1. Electron Configuration:
* Alkali metals (Group 1): Have one valence electron (electron in the outermost shell).
* Alkaline earth metals (Group 2): Have two valence electrons.
2. Noble Gas Configuration:
* Atoms are most stable when their outermost shell is full.
* Noble gases (Group 18) have a full outer shell, making them very unreactive.
3. Formation of Cations:
* Alkali metals easily lose their single valence electron to form a +1 cation, achieving the stable electron configuration of the preceding noble gas. For example, sodium (Na) loses one electron to become Na⁺, which has the same electron configuration as neon (Ne).
* Alkaline earth metals lose their two valence electrons to form a +2 cation, also achieving the stable electron configuration of the preceding noble gas. For example, magnesium (Mg) loses two electrons to become Mg²⁺, which has the same electron configuration as neon (Ne).
In summary:
* Alkali metals and alkaline earth metals have a strong tendency to lose electrons to achieve a stable electron configuration.
* This loss of electrons results in the formation of positively charged ions (cations).
Example:
* Sodium (Na): [Ne]3s¹ → Na⁺ + e⁻ (loses one electron, becomes like Neon)
* Magnesium (Mg): [Ne]3s² → Mg²⁺ + 2e⁻ (loses two electrons, becomes like Neon)
The ease with which alkali metals and alkaline earth metals form cations is a key factor in their chemical reactivity. They readily participate in ionic bonding, forming salts with nonmetals.
