Source: The Atom lecture, Purdue University General Chemistry
Difficulty: Intermediate | Prerequisites: Ions and Ion Formation notes, Light and Electromagnetic Radiation notes
Tags: Bohr model, electron orbits, energy levels, orbitals, electron configuration, aufbau principle, sublevels, s p d f, periodic table blocks, noble gas notation, abbreviated electron configuration, quantum, general chemistry
Once scientists understood that atoms emit and absorb specific wavelengths of light, the next question was: why? The Bohr model was the first attempt to explain this by placing electrons in fixed circular orbits around the nucleus. It works beautifully for hydrogen but fails for everything else. The modern orbital model replaced it, describing electrons as occupying specific three-dimensional regions (orbitals) with quantised energies. Electron configuration notation is the shorthand chemists use to describe which orbitals are occupied in a given atom. This is the foundation for understanding chemical bonding, periodic trends, and reactivity.
The Bohr model places electrons in circular orbits with fixed energies, but only works for hydrogen. The modern model uses orbitals (s, p, d, f) with quantised energies that fill from lowest to highest. Electron configurations describe which orbitals are occupied, and the periodic table is organised so you can read configurations straight off it.
Bohr model
A model of the atom in which electrons orbit the nucleus in specific circular paths at fixed distances, each path corresponding to a specific energy level. Think of it as a planetary model: electrons circle the nucleus the way planets circle the sun, but only in certain allowed orbits.
Energy level (principal energy level)
A shell around the nucleus where electrons can exist. Designated by the principal quantum number n (n = 1, 2, 3...). Higher n means greater distance from the nucleus and higher energy.
Orbital
A specific three-dimensional region around the nucleus where an electron is most likely to be found. Each orbital can hold a maximum of two electrons.
Sublevel (subshell)
A subdivision of an energy level, labelled s, p, d, or f. Each sublevel contains a specific number of orbitals: s has 1 orbital (2 electrons max), p has 3 orbitals (6 electrons max), d has 5 orbitals (10 electrons max), f has 7 orbitals (14 electrons max).
Quantised
Describes something that exists only in specific, discrete amounts rather than any arbitrary value. Orbital energies are quantised: electrons can only have certain energies, not values in between.
Electron configuration
A notation showing the distribution of electrons among the orbitals of an atom. Written as a series of sublevel labels with superscript electron counts, e.g. 1s^2 2s^2 2p^6.
Aufbau principle
The rule that electrons fill orbitals starting from the lowest energy level and working upward. "Aufbau" is German for "building up."
Abbreviated (noble gas) electron configuration
A shorthand that replaces the inner electron shells with the symbol of the nearest preceding noble gas in brackets. For example, sodium's full configuration 1s^2 2s^2 2p^6 3s^1 becomes [Ne] 3s^1.
Continuous spectrum
A spectrum containing all wavelengths of visible light with no gaps, such as a rainbow. Most radiation sources emit a continuous spectrum.
Emission spectrum (line spectrum)
A spectrum showing only specific wavelengths of light, appearing as distinct coloured lines. Each element has a unique emission spectrum, which acts like a fingerprint.
Electrons orbit the nucleus in specific circular paths.
Each orbit corresponds to a fixed energy level.
Electron energy increases as distance from the nucleus increases.
When an electron moves closer to the nucleus (drops to a lower orbit), it must release energy.
That released energy takes the form of a photon.
The energy of the emitted photon equals the difference between the two energy levels: ΔE = E(final) - E(initial).
The Bohr model successfully explains the hydrogen emission spectrum.
It fails for atoms with more than one electron, because electron-electron repulsions make the circular orbit picture too simple.
Scientists studied how light interacted with elements in the gas phase.
Three key findings emerged from experiments:
Electrons occupy specific volumes in the atom called orbitals, and these orbitals have specific shapes.
Orbitals have specific, discrete (quantised) energies.
Electrons fill orbitals in a specific order, from lowest energy to highest.
s sublevel: 1 orbital, holds a maximum of 2 electrons.
p sublevel: 3 orbitals, holds a maximum of 6 electrons.
d sublevel: 5 orbitals, holds a maximum of 10 electrons.
f sublevel: 7 orbitals, holds a maximum of 14 electrons.
Any single orbital, regardless of type, holds at most 2 electrons.
Fill orbitals from lowest energy upward (aufbau principle).
The filling order is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.
Note that 4s fills before 3d, and 5s fills before 4d. This is a common source of errors.
Examples from the lecture:
F (fluorine, 9 electrons): 1s^2 2s^2 2p^5
Al (aluminium, 13 electrons): 1s^2 2s^2 2p^6 3s^2 3p^1
K (potassium, 19 electrons): 1s^2 2s^2 2p^6 3s^2 3p^6 4s^1
Sc (scandium, 21 electrons): 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^1
Br (bromine, 35 electrons): 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^5
The periodic table is divided into blocks that correspond to the sublevel being filled:
s-block: groups 1 and 2 (plus helium)
p-block: groups 13 through 18
d-block: groups 3 through 12 (transition metals)
f-block: the two rows at the bottom (lanthanides and actinides)
The column number within each block tells you the number of electrons in that sublevel.
For s and p sublevels: the principal energy level number equals the period (row) number.
For the d sublevel: the principal energy level number is one less than the period number. So an element in period 4 filling a d sublevel is filling 3d, not 4d.
Identify the noble gas that comes just before your element on the periodic table.
Write that noble gas symbol in square brackets.
Then write out only the remaining electrons beyond that noble gas core.
Examples:
Li (lithium): full = 1s^2 2s^1, abbreviated = [He] 2s^1
Rb (rubidium): abbreviated = [Kr] 5s^1
Br (bromine): abbreviated = [Ar] 4s^2 3d^10 4p^5
Energy of emitted photon (Bohr model)
ΔE = E(final) - E(initial)
The energy change equals the difference between the final and initial energy levels. When the electron drops to a lower level, ΔE is negative (energy is released as a photon).
Emission spectra are the basis of spectroscopy, which is used to identify elements in distant stars, detect contaminants in water, and analyse the composition of unknown samples in forensic labs. Electron configuration determines how elements bond, which in turn determines the properties of every material you interact with.
Students often think the Bohr model works for all atoms. It does not. It only accurately predicts the spectrum of hydrogen. Multi-electron atoms require the orbital model.
A very common mistake is writing 3d before 4s in the filling order. Remember: 4s fills before 3d for elements in the fourth period.
Students sometimes confuse "orbital" with "orbit." An orbit (Bohr model) is a fixed circular path. An orbital (modern model) is a three-dimensional probability region with a specific shape.
When writing abbreviated configurations, students sometimes pick the wrong noble gas. Always use the noble gas at the end of the previous period, not the one closest in atomic number.
⚠️ You will be asked to write full and abbreviated electron configurations. Practise until these are automatic.
⚠️ The filling order (especially 4s before 3d) is tested frequently. Know the order cold and watch for the d-sublevel offset.
⚠️ Expect conceptual questions on why the Bohr model fails for multi-electron atoms.
⚠️ Questions about orbital capacity (s = 2, p = 6, d = 10, f = 14) appear in multiple-choice sections regularly.
⚠️ Using the periodic table to determine electron configuration is a skill that saves time on exams. Practise reading configurations directly off the table.
True or false: the Bohr model accurately predicts the emission spectra of all elements. (False, only hydrogen.)
Fill in the blank: a p sublevel contains ______ orbitals and can hold a maximum of ______ electrons. (3 orbitals, 6 electrons.)
True or false: 3d fills before 4s. (False, 4s fills before 3d.)
Fill in the blank: the abbreviated electron configuration of sodium (Na) is ______. ([Ne] 3s^1.)
True or false: the principal energy level for a d sublevel is the same as the period number. (False, it is one less than the period number.)
Q: Write the full electron configuration for chlorine (Cl, 17 electrons).
A: 1s^2 2s^2 2p^6 3s^2 3p^5.
Q: Write the abbreviated electron configuration for iron (Fe, 26 electrons).
A: [Ar] 4s^2 3d^6.
Q: Why does the Bohr model fail for helium and heavier atoms?
A: Because it does not account for electron-electron repulsions. With more than one electron, the simple circular orbit model cannot accurately predict energy levels or spectral lines.
Q: An electron in a hydrogen atom drops from the n = 3 level to the n = 1 level. What happens?
A: The electron loses energy equal to the difference between the two levels, and that energy is emitted as a photon. The wavelength of the photon corresponds to that specific energy difference.
Q: How many electrons can the 3d sublevel hold?
A: 10 electrons (5 orbitals, each holding a maximum of 2 electrons).
Q: Using the periodic table, how do you determine that bromine's outermost electrons are in the 4p sublevel?
A: Bromine is in period 4 and in the p-block. For p sublevels, the principal energy level equals the period number, so the outermost p electrons are in 4p.
Electron configuration directly determines an element's chemical properties and its position on the periodic table. This connects to periodic trends (ionisation energy, electronegativity, atomic radius), which you will study next. It also connects to chemical bonding: the number of valence electrons, read from the configuration, dictates how atoms form ionic and covalent bonds.
Bohr model, Bohr atom, energy levels, principal quantum number, orbitals, subshells, sublevels, s p d f, electron configuration, aufbau principle, filling order, noble gas notation, abbreviated electron configuration, core electrons, valence electrons, periodic table blocks, s-block, p-block, d-block, f-block, emission spectrum, line spectrum, continuous spectrum, quantised energy, hydrogen spectrum, general chemistry electron configuration, Purdue general chemistry