Board exam focus — Classification of Elements and Periodicity in Properties (CBSE & HBSE)

This chapter explains how the vast number of elements is organised into the modern periodic table. It covers the genesis of periodic classification, the modern periodic law, the long form of the periodic table and its division into s-, p-, d- and f-blocks. It then develops the periodic trends in atomic and ionic radii, ionisation enthalpy, electron gain enthalpy, electronegativity and valency, and relates these to chemical reactivity. It is a conceptually unifying chapter, important for understanding bonding and the chemistry of elements in both CBSE and HBSE syllabi.

Genesis of Classification, Modern Periodic Law and the Long Form Table

Why Classify Elements?

With over a hundred known elements, classification helps to study their properties systematically and predict the behaviour of unknown elements.

Early Attempts

Mendeleev's success: He predicted eka-aluminium (gallium) and eka-silicon (germanium) and arranged elements so that those with similar properties fell in the same group.

Modern Periodic Law

After Moseley's work on X-ray spectra, the law was revised:

The physical and chemical properties of elements are a periodic function of their atomic numbers.

Atomic number (not atomic mass) is the true basis of classification, which removed the anomalies of Mendeleev's table.

Long Form of the Periodic Table

• 7 periods (horizontal rows): the period number equals the principal quantum number (n) of the outermost shell.
• 18 groups (vertical columns): elements in a group have similar valence-shell configurations and hence similar properties.

Trap: The period number tells you the value of n for the outermost shell; the number of elements in a period equals the number of electrons needed to fill the orbitals available at that energy level.

s-, p-, d- and f-Blocks and Atomic & Ionic Radii

Classification into Blocks

The periodic table is divided into blocks based on the subshell into which the last electron enters.

• Metals occupy the left and centre; non-metals the upper right; metalloids (B, Si, Ge, As, Sb, Te) lie along the diagonal.
• The noble gases (Group 18) have completely filled valence shells and are chemically inert.

Atomic Radius

The distance from the nucleus to the outermost electron shell. Measured as covalent, metallic or van der Waals radius.

Trends:

• Across a period (left to right): atomic radius decreases because nuclear charge increases while electrons are added to the same shell, pulling them closer.
• Down a group: atomic radius increases because a new shell is added in each period.

Ionic Radius

• Cations are smaller than their parent atoms (loss of electrons, often loss of a shell, greater effective nuclear charge per electron).
• Anions are larger than their parent atoms (added electrons increase repulsion).

Isoelectronic species have the same number of electrons. Among them, the radius decreases as the nuclear charge (atomic number) increases. For example, for N3-, O2-, F-, Na+, Mg2+, Al3+ (all 10 electrons), size decreases in that order.

Ionisation Enthalpy, Electron Gain Enthalpy, Electronegativity & Reactivity

Ionisation Enthalpy (IE)

The energy required to remove the most loosely bound electron from an isolated gaseous atom: M(g) -> M+(g) + e-.

Trends:

• Across a period: IE generally increases (smaller size, higher nuclear charge).
• Down a group: IE decreases (larger size, increased shielding).

Anomalies: Be has higher IE1 than B (stable 2s2), and N has higher IE1 than O (stable half-filled 2p3). Examiners love these two exceptions.

Successive ionisation enthalpies always increase: IE1 < IE2 < IE3.

Electron Gain Enthalpy (Delta_eg H)

The enthalpy change when an electron is added to a neutral gaseous atom: X(g) + e- -> X-(g).

• More negative value -> greater tendency to gain an electron.
• Halogens have the most negative electron gain enthalpies.

Anomaly: The electron gain enthalpy of fluorine is less negative than that of chlorine, because the small size of the F atom causes strong inter-electronic repulsion in the compact 2p subshell. Hence chlorine, not fluorine, has the most negative electron gain enthalpy.

Electronegativity

The tendency of an atom to attract the shared pair of electrons in a covalent bond (Pauling scale).

• Across a period: increases.
• Down a group: decreases.
• Fluorine is the most electronegative element (3.98).

Periodic Trends in Chemical Reactivity

• Metallic character decreases across a period and increases down a group.
• Reactivity of metals increases down Group 1 (easier to lose electrons).
• Reactivity of non-metals increases up Group 17 (easier to gain electrons), so fluorine is the most reactive non-metal.
• The most reactive metal (Cs/Fr) is at the bottom-left; the most reactive non-metal (F) is at the top-right.

Frequently asked questions

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Do these notes follow CBSE and HBSE?

Yes. The Classification of Elements and Periodicity in Properties notes are NCERT-aligned and include guidance for both CBSE and Haryana Board (HBSE), with important questions and MCQs for revision.

What does the Classification of Elements and Periodicity in Properties chapter cover?

Concept explanations, key formulas and definitions, fully solved examples and board-pattern practice questions for Classification of Elements and Periodicity in Properties.