’s 1s² configuration is s-like, yet it is placed in group 18 because of its chemical inertness and filled first shell.

Why periodic properties change: shielding and effective nuclear charge

The nucleus attracts electrons electrostatically, but inner electrons partially shield outer electrons from nuclear charge. Effective nuclear charge summarises the net nuclear attraction experienced by an electron in a many-electron atom. Across a period, proton number usually increases while electrons enter broadly the same principal shell; effective attraction on valence electrons generally increases.

Down a group, the valence shell moves farther from the nucleus and additional inner shells screen attraction. These two ideas explain many, but not every, periodic trend. The precise trend can depend on whether one compares covalent radii, ionic radii, measured ionisation energies or other operational definitions.

Atomic radius, ionic radius and isoelectronic comparisons

Atomic radius generally decreases from left to right across a period and increases down a group. Radius requires a measurement definition: covalent, metallic and van der Waals radii are not identical. Noble-gas radii may be tabulated with a different basis and should not be blindly ranked alongside covalent radii.

A cation is usually smaller than its neutral parent atom because it has lost electrons and may have lost its outer shell. An anion is generally larger than its neutral parent atom because electron–electron repulsion is increased for the same nuclear charge.

When candidates compare Na⁺ and Mg²⁺, counting shells is not enough; both have a neon-like electron arrangement, but Mg²⁺ has 12 protons compared with 11 in Na⁺, making Mg²⁺ smaller.

Ionisation enthalpy and its exam-favourite exceptions

First ionisation enthalpy is the energy required to remove an electron from an isolated gaseous atom in its ground state to form a gaseous cation. It is always endothermic in this definition. First ionisation enthalpy generally increases across a period and decreases down a group. Noble gases tend to have high values; alkali metals have relatively low values.

Two standard exceptions in period 2 matter: B has a lower first ionisation enthalpy than Be, because B’s removed electron is from the higher-energy 2p subshell rather than Be’s filled 2s subshell. O has lower first ionisation enthalpy than N, because N has relatively stable half-filled 2p orbitals and O has a paired 2p electron that is easier to remove.

Electron gain enthalpy versus electronegativity

Electron gain enthalpy describes the energy change when an isolated gaseous atom gains an electron. A negative value means energy is released. Across a period the first electron gain enthalpy often becomes more negative, while noble gases generally do not readily accept an extra electron. The exact order includes notable exceptions: chlorine has a more negative first electron gain enthalpy than fluorine because adding an electron to fluorine’s compact 2p region introduces strong repulsion.

Electronegativity instead describes an atom’s tendency to attract the shared electron pair within a chemical bond; it is a relative scale, not simply the negative of electron gain enthalpy. Fluorine is the most electronegative element on the usual Pauling scale, even though chlorine’s first electron gain enthalpy is more exothermic.

Metallic character, non-metallic character and oxide behaviour

Metallic character—the tendency to lose electrons and display metallic properties—typically decreases across a period and increases down a group. Non-metallic character tends to show the reverse general pattern. Metal oxides are often basic, non-metal oxides often acidic, but several important oxides are amphoteric or neutral.

In period 3, oxides illustrate a broad shift from strongly basic at the metallic end to acidic toward the non-metallic end, with Al₂O₃ occupying an amphoteric position. Avoid assuming every metal oxide is basic or every non-metal oxide reacts directly with water to form an acid.

Reactivity patterns and special placements

Among alkali metals, reactivity with water generally becomes more vigorous down the group: potassium is more reactive than sodium under comparable conditions. Down group 17, halogens generally become weaker oxidising agents; fluorine is the strongest oxidising halogen under standard aqueous comparisons. These trends arise from different processes, so do not treat ‘reactivity increases down a group’ as a universal rule.

Hydrogen shares some features with both group 1 and group 17 elements, but it is chemically distinctive and not an alkali metal. Helium’s place in group 18 reflects its closed-shell, low-reactivity behaviour even though its electron configuration is 1s².

One-minute recall sheet

• Modern periodic law is based on atomic number, not atomic mass.
• There are 7 periods and 18 groups in the standard long-form table.
• Atomic radius typically decreases across a period and increases down a group.
• In an isoelectronic series, ionic radius decreases as nuclear charge increases.
• First ionisation enthalpy trends have classic Be/B and N/O exceptions.
• Chlorine has more negative first electron gain enthalpy than fluorine; fluorine is more electronegative.
• Al₂O₃ is amphoteric; CO is a neutral oxide.
• Alkali-metal reactivity generally increases down the group, while halogen oxidising strength generally decreases.

UPSC-style practice — 15 explained questions

These are original practice questions, not claims of verbatim UPSC previous-year questions. Attempt each before opening the answer. Questions are single-correct unless specified.

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