Source: Chapter 7, Sections 7.1, 7.8–7.9 | General Chemistry, Purdue University
Tags: Bohr model, electron configuration, orbitals, energy levels, aufbau principle, periodic table blocks, s block, p block, d block, quantum, sublevel
Difficulty: Introductory to Intermediate
Prerequisites: understanding of electromagnetic radiation and photon energy (see Part 1 study notes), familiarity with the periodic table layout.
Now that you know light carries quantised energy (from Part 1), this set of notes covers what that means for the structure of the atom. Scientists used emission spectra to work out that electrons occupy specific energy levels, and the Bohr model was the first attempt to describe that structure for hydrogen.
The Bohr model has limits (it only works for hydrogen), but it introduces the key idea: electrons live at fixed energy levels, and moving between levels means absorbing or releasing a photon. From there, the course moves to orbitals and electron configurations, which describe where electrons actually sit in multi-electron atoms and how the periodic table is organised around that.
The Bohr model places electrons in numbered circular orbits around the nucleus, with energy increasing as distance from the nucleus increases. When an electron drops to a lower orbit, it emits a photon. Electron configurations describe the order in which orbitals fill (lowest energy first), and the periodic table’s block structure (s, p, d, f) maps directly onto those sublevels.
Bohr model
A model of the hydrogen atom in which electrons orbit the nucleus in specific circular paths at fixed energy levels. Proposed by Niels Bohr to explain hydrogen’s line spectrum.
In simple terms: electrons travel in set circular lanes around the nucleus, like cars on a roundabout with fixed lanes.
Orbit (Bohr orbit) / energy level (n)
A fixed circular pathway around the nucleus. Each orbit is labelled with a number (n = 1, 2, 3…). Orbit 1 is closest to the nucleus and has the lowest energy.
Think of it as: the higher the orbit number, the further from the nucleus and the more energy the electron has.
ΔE (change in energy)
The difference in energy between two orbits. Calculated as ΔE = E(final) – E(initial). When an electron drops to a lower level, ΔE is negative and a photon is emitted.
Orbital
A region of space around the nucleus where there is a high probability of finding an electron. Unlike Bohr’s circular orbits, orbitals have specific three-dimensional shapes (s, p, d, f).
In simple terms: orbitals are not paths. They are probability clouds showing where an electron is most likely to be found.
Quantised
Restricted to specific, discrete values rather than a continuous range. Electron energies in atoms are quantised, meaning electrons can only have certain energies, not any arbitrary value.
Electron configuration
A notation that describes which orbitals the electrons of an atom occupy and how many electrons are in each. For example, the electron configuration of lithium is 1s²2s¹.
Sublevel (subshell)
A subdivision of an energy level, designated by the letters s, p, d, or f. Each sublevel contains a specific number of orbitals: s = 1, p = 3, d = 5, f = 7.
Principal energy level number (n)
The number that appears before the sublevel letter in an electron configuration (e.g. the "3" in 3p). For s and p sublevels, n equals the period number on the periodic table. For d sublevels, n is one less than the period number.
Aufbau principle
Electrons fill orbitals starting from the lowest energy and working upward. "Aufbau" is German for "building up."
In simple terms: fill from the bottom of the energy ladder first.
What Bohr proposed
Niels Bohr studied the line spectrum of hydrogen to understand how electrons behave.
He proposed that electrons orbit the nucleus in specific circular pathways, each at a fixed distance from the nucleus.
Each orbit has a specific energy. The further the orbit from the nucleus, the higher the energy.
Orbit numbering
Orbits are labelled with integers: n = 1, 2, 3, 4…
n = 1 is the closest orbit to the nucleus (lowest energy).
As n increases, the orbit is further away and the electron has more energy.
Energy transitions and photon emission
When an electron moves from a higher-energy orbit to a lower-energy orbit, it releases energy.
That energy is emitted as a photon of light.
The energy of the emitted photon equals the difference between the two energy levels: ΔE = E(final) – E(initial).
This is why hydrogen’s emission spectrum has discrete lines: each line corresponds to a specific electron transition between specific orbits.
Limitations of the Bohr model
The Bohr model works only for hydrogen (one-electron atoms).
It cannot accurately predict the spectra or behaviour of multi-electron atoms.
The modern quantum mechanical model replaced it, but the core idea of quantised energy levels carries forward.
From orbits to orbitals
Experiments showed that electrons occupy specific three-dimensional regions called orbitals (not the circular orbits of Bohr). Each orbital has a distinct energy and shape.
Orbitals fill in order from lowest energy to highest. This is the aufbau (building-up) principle.
Orbital types and electron capacity
Each orbital holds a maximum of 2 electrons.
s sublevel: 1 orbital, holds 2 electrons total.
p sublevel: 3 orbitals, holds 6 electrons total.
d sublevel: 5 orbitals, holds 10 electrons total.
f sublevel: 7 orbitals, holds 14 electrons total.
Filling order
The standard 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. The diagonal rule (or the periodic table itself) is the easiest way to remember this.
Writing electron configurations
Write the sublevels in filling order, with the number of electrons as a superscript. Examples:
Li (3 electrons): 1s²2s¹
Na (11 electrons): 1s²2s²2p⁶3s¹
Cl (17 electrons): 1s²2s²2p⁶3s²3p⁵
Ca (20 electrons): 1s²2s²2p⁶3s²3p⁶4s²
Br (35 electrons): 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁵
The periodic table as a map of electron configurations
The periodic table is divided into blocks (s, p, d, f) that correspond to the sublevel being filled by the last electron added.
s block: Groups 1 and 2 (plus helium). The outermost electrons are in an s sublevel.
p block: Groups 13–18. The outermost electrons are in a p sublevel.
d block: Groups 3–12 (transition metals). The outermost electrons being added are in a d sublevel.
f block: Lanthanides and actinides (the two rows pulled out at the bottom). Electrons fill f sublevels.
Columns and electron counts
The number of columns in each block matches the maximum number of electrons that sublevel can hold: s block has 2 columns (2 electrons), p block has 6 columns (6 electrons), d block has 10 columns (10 electrons), f block has 14 columns (14 electrons).
Principal energy level number (n) and the period
For s and p sublevels: n equals the period (row) number. Sodium is in Period 3, and its outermost electron is in 3s.
For d sublevels: n is one less than the period number. The first row of transition metals is in Period 4, but their d electrons are in 3d.
Patterns by group
Alkali metals (Group 1) all end in ns¹ (one electron in the outermost s sublevel).
Alkaline earth metals (Group 2) all end in ns².
Halogens (Group 17) all end in np⁵.
Noble gases (Group 18) all end in np⁶ (full outer p sublevel).
This is why elements in the same group share similar chemical properties: they have the same number of valence electrons in the same type of sublevel.
Energy change for electron transitions (Bohr model)
ΔE = E(final) – E(initial)
When an electron drops to a lower level, ΔE is negative (energy is released as a photon). The magnitude of ΔE equals the energy of the emitted photon.
Maximum electrons per sublevel
Sublevel | Number of Orbitals | Max Electrons |
|---|---|---|
s | 1 | 2 |
p | 3 | 6 |
d | 5 | 10 |
f | 7 | 14 |
Example electron configurations by group
Group | Element | Configuration |
|---|---|---|
Alkali metals | Li | 1s²2s¹ |
Alkali metals | Na | 1s²2s²2p⁶3s¹ |
Alkali metals | K | 1s²2s²2p⁶3s²3p⁶4s¹ |
Halogens | F | 1s²2s²2p⁵ |
Halogens | Cl | 1s²2s²2p⁶3s²3p⁵ |
Alkaline earth | Be | 1s²2s² |
Alkaline earth | Ca | 1s²2s²2p⁶3s²3p⁶4s² |
Noble gases | Ne | 1s²2s²2p⁶ |
Noble gases | Ar | 1s²2s²2p⁶3s²3p⁶ |
Neon signs work because of electron transitions. Electricity excites electrons in neon gas to higher energy levels. When those electrons fall back down, they emit photons at specific wavelengths, producing the characteristic orange-red glow. Different gases produce different colours for the same reason.
Electron configurations explain why elements in the same group behave similarly in chemical reactions. Sodium and potassium are both reactive metals that explode in water because they both have a single electron in their outermost s sublevel, which they readily lose.
Students often think the Bohr model applies to all atoms. It does not. It works only for hydrogen (and other single-electron species like He⁺). For anything with more than one electron, you need the quantum mechanical model.
Students frequently confuse orbits and orbitals. An orbit (Bohr) is a fixed circular path. An orbital (modern model) is a three-dimensional probability cloud. They are different concepts.
Students sometimes assume 3d fills before 4s because 3 is less than 4. In reality, 4s is lower in energy than 3d and fills first. The periodic table itself is the best guide: the 4s block comes before the 3d block.
Students often forget that the principal energy level for d sublevels is one less than the period number. Elements in Period 4 fill 3d, not 4d.
⚠️ Writing full electron configurations from scratch is a near-certainty on the exam. Practise until you can do it without the diagonal chart.
⚠️ Know the limitations of the Bohr model. A question asking "which model correctly predicts the spectrum of helium?" is testing whether you know Bohr only works for hydrogen.
⚠️ Be prepared to identify an element from its electron configuration, or write the configuration given the element.
⚠️ Understand why elements in the same group have similar properties (same number and type of valence electrons). This is a conceptual question that appears frequently.
⚠️ The d-block principal energy level being one less than the period number is a specific detail examiners like to test. Period 4 transition metals have 3d electrons, not 4d.
True or false: The Bohr model accurately predicts the emission spectrum of helium. (False. It only works for hydrogen.)
Fill in the blank: In the Bohr model, when an electron drops from a higher orbit to a lower orbit, it ______ a photon. (emits / releases)
True or false: A d sublevel can hold a maximum of 10 electrons. (True. 5 orbitals × 2 electrons each.)
Fill in the blank: For d sublevels, the principal energy level number is ______ than the period number. (one less)
True or false: Elements in the same group of the periodic table have different numbers of valence electrons. (False. Same group = same number of valence electrons.)
Q: What happens when an electron in a hydrogen atom moves from n = 4 to n = 2?
A: The electron loses energy. That energy is emitted as a photon whose energy equals the difference between the two levels: ΔE = E(2) – E(4). The photon’s wavelength corresponds to a specific line in hydrogen’s emission spectrum.
Q: Write the full electron configuration of potassium (K, atomic number 19).
A: 1s²2s²2p⁶3s²3p⁶4s¹. Note that 4s fills before 3d.
Q: Why does the Bohr model fail for helium?
A: Helium has two electrons, and their mutual repulsion introduces interactions the Bohr model cannot account for. The model only handles one-electron systems.
Q: An element has the electron configuration 1s²2s²2p⁶3s²3p⁵. Identify the element and its group.
A: Chlorine (Cl). It ends in 3p⁵, placing it in the p block, Group 17 (halogens).
Q: Explain why calcium and strontium have similar chemical properties.
A: Both are alkaline earth metals (Group 2). Both have electron configurations ending in ns², meaning they each have two valence electrons in an s sublevel. The same valence electron arrangement leads to similar reactivity.
This connects back to electromagnetic radiation and photon energy (Part 1 notes). The Bohr model’s explanation of emission lines depends entirely on the relationship E = hc/λ.
Electron configurations feed directly into the next major topics in the course: periodic trends (atomic radius, ionisation energy, electronegativity) and chemical bonding. The number and arrangement of valence electrons determine how atoms bond, so this material is foundational for everything that follows.
The idea of quantised energy levels also connects forward to molecular orbital theory and spectroscopy in later chemistry courses.
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