Q.Among alkali metals which element do you expect to be least electronegative and why?

Concept: Electronegativity decreases down a group in the periodic table.

Electronegativity measures an atom's ability to attract shared electrons in a chemical bond. Within the alkali metals (Group 1), atomic size increases as we move from lithium to francium because each successive element adds a new electron shell.

As the atomic radius grows, the valence electron sits farther from the nucleus and experiences greater shielding from inner electrons. Both effects weaken the nucleus's pull on bonding electrons, reducing electronegativity. …

Why electronegativity decreases down a group

Electronegativity measures an atom's ability to attract shared electrons in a chemical bond. Three structural factors control this attraction:

Atomic radius: The farther the valence shell sits from the nucleus, the weaker the pull on incoming electrons. Distance dilutes electrostatic force by the inverse-square law.

Nuclear charge: More protons mean stronger attraction, but this effect is largely canceled by the next factor.

Shielding: Inner-shell electrons repel valence electrons, screening them from the full nuclear charge. As you add shells moving down a group, each new layer of core electrons weakens the effective nuclear charge felt by the outermost electron.

In the alkali metals (Group 1), every element has the same valence configuration—one $s^1$ electron—but each successive element adds a complete shell. The competition between these factors is one-sided: radius and shielding grow dramatically, while the increase in nuclear charge is modest and mostly offset.

Step-by-step comparison across alkali metals

1. Lithium (Li): Valence electron in the $n=2$ shell, only one inner shell ($1s^2$) providing shielding. The $2s$ electron is relatively close to the nucleus and experiences a fairly strong effective nuclear charge.

2. Sodium (Na): Valence electron now in $n=3$, with two complete inner shells. The atom is larger, and the $3s$ electron is more shielded—electronegativity drops.

3. Potassium (K): Valence electron in $n=4$. Radius increases further, shielding intensifies. Electronegativity continues to fall.

4. Rubidium (Rb): Valence electron in $n=5$. The trend persists—larger size, more shielding, weaker hold on bonding electrons.

5. Cesium (Cs): Valence electron in $n=6$, the outermost position among the common alkali metals. The $6s$ electron is separated from the nucleus by five complete shells. The effective nuclear charge is at its minimum, the atomic radius at its maximum. Cesium holds its valence electron—and any shared electron in a bond—most weakly. …