Les énergies de première et de deuxième ionisation du néon...

Les énergies de première et de deuxième ionisation du néon $(Z=10)$, du sodium ( $Z=11)$ et du magnésium $(Z=12)$ sont les suivantes (en kilojoules par mole) : $$ \begin{array}{lccc} & \mathrm{Ne} & \mathrm{Na} & \mathrm{Mg} \\ 1^{\text {re } \text { ionisation }} & 2081 & 496 & 738 \\ 2^{\mathrm{e}} \text { ionisation } & 3952 & 4562 & 1451 \end{array} $$ On remarque que le sodium a la plus faible énergie de première ionisation et la plus forte énergie de deuxième ionisation. Comment cela peut-il s'expliquer, pour trois éléments qui se suivent dans la classification périodique?

Les énergies de première et de deuxième ionisation du néon $(Z=10)$, du sodium ( $Z=11)$ et du magnésium $(Z=12)$ sont les suivantes (en kilojoules par mole) :
$$
\begin{array}{lccc}
& \mathrm{Ne} & \mathrm{Na} & \mathrm{Mg} \\
1^{\text {re } \text { ionisation }} & 2081 & 496 & 738 \\
2^{\mathrm{e}} \text { ionisation } & 3952 & 4562 & 1451
\end{array}
$$
On remarque que le sodium a la plus faible énergie de première ionisation et la plus forte énergie de deuxième ionisation. Comment cela peut-il s'expliquer, pour trois éléments qui se suivent dans la classification périodique?

Key Concepts

Effective Nuclear Charge and Electron Shielding

Effective nuclear charge is the net positive charge experienced by an electron in an atom, and electron shielding refers to the reduction in attraction between the nucleus and an electron due to the presence of other electrons. In atoms with a single loosely bound electron, such as those with an outer s-orbital configuration, the first ionisation energy is low because the outer electron is well shielded. However, once this electron is removed, the remaining electrons experience a higher effective nuclear charge, resulting in a significant increase in the energy required for subsequent ionisations.

Periodic Trends

Periodic trends refer to the systematic variation in properties such as ionisation energy, atomic radius, and electron affinity as one moves across a period or down a group in the periodic table. These trends arise due to changes in both nuclear charge and electron-electron interactions. The example under consideration illustrates how a slight change in atomic number and electron filling can lead to significant differences in first and second ionisation energies.

Ionisation Energy

Ionisation energy refers to the amount of energy required to remove an electron from a neutral atom. It is a key property that provides insights into the electronic structure and reactivity of elements. The trend in ionisation energies across the periodic table helps explain variations in chemical behavior, as differences in electron binding energies can be traced to effective nuclear charge and electron shielding effects.

Electronic Configuration

Electronic configuration describes the distribution of electrons among the various orbitals in an atom. It is pivotal for understanding chemical properties because the arrangement of electrons influences an atom’s stability, reactivity, and the energy needed to remove an electron. In the context of ionisation energies, configurations that result in a complete or half-complete subshell are particularly stable, causing larger energy jumps when these electrons are removed.

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