Justify the following: (a) Lithium forms the normal oxide whereas potassium forms the super oxide an burning in air. (b) The melting point gap between lithium and sodium is maximum.

Lithium, when exposed to air, undergoes a reaction with oxygen that leads to the creation of lithium oxide. This compound, represented by the chemical formula \(Li_2O\), is the result of lithium's moderate reactivity, which is linked to its position at the top of the alkali metals group in the periodic table. As an element with a relatively strong attraction between its nucleus and electron, lithium tends to lose one electron, producing a stable oxide product. In contrast to heavier alkali metals, which may form peroxides or superoxides, lithium's low reactivity only allows the formation of the typical oxide, signifying a simple transfer of its outer electron to oxygen.

The formation of lithium oxide can be equationed as:\[4Li (s) + O_2 (g) \rightarrow 2Li_2O (s)\]
During this process, the oxygen molecule \(O_2\) is reduced to oxide ions \(O^{2-}\), while lithium is oxidized to lithium ions \(Li^+\). This exchange underscores the straightforward nature of chemical interactions involving lithium, where more complex oxide types, such as peroxides or superoxides, are not typically formed.

Potassium diverges from lithium's behavior by reacting with oxygen to form a much more reactive oxide known as potassium superoxide, denoted by the formula \(KO_2\). Superoxides are special in that they contain the superoxide ion \(O_2^−\), which consists of a more complex electronic structure with extra oxygen atoms.

Potassium's position lower in the alkali metals group confers it a higher reactivity due to its larger atomic radius and lower ionization energy. Consequently, when potassium burns in air, it can stabilize the extra negative charge on the oxygen molecule, thanks to the larger distance between the outermost electron and its nucleus. This unique capability stems from the metal's ability to more readily lose an electron compared to lithium, thus facilitating the formation of the superoxide.
The interaction can be summarized by the equation:\[4K (s) + O_2 (g) \rightarrow 2KO_2 (s)\]
This superoxide formation is crucial not just scientifically but also practically, as potassium superoxide is used in life support systems to both release oxygen and absorb carbon dioxide.

The melting points of alkali metals exhibit a notable trend within the group; they decrease as one moves from lithium down to cesium. This is generally implicated by the atomic structure of these metals, where a larger atomic radius results in weaker metallic bonding. Consequently, with less energy needed to disrupt the metal lattice, the metal melts at a lower temperature.

Lithium possesses the highest melting point among the alkali metals due to its small atomic radius, which strengthens the metallic bonds. This stronger bonding network demands more energy to break apart the solid structure, leading to a higher melting point. As we descend the group, each subsequent element has a larger atomic radius and a corresponding decrease in melting point. For example, the gap in melting points between lithium and sodium is particularly wide because sodium atoms are significantly larger, thus their metallic bonds are weaker. This concept is key to understanding the physical properties and practical uses of alkali metals in various applications where temperature stability is a factor.

Ionization energy refers to the energy needed to remove the most loosely bound electron from an atom in its gaseous state. In alkali metals, this ionization energy decreases down the group from lithium to francium. The reason for this trend lies in the increasing atomic radius and shielding effect, which collectively reduce the electrostatic pull of the nucleus on the outer electrons.

Lithium, being at the top of the alkali group, has the highest ionization energy because its electrons are closer to the nucleus, and there is minimal shielding from inner electrons. Therefore, it requires a greater amount of energy to remove its outermost electron compared to other alkali metals. The substantial ionization energies of alkali metals like lithium are responsible for their chemical reactions, as it influences the ease with which these metals can lose an electron and thereby engage in compounds formation, such as forming lithium oxide rather than a superoxide or peroxide. Understanding ionization energy is essential for predicting reactivity patterns and bonding preferences in alkali metals.