Why are Group 1B elements more stable than Group 1A elements even though they seem to have the same outer electron configuration, ns1, where n is the principal quantum number of the outermost shell?

Explanation:

The elements in Group 1A elements are Lithium, Sodium, Potassium, Rubidium, Caesium, and Francium(radioactive). The valence shells of the members have the electronic configuration of ns¹ where n =2,3,4,5,6,7. They are known as alkali earth metals. Their chemistry is dominated by the only electron in the outer shell of their atoms. Thus, they are powerful reducing agents because of their ability to lose this outer shell electron readily.

Thus ns' electrons in Group 1A require lesser energy for ionization in the case of elements of Group 1A; this explains why Group 1A elements are more reactive.

On the other hand, the Group 1B elements outer ns' electrons are less effectively shielded by the inner d electrons; therefore, the outer electrons in case of Group 1B feel more nuclear pull (due to positively charged protons in the nucleus) as compared to the nuclear pull in Group 1A elements. Thus electrons require more energy for ionization in the case of elements of Group 1B. This explains why Group 1B elements are less reactive than elements of Group 1A.

alkaline metals are lithium, sodium, rubidium, potassium, cesium and francium.

alkali metals have electronegativity from 0.7 to 1, lowest in periodic table of elements, which means that alkali metals (i group in periodic table) has positive oxidation number in compounds.

electronegativity is a chemical property that describes the tendency of an atom to attract a shared pair of electrons towards itself.

alkaline metals (far left in main group) have lowest ionizations energy and easy remove valence electrons,

for example, potassium is the alkali metal and has a single valence electron in the outer electron shell.

electron configuration of potassium is: ₁₉k 1s²2s²2p⁶3s²3p⁶4s¹.

they can not exist alone in nature, because they are very reactive.