Ionization Energy (original) (raw)

Last Updated : 2 Jun, 2026

Ionization energy is an important periodic property that explains how easily an atom can lose an electron. It is defined as the minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom. When an electron is removed, the atom forms a positively charged ion called a cation.

Ionization energy is usually expressed in kilojoules per mole (kJ/mol).

The ionization process can be represented as:

X (g) → X + (g) + e−

where:

• X(g) represents a gaseous atom
• X⁺(g) represents the positive ion formed
• e⁻ represents the electron removed.

Characteristics of Ionization Energy

The important characteristics of ionization energy are as follows:

• Ionization energy is always positive because energy must be supplied to remove an electron from an atom.
• Ionization energy is measured in kilojoules per mole (kJ/mol).
• Ionization energy depends on the attraction between the nucleus and the outermost electron.
• Elements with low ionization energy lose electrons easily and generally show metallic character.
• Elements with high ionization energy hold their electrons strongly and usually show non-metallic character.
• Successive ionization energies increase progressively because after the removal of each electron, the positive charge on the ion increases and the remaining electrons are held more strongly by the nucleus.

Types of Ionization Energy

An atom can lose more than one electron. Therefore, ionization energy occurs in successive stages depending on the number of electrons removed. Each successive ionization energy is greater than the previous one because after the removal of an electron.

The successive ionization energies follow the order: IE1 < IE2 < IE3

This increase occurs because each electron is removed from a more positively charged species.

**1. First Ionization Energy (IE 1 )

The minimum amount of energy required to remove the outermost electron from an isolated gaseous atom is called the first ionization energy. The first ionization energy is generally the lowest because the electron is removed from a neutral atom.

The process is represented as: X (g) → X + (g) + e−

**Example: Na (g) → Na + (g) + e−

**2. Second Ionization Energy (IE 2 )

The energy required to remove the second electron from a singly charged positive ion is called the second ionization energy. The second ionization energy is always greater than the first ionization energy because the electron is removed from a positively charged ion.

The process is represented as: X+ (g) → X2 + (g) + e−

**Example: Na+ (g) → Na 2+ (g) + e−

**3. Third Ionization Energy (IE 3 )

The energy required to remove the third electron from a doubly charged positive ion is called the third ionization energy. The third ionization energy is greater than the second ionization energy because the attraction between the nucleus and the remaining electrons becomes stronger.

The process is represented as: X2+ (g → X3+ (g) + e−

E**xample: Al 2+ (g) → Al 3+ (g) + e−

Factors Affecting Ionization Energy

Ionization energy depends on the force of attraction between the nucleus and the outermost electron. The important factors affecting ionization energy are:

**1. Atomic Size: Atomic size refers to the distance between the nucleus and the outermost electron.

• When the atomic size increases, the outermost electron moves farther away from the nucleus and experiences less attraction.
• Therefore, less energy is required to remove the electron, and ionization energy decreases.

**2. Nuclear Charge: Nuclear charge is the positive charge present in the nucleus due to protons.

• As nuclear charge increases, the attraction between the nucleus and electrons also increases.
• As a result, more energy is needed to remove an electron, so ionization energy increases.

**3. Shielding Effect: Inner-shell electrons block the attraction between the nucleus and the outermost electron. This reduction in attraction is called the shielding effect or screening effect.

• When shielding increases, the outermost electron experiences less effective nuclear attraction and can be removed more easily.
• Therefore, ionization energy decreases with increase in shielding effect.

**4. Electronic Configuration: Atoms having completely filled or half-filled orbitals are more stable.

• Due to this extra stability, more energy is required to remove an electron from such atoms.
• Therefore, atoms with stable electronic configurations generally have higher ionization energy.

Trends in Periodic Table

Ionization Energy follows a specific trend while moving from Left to Right and Top to Bottom in the periodic table.

**1. Across a Period

Ionization energy generally increases from left to right across a period.

• This is because nuclear charge increases , atomic size decreases , electrons are held more strongly by the nucleus.
• Therefore, more energy is required to remove an electron.

**Example: Lithium has lower ionization energy than fluorine because fluorine has a smaller atomic size and stronger nuclear attraction.

**2. Down a Group

Ionization energy generally decreases from top to bottom in a group.

• This happens because atomic size increases , shielding effect increases , outermost electrons are farther from the nucleus
• As a result, electrons can be removed more easily and less energy is required.

**Example: Lithium has higher ionization energy than sodium and potassium because lithium is smaller in size and its outermost electron is more strongly attracted by the nucleus.

Exceptions in Ionization Energy

Although ionization energy generally increases across a period, some elements do not follow this regular trend. These exceptions arise due to differences in electronic configuration and the extra stability of completely filled and half-filled orbitals.

**1. Exception between Beryllium and Boron

According to the general trend, boron should have higher ionization energy than beryllium because ionization energy increases across a period. However, beryllium has higher ionization energy than boron.

• Beryllium has a completely filled 2s orbital, which is more stable.
• In boron, the electron is removed from the 2p orbital, which is higher in energy and less strongly held.
• Therefore, less energy is required to remove an electron from boron.

IE (Be) > IE (B)

**2. Exception between Nitrogen and Oxygen

Generally, ionization energy increases from nitrogen to oxygen, but nitrogen has higher ionization energy than oxygen.

• Nitrogen has a half-filled p orbital, which provides extra stability.
• In oxygen, one of the p orbitals contains a pair of electrons, causing electron-electron repulsion.
• As a result, it is easier to remove an electron from oxygen than from nitrogen.

IE (N) > IE (O)

Ionization Energy vs Ionization Enthalpy

**Related Articles:

• Nuclear Binding Energy
• Molecules and Compounds
• Gram Atomic and Gram Molecular Mass