The electron configurations described in this chapter all refer to gaseous atoms in their ground states. An atom may absorb a quantum of energy and promote one of its electrons to a higher-energy orbital. When this happens, we say that the atom is in an excited state. The electron configurations of some excited atoms are given. Identify these atoms and write their ground-state configurations: (a) \(1 s^{1} 2 s^{1}\) (b) \(1 s^{2} 2 s^{2} 2 p^{2} 3 d^{1}\) (c) \(1 s^{2} 2 s^{2} 2 p^{6} 4 s^{1}\) (d) \([\mathrm{Ar}] 4 s^{1} 3 d^{10} 4 p^{4}\) (e) \([\mathrm{Ne}] 3 s^{2} 3 p^{4} 3 d^{1}\)

The ground state of an atom is the lowest energy arrangement of its electrons. In the ground state, electrons occupy the available orbitals starting from the lowest energy levels, up to the highest, following the Aufbau principle, Pauli exclusion principle, and Hund's rule. For example, the ground state electron configuration of a hydrogen atom would be written as \(1s^1\).

It's important to see the ground state as the 'default' position for electrons within an atom. From there, any additional energy absorbed by an atom can shift the electrons to higher energy levels, causing the atom to enter what is known as an excited state. The ground state configurations form the basis for understanding chemical properties and reactivity of elements.

In contrast to the ground state, an excited state occurs when an atom absorbs energy and one or more electrons are promoted to higher energy orbitals. This results in an electron configuration that is higher in energy and not as stable as the ground state configuration. These states are temporary and the atom will eventually release energy to return to its ground state.

Taking element (a) from our exercise as an example, \(1s^1 2s^1\) represents hydrogen in an excited state, since one of its electrons has absorbed energy to move to the 2s orbital, deviating from its ground state configuration of \(1s^1\). Understanding the differences between ground and excited states allows us to explain phenomena such as spectral lines and gives insight into an atom's electron distribution under different conditions.

Quantum energy levels are central to the concept of electron configurations. These levels are regions around the nucleus where electrons are likely to be found; they are defined by principal quantum numbers (n). Electrons in lower quantum numbers (closer to the nucleus) are in lower energy states, and as n increases, so does the energy of the orbital and the electrons within it.

For example, the electron configuration \(1s^2 2s^2 2p^2 3d^1\) describes the distribution of electrons across different energy levels. In this configuration, \(1s^2\) represents the two electrons in the first energy level (closest to the nucleus and lowest in energy), proceeding to higher levels with the \(2s^2\), \(2p^2\), and \(3d^1\). The presence of an electron in the \(3d\) orbital indicates a higher quantum energy level than the ground state and typically represents an excited state.

Noble gas notation is a useful shortcut when writing electron configurations for elements with a large number of electrons. It builds upon the electron configuration of the nearest noble gas that precedes the element in the periodic table, as noble gases have fully filled electron shells, making them stable and chemically inert.

For instance, the noble gas notation for a selenium atom not in its ground state is given as \([Ar] 4s^1 3d^{10} 4p^4\). Here, \([Ar]\) represents the electron configuration up to argon, the previous noble gas. From this point, the notation describes the additional electron configuration beyond the noble gas. Noble gas notation simplifies the process of writing long electron configurations and emphasizes the relationship of elements with their closest noble gases.