Q.Discuss and compare the trend in ionisation enthalpy of the elements of group 1 with those of group 17 elements.

Concept: Periodic Trend in Ionisation Enthalpy

Ionisation enthalpy is the energy required to remove the most loosely bound electron from an isolated gaseous atom. It depends on atomic size, nuclear charge, and shielding effect.

Reasoning:

1. Group 1 (Alkali Metals): These elements have a single valence electron far from the nucleus, with strong shielding from inner electrons. As we move down the group, atomic size increases sharply and effective nuclear charge felt by the valence electron decreases. Hence, ionisation enthalpy decreases down the group.
2. Group 17 (Halogens): These elements have high effective nuclear charge and small atomic size. Their valence electrons are closer to the nucleus and experience greater attraction. As we move down the group, atomic size increases, but the increase in nuclear charge is less effectively shielded, so ionisation enthalpy decreases, but the values remain much higher than those of group 1. …

The Core Concept: Why Ionisation Enthalpy Varies

Ionisation enthalpy is the energy required to remove the most loosely bound electron from an isolated gaseous atom. The key factors are:

• Atomic size — larger atoms have electrons farther from the nucleus, easier to remove
• Effective nuclear charge ($Z_{\text{eff}}$) — more protons pulling on the electron makes it harder to remove
• Shielding effect — inner electrons block the nuclear pull, making outer electrons easier to remove

Group 1 and group 17 sit at opposite ends of the periodic table, so their ionisation enthalpies differ dramatically. Let's trace the trends.

Step-by-Step Comparison

1. Trend Down Group 1 (Alkali Metals)

The elements: Li, Na, K, Rb, Cs, Fr.

As we move down:

• Atomic radius increases significantly — each new shell adds distance
• Shielding increases — more inner electrons block the nuclear charge
• $Z_{\text{eff}}$ on the valence electron stays low — the single $ns^1$ electron feels only about +1 effective charge

Result: Ionisation enthalpy decreases steadily down the group.

Why the drop? The outermost electron is farther from the nucleus and better shielded. Less energy is needed to pluck it away.

2. Trend Down Group 17 (Halogens)

The elements: F, Cl, Br, I, At.

Same direction — ionisation enthalpy decreases down the group — but for a slightly different reason:

• Atomic size increases, shielding increases
• But halogens have high $Z_{\text{eff}}$ (they're one electron short of a noble gas configuration)
• The valence electrons are in the $np^5$ subshell, experiencing strong nuclear pull

The decrease is less steep than in group 1 because the high $Z_{\text{eff}}$ partially offsets the size effect.

3. Comparing the Two Groups Side by Side

The stark difference in magnitude is the real story. Even the highest IE in group 1 (Li at 520) is barely half the lowest IE in group 17 (I at 1008).

4. Why Such a Huge Gap Between Groups?

The root cause is electronic configuration:

• Group 1: One electron in the outermost $s$ orbital. Removing it gives a stable noble gas configuration ($ns^2 np^6$). This is energetically favourable — the atom wants to lose that electron. That's why alkali metals are so reactive and have low IE.

• Group 17: Five electrons in the outermost $p$ subshell. Removing one electron leaves an $np^4$ configuration, which is not particularly stable. The atom strongly holds onto its electrons because it's closer to completing its octet by gaining one electron, not losing it. Hence, very high IE. …