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What is ionisation enthalpy, and how does it vary in the periodic table?

Class 11 ChemistryNCERT chapter: Classification of Elements and Periodicity in Propertiesइसे हिंदी में पढ़ें

Exam-ready answer

Ionisation enthalpy is the energy required to remove an electron from an isolated gaseous atom in its ground state: X(g) → X⁺(g) + e⁻. It is expressed in kJ mol⁻¹ and is always positive. It generally increases from left to right across a period, because nuclear charge increases and atomic size decreases, and decreases down a group, because size and shielding increase.

Explanation

Ionisation enthalpy is a measure of how easily an atom loses an electron. To remove an electron, energy must be supplied to overcome the attraction of the nucleus, so the process is endothermic. The energy for removing the first electron is the first ionisation enthalpy (ΔᵢH₁). The second ionisation enthalpy is always higher, because the electron is now removed from a positive ion, which holds its remaining electrons more strongly.

Across a period, electrons are added to the same shell while the nuclear charge increases by one each step. The inner electrons shield the outer ones only partly, so the effective nuclear charge rises and the atom becomes smaller. The outer electrons are held more tightly and ionisation enthalpy increases. That is why alkali metals have the lowest values and noble gases the highest values in each period.

Down a group, a new shell is added each time. The outermost electron is farther from the nucleus and better shielded by inner electrons, so it is removed more easily and ionisation enthalpy decreases. There are two well-known exceptions in the second period: boron is lower than beryllium because its electron comes from a 2p orbital, and oxygen is lower than nitrogen because removing a paired 2p electron from oxygen relieves electron repulsion, while nitrogen has a stable half-filled 2p³ configuration.

Points to remember

  1. Ionisation enthalpy: energy required to remove an electron from an isolated gaseous atom in its ground state; unit kJ mol⁻¹.
  2. Process: X(g) → X⁺(g) + e⁻; it is always endothermic (positive).
  3. Across a period it generally increases (higher effective nuclear charge, smaller size); down a group it decreases (larger size, more shielding).
  4. Second ionisation enthalpy is higher than the first: ΔᵢH₂ > ΔᵢH₁.
  5. Exceptions: B < Be (2p electron easier to remove than 2s) and O < N (N has stable half-filled 2p³).

Exam tip

Always write the equation X(g) → X⁺(g) + e⁻ and stress 'isolated gaseous atom in its ground state'. For the trend, give the reason in terms of effective nuclear charge, atomic size and shielding, and quote the B/Be and O/N exceptions with configurations.

Questions students also ask

Why is the first ionisation enthalpy of nitrogen higher than that of oxygen?
Nitrogen has the configuration 2s²2p³, with three 2p electrons in different orbitals, a stable half-filled arrangement. In oxygen (2s²2p⁴), two electrons share one 2p orbital and repel each other. Removing one of these paired electrons is easier, so oxygen has the lower ionisation enthalpy.
Why do noble gases have the highest ionisation enthalpy in a period?
Noble gases have completely filled valence shells and the highest effective nuclear charge in their period, with the smallest size. Their outer electrons are held very tightly, and a full shell is very stable, so a large amount of energy is needed to remove an electron.
How is ionisation enthalpy related to metallic character?
Metals tend to lose electrons to form positive ions. Elements with low ionisation enthalpy, such as alkali metals, lose electrons easily and are strongly metallic and good reducing agents. As ionisation enthalpy increases across a period, metallic character decreases and non-metallic character increases.

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