Understanding Ionisation Energy and Ionisation Potential

Ionisation energy and ionisation potential are important concepts describing the minimum energy required to remove an electron from a gaseous atom or ion. These topics help in understanding atomic structure, periodic trends, and chemical reactivity, which are essential for JEE Main preparation.

Definition of Ionisation Energy

Ionisation energy is defined as the minimum energy needed to remove the most loosely bound electron from a neutral gaseous atom, converting it into a positive ion. This process is always endothermic, as energy must be supplied for electron removal.

Ionisation energy is expressed in units of kJ/mol or electron volts (eV). For example, the first ionisation energy is the energy required to remove the outermost electron from a neutral atom.

Metals generally have low ionisation energies due to their tendency to lose electrons, while non-metals exhibit higher ionisation energies because they hold on to their electrons more tightly. Find further explanation in the Ionization Energy Explained page.

Factors Affecting Ionisation Energy

Several atomic properties influence ionisation energy. These factors determine the nuclear attraction on the outermost electron and the ease of its removal.

• Atomic radius: Larger atomic radius reduces nuclear attraction.
• Effective nuclear charge: Higher charge increases attraction on electrons.
• Electron shielding: Inner electrons decrease the effective pull on outer electrons.
• Electron configuration: Stable or half-filled subshells require more energy.

Bohr’s Atomic Model and Ionisation Energy

Bohr’s model describes electrons revolving in discrete orbits with specific energies. The removal of an electron requires it to absorb energy equal to the difference between its present energy state and the energy at infinity (where the electron is free from nuclear attraction).

The energy of an electron in the nth orbit according to Bohr’s theory is:

$E_{n} = -\dfrac{2\pi^2 m e^4}{(4\pi \varepsilon_{0})^2 h^2} \cdot \dfrac{Z^2}{n^2}$

where $E_n$ is the energy, $m$ is the mass of electron, $e$ is its charge, $h$ is Planck’s constant, $Z$ is atomic number, $n$ is principal quantum number, and $\varepsilon_0$ is permittivity of free space.

Ionisation Energy Formula Derivation

Ionisation energy is calculated as the difference between the energy of an electron in its initial orbit ($n_1$) and when it is removed ($n_2 = \infty$):

$E_{n_1} = -R_H \dfrac{Z^2}{n_1^2}$, $E_{n_2} = -R_H \dfrac{Z^2}{\infty^2} = 0$

$\Delta E = E_{n_2} - E_{n_1} = 0 - \left(-R_H \dfrac{Z^2}{n_1^2}\right) = R_H \dfrac{Z^2}{n_1^2}$

Here, $R_H$ is the Rydberg constant with a value of $2.18 \times 10^{-18}$ J/atom. More details on atomic energy levels are available at Understanding Atoms And Nuclei.

Trends of Ionisation Energy in the Periodic Table

Across a period, ionisation energy increases from left to right, primarily due to a decreasing atomic radius and increasing effective nuclear charge. Down a group, ionisation energy decreases, since the atomic radius increases and outer electrons experience greater shielding.

Concept of Ionisation Potential

Ionisation potential is defined as the potential difference required to remove the outermost electron from a gaseous atom, forming a unipositive cation. It is numerically equal to ionisation energy expressed in electron volts per atom.

Both ionisation energy and ionisation potential represent the same physical process and, for single electrons, the terms are often used interchangeably. Their relation is given by $V = \dfrac{E}{e}$, where $E$ is energy in joules and $e$ is the elementary charge.

The difference between ionisation energy and ionisation potential mainly lies in their units: ionisation energy in kJ/mol or eV/atom, ionisation potential in volts. Further distinction is covered at Ionization Energy And Potential Overview.

Ionisation Potential Formula

The formula for ionisation potential is

$V = \dfrac{E}{e} = 13.6\dfrac{Z^2}{n^2} \; \text{V}$

For a hydrogen atom ($Z=1$, $n=1$), $V = 13.6$ V. This value corresponds to the energy needed to remove the single electron from a hydrogen atom in its ground state.

Differences between Ionisation Energy and Ionisation Potential

Relation to Ionisation Enthalpy

Ionisation enthalpy is the enthalpy change when one mole of electrons is removed from one mole of gaseous atoms. For individual atoms, ionisation enthalpy, ionisation energy, and ionisation potential convey closely related concepts, but units differ.

For advanced problems or trends, refer to Overview Of Electrostatics.

Successive Ionisation Energies

After the first ionisation, removing additional electrons requires more energy due to increased positive charge and greater attraction between the nucleus and the remaining electrons. Thus, second and higher ionisation energies are always greater than the first.

Applications and Importance in Chemistry

Knowledge of ionisation energy and potential is essential for predicting the chemical reactivity of elements, the formation of ions, and understanding bonding in molecules. These concepts are also important in the study of nuclear reactions as explained at Nuclear Fission And Fusion Explained.

Ionisation energy data help in explaining stability of electron configurations, abnormal trends, and the nature of metallic and non-metallic character.

Summary of Key Points

• Ionisation energy is the energy to remove the outermost electron.
• Ionisation potential is the corresponding potential difference.
• Both reflect atomic structure and periodic properties.
• Trends in the periodic table are governed by atomic size and nuclear charge.
• Successive ionisation energies always increase for a given atom.
• Applications extend to atomic models, bonding, and chemical reactivity.

For more on kinetic theory involving ionisation in gases, refer to Kinetic Theory Of Gases Overview.