Q.Explain the following:

(a) Electronegativity increases left to right

Concept: Electronegativity measures an atom's ability to attract shared electrons in a bond. Across a period, nuclear charge increases while shielding remains nearly constant.

Moving left to right, each element adds one proton to the nucleus and one electron to the same valence shell. The effective nuclear charge $Z_{\text{eff}}$ felt by valence electrons increases because:

1. Nuclear charge rises by +1 each step.
2. Added electrons enter the same shell, providing negligible additional shielding.
3. Valence electrons are pulled more strongly toward the nucleus. …

(a) Why electronegativity increases across a period

Electronegativity measures an atom's ability to attract shared electrons in a chemical bond. To understand the left-to-right trend, picture what changes as you move across a period.

1. Nuclear charge climbs steadily.

Each step to the right adds one proton to the nucleus. Sodium has 11 protons, magnesium 12, aluminium 13, and so on across period 3. The positive charge pulling on electrons grows stronger.

2. Electrons enter the same principal shell.

All elements in a given period fill the same outermost shell. Across period 3, for instance, every atom adds electrons to the $n=3$ shell. Because the shell doesn't change, the distance between the nucleus and the valence electrons stays roughly constant—or even shrinks slightly.

3. Shielding remains nearly unchanged.

The inner, filled shells screen the valence electrons from the full nuclear charge. As you move across a period, you add electrons to the same outer shell, not to new inner shells. Electrons in the same shell shield one another poorly. The inner-shell shielding therefore stays almost the same.

4. Effective nuclear charge rises.

The net effect is that valence electrons feel a stronger and stronger pull. The nucleus gains charge faster than shielding can compensate. An atom on the right of the period grips its own electrons—and any electrons it shares in a bond—much more tightly than an atom on the left.

5. Atomic radius shrinks.

The stronger effective nuclear charge pulls the electron cloud inward. Chlorine is noticeably smaller than sodium, even though both have three shells. A smaller atom brings the bonding electrons closer to its nucleus, amplifying the attractive force.

Put together, a rising nuclear charge, constant shielding, and a shrinking radius all conspire to make atoms more electronegative as you move left to right.

(b) Why ionisation enthalpy decreases down a group

Ionisation enthalpy is the energy needed to remove the outermost electron from a gaseous atom. The top-to-bottom trend in a group reflects two competing effects—distance and shielding—that both work in the same direction.

1. The outermost electron moves to a higher shell.

Lithium's valence electron sits in the $n=2$ shell, sodium's in $n=3$, potassium's in $n=4$, and so on down Group 1. Each step down adds a complete new shell. The valence electron is now farther from the nucleus.

2. Electrostatic force weakens with distance.

Coulomb's law tells us that the force between two charges falls off as $1/r^2$. An electron in the fourth shell experiences a much weaker pull than one in the second shell, even if the nuclear charge is higher.

3. Inner shells provide additional shielding. …