Electron configuration and energy levels: describe how electrons occupy energy levels, write electron configurations, identify valence electrons, and relate ground and excited states to spectra.

What this topic is asking

Within standard CH.2, Virginia expects you to describe how electrons are arranged in energy levels, to write an electron configuration, to identify the valence electrons that drive chemistry, and to relate the ground and excited states of an atom to the line spectra it produces. The arrangement of electrons explains an element's position in the periodic table and how it bonds.

Energy levels and sublevels

The first level has only an s sublevel ($1s$, up to 2 electrons). The second level has s and p ($2s$ and $2p$, up to 8 electrons). Filling from the bottom up gives the ground-state arrangement and explains the structure of the periodic table, where each period corresponds to filling a new principal energy level.

Writing an electron configuration

An electron configuration records the occupied sublevels and how many electrons each holds, in order of increasing energy. For sodium (11 electrons): $1s^2\,2s^2\,2p^6\,3s^1$. The superscripts add up to the total number of electrons, which for a neutral atom equals the atomic number. The Bohr-style "shell" notation (2, 8, 1 for sodium) carries the same information at the level the SOL emphasizes: how many electrons are in each principal energy level.

Valence electrons

For main-group elements, the number of valence electrons matches the group number pattern: Group 1 has one, Group 2 has two, and Groups 13 to 18 have three to eight. Atoms gain, lose or share valence electrons to reach a stable, full outer level (the octet of eight for most main-group elements), which is the basis of bonding.

Ground state, excited state and spectra

In the ground state every electron is in the lowest energy position available. If the atom absorbs energy (heat, electricity, light), an electron can jump to a higher level, an excited state. The excited state is unstable, so the electron falls back, releasing the energy difference as a photon of light.

Because the energy levels are fixed (quantized), only specific energy differences are possible, so the emitted light has only specific wavelengths and colors. This produces a bright-line emission spectrum unique to each element, which is both how flame tests identify metals and historic evidence that electrons occupy discrete energy levels rather than any arbitrary energy.

Try this

Q1. Write the electron configuration of fluorine (atomic number $9$). [1 point]

• Cue. $1s^2\,2s^2\,2p^5$, giving $7$ valence electrons.

Q2. State what happens to an electron when an atom moves from the ground state to an excited state. [1 point]

• Cue. It absorbs energy and jumps from a lower energy level to a higher one.