Question

The ionisation energy of nitrogen is more than oxygen because (1) more attraction of electrons by the nucleus (2) the extra stability of half-filled \(p\) -orbitals (3) the size of nitrogen atom is smaller (4) more penetrating effect

Short answer

The ionization energy of nitrogen is higher due to the extra stability of half-filled p-orbitals.

Step-by-step solution

Understanding Ionization Energy

Ionization energy is the energy required to remove an electron from an atom in its gaseous phase. Factors such as nuclear charge, atomic size, electron configuration, and subshell stability can affect ionization energy.

Comparing Electron Configurations

Nitrogen has an electron configuration of \([He] 2s^2 2p^3\), while oxygen has \([He] 2s^2 2p^4\). Focus on the p-orbital configuration where nitrogen has half-filled \(2p^3\) orbitals and oxygen has \(2p^4\) orbitals.

Stability of Half-filled Orbitals

Half-filled orbitals (2p^3 in nitrogen) are more stable because of symmetric distribution of electrons, leading to less repulsion between them. This stability makes it harder to remove an electron from nitrogen compared to an element like oxygen.

Identifying the Correct Reason

Given these factors, the higher ionization energy of nitrogen is primarily due to the extra stability of its half-filled \(2p^3\) orbitals.

Key concepts

Electron Configuration

Understanding electron configuration is crucial when discussing ionization energy. Electron configuration refers to the distribution of electrons in an atom's orbitals. For instance, nitrogen has an electron configuration of \( [He] 2s^2 2p^3 \). Oxygen, on the other hand, has \( [He] 2s^2 2p^4 \). The key difference here lies in the filling of the \(2p\) orbitals.

• Nitrogen has three electrons in its \(2p\) orbitals, making a \(2p^3\) configuration.
• Oxygen has four electrons in its \(2p\) orbitals, making a \(2p^4\) configuration.

The arrangement of these electrons impacts their stability and the energy needed to remove them (ionization energy). Therefore, grasping how electrons occupy different orbitals helps in understanding why nitrogen requires more energy for ionization compared to oxygen.

Half-filled p-orbitals Stability

The stability of half-filled orbitals is another important factor affecting ionization energy. Half-filled orbitals, like the \(2p^3\) in nitrogen, exhibit exceptional stability due to their symmetric electron distribution. This symmetric distribution leads to reduced electron-electron repulsion.

Recall that nitrogen's configuration of \(2p^3\) means each \(2p\) orbital has one electron. This arrangement provides extra stability to the atom:

• The electrons are equally spaced.
• There is less repulsion among them.

In contrast, oxygen's \(2p^4\) configuration means at least one orbital has two electrons, causing greater electron-electron repulsion and reducing overall stability. As a result, removing an electron (ionization) from nitrogen needs more energy because breaking out of its stable configuration is harder.

Nuclear Charge Effect

The effect of nuclear charge also plays a significant role in determining ionization energy. Nuclear charge refers to the total positive charge of the nucleus, which is essentially the number of protons in an atom. This charge impacts how tightly electrons are held.

For nitrogen and oxygen, consider the following points:

• Nitrogen has a nuclear charge of +7 (seven protons).
• Oxygen has a nuclear charge of +8 (eight protons).

Although oxygen has a higher nuclear charge, the additional stability provided by nitrogen's half-filled \(2p^3\) orbitals outweighs the increased attraction due to a higher nuclear charge. Therefore, even though positively charged oxygen attracts its electrons more strongly, the inherent stability of nitrogen's electron configuration requires more energy to disrupt.

Understanding these concepts helps explain why nitrogen has a higher ionization energy compared to oxygen despite its lower nuclear charge.