Classification of Elements and Periodicity in Properties: Class 11 NCERT Guide

Understanding the Classification of Elements in the Periodic Table

The Periodic Table arranges elements in order of increasing atomic number. This classification groups elements with similar properties into columns called groups and rows called periods.

• Groups: Vertical columns, elements share the same number of valence electrons.
• Periods: Horizontal rows, elements have the same number of electron shells.

Elements are broadly classified as:

• Representative elements (s- and p-block)
• Transition elements (d-block)
• Inner transition elements (f-block)

For example, sodium (Na) in Group 1 has 1 valence electron, while chlorine (Cl) in Group 17 has 7 valence electrons. This arrangement helps predict element properties and chemical behaviour.

Key Periodic Trends in Atomic and Ionic Radii

Atomic radius is the distance from the nucleus to the outermost electron. It shows clear trends in the Periodic Table:

• Across a period: Atomic radius decreases due to increasing nuclear charge pulling electrons closer.
• Down a group: Atomic radius increases as new electron shells are added.

Ionic radius depends on whether the atom loses or gains electrons:

• Cations (positive ions) are smaller than their atoms due to electron loss.
• Anions (negative ions) are larger due to electron-electron repulsion.

Understanding these trends is essential for predicting bonding and reactivity.

Ionization Enthalpy and Its Variation in the Periodic Table

Ionization enthalpy (IE) is the energy required to remove an electron from a gaseous atom or ion. It indicates how strongly an atom holds its electrons.

• Trend across a period: IE increases due to stronger nuclear attraction.
• Trend down a group: IE decreases because outer electrons are farther and shielded.

Formula:

$$ ext{X (g)} ightarrow ext{X}^+ (g) + e^-$$

Example:

• IE of Li = 520 kJ/mol
• IE of F = 1681 kJ/mol

This trend explains why alkali metals easily lose electrons and halogens tend to gain electrons.

Electron Gain Enthalpy and Electronegativity Explained

Electron gain enthalpy (EGE) is the energy change when an atom gains an electron. A more negative value means the atom releases energy and readily gains electrons.

• Across a period: EGE becomes more negative (except for some anomalies).
• Down a group: EGE becomes less negative due to increased atomic size.

Electronegativity measures an atom's ability to attract electrons in a bond.

• Trend across a period: Electronegativity increases.
• Trend down a group: Electronegativity decreases.

These properties explain why halogens are strong oxidizers and why electronegativity affects bond polarity.

Chemical Reactivity and Valence States in Periodic Trends

Chemical reactivity depends on how easily atoms gain, lose, or share electrons, which relates to their valence electrons.

• Group 1 (alkali metals): Highly reactive, reactivity increases down the group as ionization enthalpy decreases.
• Group 17 (halogens): Reactivity decreases down the group as electron gain enthalpy becomes less negative.

Valence states often equal the number of valence electrons or $8 - ext{valence electrons}$.

Example:

• Oxygen (Group 16) has 6 valence electrons; common valence states are 2 and 6.
• Nitrogen (Group 15) has 5 valence electrons; valence states include 3 and 5.

Some first elements in groups show diagonal relationships, e.g., Aluminium and Beryllium, due to similar size and electronegativity, causing exceptions in trends.

Worked Example: Predicting Trends in Ionization Enthalpy

Problem: Arrange the following elements in order of increasing ionization enthalpy: Na, Mg, Al, Si.

Solution:

• Across Period 3, ionization enthalpy generally increases.
• From data:
• Na: 496 kJ/mol
• Mg: 737 kJ/mol
• Al: 578 kJ/mol
• Si: 786 kJ/mol

Order: Na < Al < Mg < Si

Explanation: Aluminium has a lower IE than magnesium due to electron configuration stability (Al’s electron removed from a 3p orbital, Mg’s from 3s).

This example shows exceptions in periodic trends due to sublevel electron arrangements.