Ground State vs Excited State Electron Configurations

When we write an electron configuration such as 1s² 2s² 2p⁶ 3s¹ for sodium, we’re describing the atom’s ground state: its lowest-energy arrangement of electrons. But atoms don’t always stay there. Heat, light or electricity can push electrons into higher orbitals, creating an excited state. Excited states are short-lived, but they’re responsible for the colours of flames, neon signs, fireworks and stars. This article compares the two and shows how to recognise each.

Quick comparison

Feature	Ground state	Excited state
Energy	lowest possible for the atom	higher than the ground state
Electron arrangement	follows the Aufbau principle, Hund’s rule and Pauli principle	one or more electrons in a higher orbital than normal
Stability	stable; the atom stays here indefinitely	unstable; usually lasts nanoseconds
How it arises	the normal state of an isolated atom	absorption of energy (heat, light, electrical discharge, collisions)
What happens next	nothing, unless energy is added	electron falls back, emitting a photon (or losing energy as heat)
Number of possible arrangements	one (for a given atom, ignoring spin details)	many

The ground state

The ground state is the arrangement predicted by:

• the Aufbau principle: fill orbitals from lowest energy upward
• the Pauli exclusion principle: at most two electrons per orbital, with opposite spins
• Hund’s rule: fill equal-energy orbitals singly before pairing

See the Aufbau principle, the Pauli exclusion principle and Hund’s rule.

Electron configurations in data tables, on the periodic table and on this site’s element pages are ground states unless stated otherwise.

Excited states

An excited state is any arrangement with more energy than the ground state. Usually, one electron has been promoted from its normal orbital to a higher one.

The number of electrons stays the same: the atom is still neutral and still the same element. Only the arrangement has changed.

How atoms get excited

• Heat: in a flame, collisions between fast-moving particles transfer energy to electrons. See flame test colours.
• Electricity: in a discharge tube or neon sign, electrons accelerated by a voltage collide with atoms.
• Light: an atom absorbs a photon whose energy exactly matches the gap between two levels. See photons and energy levels.
• Chemical reactions: some reactions leave products in excited states, which then emit light (chemiluminescence, as in glow sticks). See luminol and blood detection.

What happens next

Excited states are unstable. Within about 10⁻⁸ seconds (typical for an allowed transition), the electron drops back to a lower level, releasing the energy difference, usually as a photon of a specific wavelength:

ΔE = E(excited) − E(lower) = hf = hc/λ

The electron may fall straight to the ground state or cascade down in several steps, emitting several photons.

Worked examples: ground or excited?

Example 1: sodium (11 electrons)

• Ground state: 1s² 2s² 2p⁶ 3s¹
• Excited state: 1s² 2s² 2p⁶ 3p¹

In the excited state, the outer electron has jumped from 3s to 3p. When it falls back, it emits the famous yellow light at 589 nm, the colour of sodium street lamps and sodium flame tests.

Example 2: carbon (6 electrons)

• Ground state: 1s² 2s² 2p²
• Excited state: 1s² 2s¹ 2p³

One 2s electron has been promoted to 2p. This particular arrangement is also used (as a conceptual step) to explain why carbon forms four bonds, before hybridisation is introduced.

Example 3: identify the state

Which of these are excited states?

(a) 1s² 2s² 2p⁶ 3s² 3p⁵ (b) 1s² 2s² 2p⁵ 3s¹ (c) 1s² 2s¹ (d) 1s¹ 2s¹ (e) 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹

Answers:

• (a) Ground state of chlorine (17 electrons).
• (b) Excited state of neon (10 electrons): a 2p electron has moved to 3s. The ground state is 1s² 2s² 2p⁶.
• (c) Ground state of lithium.
• (d) Excited state of helium: one electron has left 1s even though 1s has room.
• (e) Excited state of potassium: the ground state is …3p⁶ 4s¹, because 4s is lower in energy than 3d in potassium.

How to spot an excited state

1. Count the electrons and write the ground-state configuration.
2. Compare. If any lower orbital has a vacancy while a higher orbital is occupied, it’s an excited state.
3. Watch for tricky cases like (e), where the order of 4s and 3d matters.

Also watch out for configurations that are impossible, not excited, such as 1s³ (breaks the Pauli principle) or 2d¹ (no d subshell exists for n = 2). See quantum numbers explained.

Why excited states matter

Line spectra

Every emission line in an element’s spectrum corresponds to an electron falling from one particular excited state to a lower one. Because the energy levels are fixed, the lines are at fixed wavelengths. See the hydrogen emission spectrum and emission vs absorption spectra.

Colour and light technology

• Flame tests and fireworks: heat creates excited states that emit characteristic colours. See fireworks and electron transitions.
• Neon and sodium lamps: electrical discharge creates excited states.
• Fluorescence: a molecule absorbs UV light, loses a little energy as heat, then emits visible light from a lower excited state.
• Lasers: many atoms are kept in a particular excited state (a population inversion), and stimulated emission produces a powerful beam of identical photons.

Chemical reactions

Excited molecules can be much more reactive than ground-state ones. Photochemistry, including photosynthesis and the formation of ozone and smog, starts when light creates excited states. See reaction rates and catalysts.

Energy level diagrams

Energy level diagrams make the difference easy to see. Each horizontal line represents an orbital’s energy, with the lowest at the bottom. In the ground state, electrons occupy the lowest available lines. In an excited state, an arrow shows an electron lifted to a higher line, leaving a gap below. A downward arrow drawn from the excited level to a lower one represents emission, and its length corresponds to the photon’s energy: longer arrows mean higher-energy, shorter-wavelength light.

Drawing the diagram before answering a question about spectra, emission or absorption is often the quickest way to avoid confusion between the two states.

Common mistakes

• Thinking an excited atom has gained or lost electrons. It has the same number; they’re just rearranged. (Losing an electron completely is ionisation, not excitation.)
• Assuming excited states follow Hund’s rule. They may not; that’s partly what makes them higher in energy.
• Confusing exceptions with excited states. Chromium’s [Ar] 3d⁵ 4s¹ is its ground state, not an excited state. See electron configuration exceptions.
• Forgetting 4s fills before 3d. 3d¹ in potassium indicates an excited state.

Key takeaways

• The ground state is the lowest-energy arrangement, following Aufbau, Pauli and Hund.
• An excited state has an electron in a higher orbital than normal; it’s unstable and short-lived.
• Atoms are excited by heat, electricity, light or chemical reactions, and emit photons when they relax.
• To spot an excited state, look for a vacancy in a lower orbital while a higher one is occupied.
• Excited states explain line spectra, flame colours, fluorescence, lasers and photochemistry.