Animated Solution for Chemistry - Periodicity in Properties: The process that is not endothermic in nature is

Imagine the atomic world as a bustling dance floor. Every atom is constantly trying to find its most comfortable, stable position. Sometimes, achieving this stability requires an atom to absorb energy from its surroundings—this is an endothermic process. Other times, an atom naturally falls into a more stable state, releasing a burst of energy in the process—this is an exothermic process.

In this problem, we are on a mission to find the process that is not endothermic. In simpler terms, we are hunting for the exothermic reaction among the given options. Let's analyze the energetic struggles of each atom.

Let's look at the first scenario: $O^{-}(g)+e^{-}\to O^{2-}(g)$.

At first glance, you might think, "Oxygen wants a full octet, so it should happily accept another electron!" But there is a massive catch here. The oxygen atom already has a negative charge ($O^{-}$). When you try to force another negatively charged electron into its electron cloud, you encounter fierce electrostatic repulsion.

It's like trying to push the negative poles of two strong magnets together. They fight back! To overcome this intense repulsion and force the electron into the atom, you must actively pump energy into the system. Therefore, the second electron gain enthalpy is always positive, making this process strictly endothermic.

Next, we examine the ionization of sodium: $Na(g)\to N{a}^{+}(g)+e^{-}$.

In this process, we are ripping an electron away from a neutral sodium atom. The positively charged nucleus is holding onto that electron with an attractive electrostatic force.

To break this bond and pull the electron away into the vacuum, you have to do work. You must supply energy to overcome the nuclear attraction. Because of this, ionization enthalpy is universally positive. You will never find an atom that spontaneously ejects an electron without an energy input. Thus, this process is also endothermic.

Now, let's consider the noble gas, Argon: $Ar(g)+e^{-}\to A{r}^{-}(g)$.

Argon is the zen master of the periodic table. It possesses a perfectly full valence shell with an $n{s}^{2}n{p}^6$ configuration. It has achieved the ultimate chemical nirvana—the stable octet.

If you try to hand Argon an extra electron, it simply has no room for it in its current energy level. The new electron must be forced into a completely new, higher-energy shell. This creates a highly unstable and energetically unfavorable state. Because the atom strongly resists this disruption to its stability, you must force the electron in by supplying energy. Consequently, the electron gain enthalpy of noble gases is always positive (endothermic).

Finally, we arrive at our exception: $H(g)+e^{-}\to H^{-}(g)$.

A neutral hydrogen atom has a single electron in its $1s$ orbital ($1s^1$). The $1s$ subshell can hold a maximum of two electrons. If hydrogen gains just one more electron, it achieves the $1s^2$ configuration.

This is a magical moment for hydrogen because $1s^2$ is the exact electron configuration of the highly stable noble gas, Helium. By accepting this electron, hydrogen drops into a much more stable, lower-energy state. Because it naturally "wants" to reach this stability, the system releases its excess energy into the surroundings.

This release of energy means the process has a negative $\Delta H$, making it exothermic.

By carefully analyzing the energetic motivations of each atom, we've uncovered the truth. While adding an electron to a negative ion, ionizing a neutral atom, and disrupting a noble gas all require an input of energy, completing the $1s$ shell of hydrogen releases energy.

Therefore, the formation of the hydride ion ($H^{-}$) is the only process that is not endothermic.