Neon (atomic number 10), located in group 18 as a noble gas, is characterized by its complete valence shell, having eight electrons in its outermost shell (1s² 2s² 2p⁶). This stable electron configuration makes neon largely unreactive, as it does not tend to gain, lose, or share electrons with other elements. In contrast, sodium (atomic number 11), positioned in group 1 as an alkali metal, has a single electron in its outermost shell (1s² 2s² 2p⁶ 3s¹). This makes sodium highly reactive since it readily loses that one valence electron to achieve a stable octet in the next lower energy level. The drastic difference in reactivity between these two elements can be attributed to their electron configurations and their positions in the periodic table; neon's filled valence shell imparts stability, while sodium's lone valence electron drives it toward chemical reactions to attain a more stable state.
How does the reactivity of neon (atomic number 10, period 2, group 18) differ from that of sodium (atomic number 11, period 3, group 1)? Explain your reasoning based on the properties of neon and sodium that you can infer based on their atomic numbers, periods, and groups.
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