Periodic Table Class 11: Complete Chapter Revision for NEET by PW

The Periodic Table provides a systematic arrangement of elements based on their atomic numbers and electronic configurations. To revise this chapter effectively, you need to understand the modern periodic law, classification of elements into different blocks, electronic configuration, and the factors responsible for periodic trends.

PW Periodic Table Class 11 notes help you revise important concepts such as effective nuclear charge, shielding effect, atomic and ionic radii, ionisation energy, electron gain enthalpy and electronegativity. Key exceptions and trends are also included for quick recall while preparing for Class 11 Chemistry and NEET.

Modern Periodic Table and Periodic Law

The modern periodic table classifies elements according to their atomic numbers and recurring chemical properties.

The Modern Periodic Law states that the physical and chemical properties of elements are periodic functions of their atomic numbers.

The periodic table is organised into:

• 7 periods: Horizontal rows corresponding to the highest principal quantum number.

• 18 groups: Vertical columns containing elements with similar outer electronic configurations and chemical properties.

• 4 blocks: s, p, d and f blocks, based on the subshell into which the differentiating electron enters.

• 118 elements: The currently recognised elements.

Periods and Number of Elements

The 1st period is the shortest period, containing only two elements. The sixth and seventh periods can accommodate 32 elements because they include the f-block elements.

Electronic Configuration and Energy Filling

The arrangement of elements in the periodic table is closely related to the Aufbau principle and the (n + l) rule.

The general subshell filling sequence across a period can be represented as: ns → (n − 2)f → (n − 1)d → np

A useful pattern for remembering the subshells filled across periods 1 to 7 is: s → sp → sp → sdp → sdp → sfdp → sfdp

Steps to Write Electronic Configuration

To write the electronic configuration of an element:

1. Identify its atomic number (Z).

2. Identify the preceding noble gas.

3. Write the noble gas in square brackets.

4. Fill the remaining electrons according to the appropriate subshell filling sequence.

Example: Electronic Configuration of Z = 82

For an element with Z = 82:

• Preceding noble gas: Xe (Z = 54)

• Period: 6th

• Remaining electrons: 82 − 54 = 28

• Electronic configuration: [Xe] 4f¹⁴ 5d¹⁰ 6s² 6p²

Classification of Elements into Blocks

Elements are divided into four blocks according to the subshell receiving the differentiating electron.

The s-block and p-block elements are collectively known as representative or main-group elements.

Lanthanoids and Actinoids

• Lanthanoids: Z = 58 to 71, involving the 4f series.

• Actinoids: Z = 90 to 103, involving the 5f series.

The d-block elements are commonly called transition elements, while the f-block elements are called inner transition elements.

Shielding Effect and Effective Nuclear Charge

In a multi-electron atom, inner electrons partially shield the outer electrons from the attractive force of the nucleus. This is known as the shielding or screening effect.

The effective nuclear charge experienced by an electron can be represented as:

Zeff = Z − σ

where:

• Z = atomic number

• σ = screening constant

• Zeff = effective nuclear charge

Shielding Power of Subshells

The general order of shielding ability is: s > p > d > f

Poor shielding by inner d- and f-electrons can result in contraction of atomic size.

Periodic Contractions

Atomic and Ionic Radius

Atomic size cannot be measured as a sharp boundary because the electron cloud does not have a definite outer edge. Atomic radius is therefore determined using distances between nuclei in suitable bonded or closely packed structures.

Types of Atomic Radius

1. Covalent radius: Half the internuclear distance between two identical atoms joined by a single covalent bond.

2. Metallic radius: Half the distance between the nuclei of two adjacent metal atoms in a metallic crystal.

3. van der Waals radius: Half the distance between the nuclei of two non-bonded adjacent atoms.

The general order is: van der Waals radius > metallic radius > covalent radius

Ionic Radius

The general relationship is:

Anion radius > Parent atom radius > Cation radius

Anions are larger because the addition of electrons increases electron-electron repulsion, while cations are smaller because electron loss reduces electron-electron repulsion and may remove an entire shell.

Isoelectronic Species

Isoelectronic species have the same number of electrons.

For species containing 10 electrons: C⁴⁻ > N³⁻ > O²⁻ > F⁻

For cations: Na⁺ > Mg²⁺ > Al³⁺ > Si⁴⁺

For isoelectronic species, size generally decreases as the number of protons increases.

Periodic Trends in Atomic Radius

• Across a period: Atomic radius generally decreases from left to right because effective nuclear charge increases.

• Down a group: Atomic radius generally increases because new electron shells are added.

Noble-gas atomic sizes are generally discussed using van der Waals radii.

Group 13 Anomaly

The atomic radius order in Group 13 is:

B < Ga < Al < In < Tl

Gallium is smaller than aluminium because the intervening 3d electrons have poor shielding ability, resulting in transition contraction.

Ionisation Energy

Ionisation energy is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.

Ionisation is an endothermic process.

Successive Ionisation Energies

The successive ionisation energies of an element always increase: IE₁ < IE₂ < IE₃ < ... < IEₙ

The relationship between successive ionisation steps can also be represented as: IE₂(M) = IE₁(M⁺)

and

IE₃(M) = IE₁(M²⁺)

Electron Affinity and Electron Gain Enthalpy

Electron affinity refers to the tendency of an isolated gaseous atom to accept an electron, whereas electron gain enthalpy (ΔegH) represents the enthalpy change when an electron is added to an isolated gaseous atom.

For an exothermic electron-gain process:

ΔegH < 0

and electron affinity is conventionally represented as positive.

Important Trends and Exceptions

The first electron gain enthalpy is generally exothermic, but it is endothermic for atoms with particularly stable configurations, such as:

• N

• Be

• Mg

• Noble gases

The addition of a second electron to an already negatively charged ion is always endothermic because of strong electron-electron repulsion.

Why Does Chlorine Have Higher Electron Affinity Than Fluorine?

Although fluorine is above chlorine in the group, the very compact 2p subshell of fluorine produces greater electron-electron repulsion for the incoming electron. The larger 3p subshell of chlorine accommodates the additional electron with comparatively less repulsion.

Therefore: Cl > F

Other important comparisons include:

• S > O

• P > N

• Si > C

Important Group Orders

Halogens: Cl > F > Br > I > At

Chalcogens: S > Se > Te > Po > O

Electronegativity and Its Applications

Electronegativity is the tendency of an atom in a covalent bond to attract the shared pair of electrons towards itself. It is a relative and dimensionless property.

Electronegativity generally increases with:

• Increasing effective nuclear charge

• Positive oxidation state

• Increasing s-character

The s-character order is: sp > sp² > sp³

Electronegativity Values

Some commonly used Pauling electronegativity values are:

Electronegativity and Acidic Character

The acidic character of oxides generally increases with increasing electronegativity and positive oxidation state of the central element.

• Across a period: Acidic character generally increases.

• Down a group: Basic character generally increases.

Amphoteric Oxides

Amphoteric oxides can react with both acids and bases. Examples include oxides of: Zn, Al, Sb, As, Pb, Sn, Ga and Cr³⁺

Neutral Oxides

Some examples of neutral oxides are: NO, N₂O, CO and H₂O

The Periodic Table chapter connects the arrangement of elements with their electronic configurations and chemical properties. These PW Periodic Table Class 11 notes bring together periodic law, block classification, shielding effect, effective nuclear charge and major periodic trends in one place, while the comparison tables and quick revision points make important exceptions and trends easier to recall for Class 11 Chemistry and NEET preparation.

The Periodic Table chapter explains how elements are arranged and how their properties change across periods and groups. These PW Periodic Table notes bring together important concepts for focused revision. Revising these trends along with their important exceptions can help in solving conceptual and application-based questions in Class 11 and NEET.