Electron Configuration and Spectra

Energy Levels and Shell Notation

Electrons are not scattered randomly — they occupy energy levels (also called shells or principal energy levels), numbered n = 1, 2, 3, … starting closest to the nucleus. Lower-numbered levels are nearer the nucleus and lower in energy.

New York uses a compact electron-configuration notation that lists how many electrons are in each shell, separated by dashes. For example:

• Sodium (11 electrons): 2-8-1
• Oxygen (8 electrons): 2-6
• Chlorine (17 electrons): 2-8-7

The digits always add up to the total number of electrons — and for a neutral atom, to the atomic number.

Reading it off the Periodic Table

You do not have to memorize these. The Periodic Table on the 2025 Reference Tables prints the ground-state electron configuration directly under each element's symbol. On exam day, find the element and copy the shell numbers. This is one of the fastest free points on the test.

Maximum Electrons per Shell

Each shell can hold a limited number of electrons. The Regents-level rule of thumb uses the formula 2n², where n is the shell number.

In practice for light elements, electrons fill 2, then 8, then 8 again before the third shell takes more — which is why the Periodic Table's printed configurations are the safe reference rather than blindly filling to 18.

Ground State vs. Excited State

This distinction is tested almost every administration.

• Ground state: the lowest-energy arrangement, with electrons in the lowest available levels. It exactly matches the configuration printed on the Periodic Table. Sodium's ground state is 2-8-1.
• Excited state: an electron has absorbed energy and jumped to a higher level, leaving a gap below. The total still equals the atomic number, but the arrangement differs from the table. A sodium excited state could be 2-7-2 (an electron jumped from the second to the third shell).

How to tell them apart

A quick procedure for any configuration:

1. Add the digits — they must equal the element's number of electrons.
2. Compare to the Periodic Table's ground-state configuration.
3. If it matches exactly → ground state. If the total is right but a lower shell is not full while a higher shell has electrons → excited state.
4. If the total is wrong → it is not a valid configuration for that atom at all.

Example: For sodium (11 electrons), 2-8-1 is ground state; 2-7-2 is excited (total 11, but the second shell is not full); 2-8-2 is impossible (total 12, that is magnesium).

Valence Electrons

Valence electrons are the electrons in the outermost occupied shell. They are the electrons involved in bonding, so they control an element's chemistry.

• Chlorine (2-8-7) has 7 valence electrons.
• Sodium (2-8-1) has 1 valence electron.
• For representative (main-group) elements, the number of valence electrons matches the group number's ones digit (Group 1 → 1, Group 17 → 7).

Atoms tend to gain, lose, or share electrons to reach a stable octet (8 valence electrons, like the noble gases in Group 18). This is the basis for ion formation covered earlier.

Spectra: Light from Electron Transitions

This is where atomic structure connects to the Waves and Electromagnetic Radiation blueprint band.

• When an atom absorbs energy, an electron jumps from a lower level to a higher one (ground → excited).
• The excited electron is unstable and falls back down, releasing the energy as a photon of light.
• Because energy levels are fixed, only specific energy gaps are possible, so only specific colors (wavelengths) are emitted.

The result is a bright-line spectrum (emission spectrum): a set of discrete colored lines on a black background. Each element produces a unique pattern — a true fingerprint. Spectroscopists identify unknown elements, and even distant stars, by matching bright-line patterns.

Worked exam scenario

A cluster shows the bright-line spectra of elements A, B, and C plus a mixture. If the mixture's lines include every line of A and every line of C but none unique to B, the mixture contains A and C only. The presence of an element requires all of its characteristic lines to appear.

Linking spectra to the wave band

The Waves and Electromagnetic Radiation strand (5–7% of the exam) asks you to evaluate claims about radiation and matter. Two facts close the loop:

• Higher-energy transitions (a larger drop between levels) emit higher-frequency, shorter-wavelength light — toward the blue/violet end.
• Lower-energy transitions emit lower-frequency, longer-wavelength light — toward the red end.

So the color of an emitted line is direct evidence of the size of the energy gap an electron fell across. This is why flame tests work: sodium glows yellow, copper glows blue-green, each a fingerprint of its electron transitions.

Common Traps

• Calling an excited state invalid. As long as the digits sum to the right total, an excited state is a real configuration — just higher in energy than the ground state.
• Confusing absorption with emission. Energy is absorbed going up (ground → excited) and released as light coming down (excited → ground).
• Miscounting valence electrons by including inner shells — only the outermost occupied shell counts.
• Forgetting the table prints the answer. The ground-state configuration is supplied; reach for it before calculating.