Q.For alkali metals, which one of the following trends is incorrect ?

Group 1 -- lithium, sodium, potassium, rubidium, caesium and francium -- are the alkali metals: soft, highly reactive metals whose common valence-shell configuration is ns¹ (n = period number).

Because they are so reactive, alkali metals never occur free in nature, only as compounds; francium is intensely radioactive with a longest-lived isotope of just 21 minutes half-life. Being the first element of every period, each alkali metal has the largest atomic and ionic radius in its period, and both radii increase down the group as shells are added. Having only one loosely-held valence electron, alkali metals readily lose it to form M⁺, giving them the lowest first ionisation enthalpy in their period -- and a very high second ionisation enthalpy, since removing a second electron would break up the already noble-gas-like M⁺ configuration. Ionisation enthalpy falls down the group as atomic size and screening increase.

Lithium salts are unusually soluble (e.g. LiClO₄ up to 12× more soluble than NaClO₄) because the very small Li⁺ ion is strongly hydrated; hydration enthalpy magnitude falls steadily down the group (Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺). Electronegativity is comparatively low and falls further down the group, so alkali metals typically form ionic compounds with more electronegative elements.

A distinctive flame test -- heating a moistened salt on a platinum wire -- gives each metal its own characteristic colour (Li crimson red, Na yellow, K lilac, Rb reddish violet, Cs blue), because the flame's heat excites the valence electron to a higher level, and the energy released as it falls back is emitted as visible light of a characteristic wavelength.