Class 11 Chemistry Notes

Classification of Elements and Periodicity in Properties Class 11 Notes

Complete, exam-ready notes on the classification of elements and periodicity in properties: the modern periodic law, the s, p, d and f blocks, groups and periods, and the periodic trends — atomic and ionic radii, ionisation enthalpy, electron gain enthalpy and electronegativity — written for CBSE, JEE and NEET revision.

Class11SubjectChemistryCoversCBSE · JEE · NEET

Written byDeep Narayan· Science & Mathematics EducatorReviewed byPushpanjali

What is periodic classification in one line?

The periodic table arranges all elements by increasing atomic number into groups and periods so that elements with similar properties fall into the same column.

From Mendeleev to the Modern Periodic Law

Modern periodic law

Mendeleev ordered elements by atomic mass and left gaps for undiscovered elements. Moseley later showed the true ordering principle is atomic number: the properties of elements are a periodic function of their atomic number, Z.

Mendeleev predicted the properties of unknown elements (Ga, Ge, Sc) from gaps in his table.
Moseley's X-ray work fixed the ordering by nuclear charge, removing the atomic-mass anomalies.
The modern table has 118 elements in 18 groups and 7 periods.

Electronic Configuration and the Blocks

Blocks of the table

s:\;ns^{1-2},\quad p:\;ns^2np^{1-6},\quad d:\;(n{-}1)d^{1-10}ns^{1-2},\quad f:\;(n{-}2)f^{1-14}

The block an element belongs to is set by the subshell being filled last. s-block = Groups 1–2, p-block = Groups 13–18, d-block (transition metals) = Groups 3–12, and f-block (lanthanides + actinides) sits outside the main body.

Group = number of valence electrons (or, for d/f, related to the outer configuration).
Period = the principal quantum number of the outermost shell.
Elements with the same valence-shell configuration share similar chemical behaviour within a group.

Atomic and Ionic Radii

Periodic trends in size

Atomic radius decreases across a period because the increasing nuclear charge pulls the same shell closer. It increases down a group as new shells are added.

Cations are smaller than their parent atoms; anions are larger.
Isoelectronic species (same electron count, e.g. Na⁺, Mg²⁺, Al³⁺, O²⁻, F⁻) shrink as charge increases.
d- and f-orbital contraction: the lanthanide contraction makes post-lanthanide elements smaller than otherwise expected.

Ionisation Enthalpy

Ionisation enthalpy

M(g) \to M^+(g) + e^-,\qquad \Delta H = IE_1

The energy to remove one mole of electrons from a mole of gaseous atoms, giving IE₁, IE₂, IE₃. It decreases down a group and increases across a period — with two stable bumps at half-filled (np³) and fully filled (np⁶) configurations.

Across a period IE rises: nuclear charge grows and radius shrinks.
Down a group IE falls: bigger atom, better shielding, outer electron held weakly.
IE₂ > IE₁ > IE₃ always — removing a further electron from a more positive ion needs more energy.
Exceptions: Be (1s²2s²) shows higher IE₁ than B (2s²2p¹); N (2p³, half-filled) beats O (2p⁴).

Electron Gain Enthalpy and Electronegativity

Electron gain enthalpy

X(g) + e^- \to X^-(g),\qquad \Delta H = \Delta_{eg}H

The enthalpy change when a gaseous atom gains an electron. It is generally most negative (most exothermic) close to the top right — chlorine beats fluorine because fluorine's small size causes repulsion in the tightly packed shell.

Electronegativity

\chi \approx \frac{IE + \Delta_{eg}H}{2}\;(\times 0.5\;\text{scale})

Electronegativity measures the ability of an atom in a molecule to attract the shared electron pair. It increases across a period (F is highest, χ = 4.0) and decreases down a group.

Trends in Chemical and Physical Properties

Valence electrons define the combining ratio and the oxides formed: metals give basic oxides, non-metals acidic oxides.
Metallic character decreases across a period and increases down a group.
Reducing power follows ionisation enthalpy; oxidising power follows electron gain enthalpy.
Noble gases are the least reactive — high IE plus a stable closed shell (except Kr and Xe which form compounds with F and O).

Solved Examples

Example: Arrange N, O and F in order of increasing ionisation enthalpy.

Solution: F has the highest IE (small size, high Z), then N, then O. Nitrogen's IE₁ is higher than oxygen's despite being to its left, because the half-filled 2p³ configuration is extra stable.

Example: Which is smaller — Na⁺ or F⁻ — and why?

Solution: Na⁺ and F⁻ are isoelectronic ions (10 electrons each). For the same electron count, the ion with the higher nuclear charge pulls the electrons in tighter, so Na⁺ is smaller than F⁻.

Revision

Key formulas at a glance

Memorise these before attempting numericals — most exam questions hinge on one of them.

Blocks

s\,ns^{1-2}\;|\;p\,ns^2np^{1-6}\;|\;d\,(n{-}1)d^{1-10}ns^{1-2}\;|\;f\,(n{-}2)f^{1-14}

IE definition

M(g) \to M^+(g) + e^-,\; \Delta H = IE_1

Electronegativity (Mulliken)

\chi \propto \frac{IE + \Delta_{eg}H}{2}

Electron gain

X(g) + e^- \to X^-(g)

Isoelectronic size

\text{radius } \propto 1/Z\;(\text{same e}^-\text{ count})

Period rule

\text{period} = n_{(\text{outer shell})}

Group rule

\text{group} = \text{valence electrons}

Exam tips

How this chapter is asked

Where this topic appears in CBSE, JEE Main and NEET papers.

Modern periodic law: properties are a periodic function of atomic number, not atomic mass.
Across a period radius and metallic character fall; IE, egH and electronegativity rise.
Down a group radius rises; IE, egH and electronegativity fall.
Half-filled (np³) and fully filled (np⁶) configurations are extra stable — IE bumps.
F is the most electronegative element (χ = 4.0).
Chlorine has a more negative electron gain enthalpy than fluorine.
Cations smaller than atoms, anions larger; lanthanide contraction shrinks post-La elements.

FAQ

Common questions

What is the modern periodic law?

The properties of elements are a periodic function of their atomic number. The modern table therefore arranges elements by increasing Z into 18 groups and 7 periods, replacing Mendeleev's atomic-mass ordering.

What are the s, p, d and f blocks?

Blocks name the subshell being filled last: s-block (Groups 1–2, ns¹⁻²), p-block (Groups 13–18, ns²np¹⁻⁶), d-block transition metals (Groups 3–12) and f-block lanthanides/actinides.

Why does ionisation enthalpy increase across a period?

Across a period, atomic size shrinks while nuclear charge grows, so the outer electrons are held more strongly. It decreases down a group because added shells shield the outer electron and enlarge the atom.

Why is fluorine's electron gain enthalpy less negative than chlorine's?

Fluorine has a very small size; the incoming electron feels strong repulsion from the tightly packed 2p shell, making the gain less favourable than for chlorine, which is larger with better space for the extra electron.

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