The electron configuration of oxygen is represented as $1s^2 \, 2s^2 \, 2p^4$.

Oxygen is the eighth element in the periodic table, indicating that it possesses a total of eight electrons. The electron configuration of an atom details how these electrons are distributed across various energy levels and sublevels. For oxygen, this configuration is $1s^2 \, 2s^2 \, 2p^4$.

In this notation, the first number in each group signifies the principal energy level, while the letter denotes the type of sublevel (s, p, d, or f). The superscript indicates the number of electrons present in that particular sublevel. Specifically, for oxygen, there are two electrons in the $1s$ sublevel, two electrons in the $2s$ sublevel, and four electrons in the $2p$ sublevel.

To elaborate, the $1s^2$ notation indicates that there are two electrons in the first energy level, which is closest to the nucleus, within the s sublevel. The $2s^2$ notation signifies that there are two electrons in the second energy level, also within the s sublevel. Meanwhile, the $2p^4$ notation indicates that there are four electrons in the second energy level, specifically in the p sublevel.

It is important to note that the s sublevel can accommodate a maximum of two electrons, while the p sublevel can hold up to six electrons. Consequently, the $2p$ sublevel of oxygen is not completely filled, which contributes to oxygen’s reactivity. This unfilled sublevel drives oxygen to readily form bonds with other atoms, allowing it to achieve a stable electron configuration by filling its outermost energy level. This characteristic is one reason why oxygen plays a crucial role in chemical reactions, including those that are essential for sustaining life on Earth.