Electron configuration and energy levels: write Regents electron configurations, distinguish ground state from excited state, and explain how electrons absorb and emit specific amounts of energy as photons.

What this topic is asking

The Core Curriculum asks you to describe how electrons occupy energy levels (shells) around the nucleus, to write an atom's electron configuration in the New York shell notation, to tell a ground state from an excited state, and to explain that electrons absorb and release specific amounts of energy when they move between levels. The Periodic Table in the Reference Tables prints the ground-state configuration under every element, so much of this is reading, but the ground-versus-excited distinction is tested directly.

Energy levels and the New York configuration

You do not need orbital notation (no $1s^2 2s^2 2p^6$, no Hund's rule) for the Regents; the shell notation is the expected form. The total of the numbers in a configuration equals the atomic number for a neutral atom, which is a quick way to check your answer.

Ground state versus excited state

This is the single most tested idea in this topic. To decide whether a configuration is ground or excited: first add the numbers to confirm the element (the total must match the atomic number); then check whether the lower levels are filled before the higher ones. Sodium's ground state is $2\text{-}8\text{-}1$; the configuration $2\text{-}7\text{-}2$ has the same $11$ electrons but the second level is not full while the third is occupied, so it is an excited state of sodium.

Absorbing and emitting energy

When an atom absorbs energy (from heat or electricity), an electron can jump from a lower level to a higher one, putting the atom in an excited state. The electron is unstable there and falls back, releasing the energy it absorbed as a photon of light. Because the energy levels are fixed, only specific energy differences are possible, so each element emits only certain colors of light.

Valence electrons

The valence electrons are the electrons in the outermost occupied energy level. They are the electrons involved in bonding, so this idea links straight into the bonding module. Sulfur ($2\text{-}8\text{-}6$) has $6$ valence electrons; sodium ($2\text{-}8\text{-}1$) has $1$. For the main-group elements, the number of valence electrons matches the group number pattern on the periodic table, which is why elements in the same group have similar chemistry.

Try this

Q1. Write the ground-state electron configuration of an aluminum atom (atomic number $13$). [1 point]

• Cue. $2\text{-}8\text{-}3$ (the numbers total $13$).

Q2. Explain how an excited electron produces a line in an element's emission spectrum. [1 point]

• Cue. It falls back to a lower level and releases the fixed energy difference as a photon of a specific color.