Explain why the electron affinity of nitrogen is approximately zero, while the elements on either side, carbon and oxygen, have substantial positive electron affinities.

Understanding the electron configuration of an atom is crucial in explaining its chemical behavior, including electron affinity. The electron configuration details how electrons are distributed in an atom's orbitals. Electrons fill orbitals in a way that minimizes the energy of the atom, starting with the lowest energy levels.

Let’s take nitrogen as an example. It has an electron configuration of 1s2 2s2 2p3, which fills the 2p orbitals exactly to half. This half-filled state is known for its stability because it offers a favorable balance of electron repulsion and orbital occupation. When an added electron would disturb this stability, the atom may exhibit less desire to capture the electron, explaining nitrogen’s low electron affinity.

How does this relate to chemical stability? Atoms adopt electron configurations that confer stability. A fully or half-filled set of orbitals represents a particularly stable configuration. In contrast, other elements like carbon (1s2 2s2 2p2) and oxygen (1s2 2s2 2p4) have orbitals that are not half or fully filled, making them more inclined to accept an additional electron to move towards a more stable configuration, resulting in a release of energy – evidence of higher electron affinity.

This intricacy in electron distribution is integral in predicting how an atom will interact in chemical processes.

Periodic trends play a fundamental role in understanding why different elements exhibit the electron affinities they do. As you move across a period from left to right on the periodic table, the effective nuclear charge increases as protons are added to the nucleus, leading to a tighter hold on the valence electrons. This generally causes an increase in electron affinity, with some notable exceptions.

Nitrogen is a quintessential exception. Despite nitrogen being situated between carbon and oxygen in the periodic table, it does not have a higher electron affinity. This is because of its electron configuration, where the added electron would break the half-filled stability.

The periodic table patterns, however, are not absolute and can be influenced by atomic structure nuances. For instance, the added stability of a half-filled p-orbital in nitrogen overrides the expected trend. While understanding periodic trends offers a great starting point, it's the underlying atomic structure that ultimately dictates an atom's electron affinity.

Chemical stability is an attribute that describes an atom's resistance to change in its electron configuration. Elements strive for stability, often achieved through a full or half-filled valence shell, which in turn influences their chemical properties and reactions.

Nitrogen’s half-filled 2p orbital configuration provides it with considerable stability. Adding another electron would mean introducing electron-electron repulsions in a previously stable p-orbital, thus reducing stability. As a result, nitrogen's electron affinity is close to zero because the system requires additional energy to overcome the electron-electron repulsions and force the extra electron into a less favorable position. In contrast, elements that are not in such stable configurations, like carbon and oxygen, gain additional stability upon the addition of an electron and thus release energy.

Why is stability favored? Chemically stable configurations are less reactive and more resistant to forming compounds, which is why noble gases with complete valence shells are largely unreactive. An element’s pursuit of stability is a driving force in chemical reactions and is the underlying reason why atoms have varying affinities for additional electrons.