The world around us is made up of atoms, and the properties of these atoms are largely determined by how their electrons are arranged. The Aufbau Principle describes the order in which electrons fill atomic orbitals known as electron configuration.
What is the Aufbau Principle?
“Aufbau” is a German word meaning “building up.”
In simple terms, it states:
Electrons first occupy the orbitals with the lowest energy levels before filling higher energy levels.
Orbital Structure and Energy Levels
Atoms have several energy levels where electrons can be found. These energy levels are called principal energy levels and are labeled as n=1, n=2, n=3, etc., where n represents the energy level or shell.
Within these energy levels, there are orbitals which are regions of space where electrons are likely to be found.
There are four types of orbitals: s, p, d, and f.
- s-orbitals: These are spherical and can hold up to 2 electrons.
- p-orbitals: These are dumbbell-shaped and can hold up to 6 electrons.
- d-orbitals: These are more complex in shape and can hold up to 10 electrons.
- f-orbitals: These have more complex shapes and can hold up to 14 electrons.
Each energy level can hold a certain number of electrons depending on the type of orbitals it contains.
The filling order follows this sequence:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d
Mnemonics to Remember the Aufbau Order
“School, School, Public School, Public School, Divisional Public School, Divisional Public School, Federal Divisional Public School, Federal Divisional Public School (ss,ps,ps, dps, dps, fdps, fdps).”
How to Use:
“School, School” refers to 1s and 2s.
“Public School, Public School” represents 2p, 3s and 3p, 4s.
“Divisional Public School, Divisional Public School” refers to 3d, 4p, 5s and 4d, 5p, 6s.
“Federal Divisional Public School, Federal Divisional Public School” represents 4f, 5d, 6p, 7s and 5f, 6d, 7p, 8s.
According to the Aufbau Principle:
- Electrons will fill the lowest-energy orbitals first.
- The 1s orbital is the lowest in energy, followed by 2s, 2p, and so on.
- As more electrons are added to an atom, they fill the available orbitals in the order of increasing energy.
Example of Electron Configuration using the Aufbau Principle:
Let’s take an example using the element Carbon (C):
- The atomic number of Carbon is 6.
- The electron configuration for Carbon is 1s² 2s² 2p².
- This means that Carbon has two electrons in the 1s orbital, two electrons in the 2s orbital, and two electrons in the 2p orbital.
Similarly:
Calcium (Ca):
1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
Iodine (I):
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵
Total electrons = 53
Key Principles Supporting the Aufbau Principle
Three important rules dictate how electrons fill orbitals:
Pauli Exclusion Principle: Each orbital can hold a maximum of 2 electrons, and they must have opposite spins.
Hund’s Rule: Within a sublevel, electrons occupy empty orbitals before pairing up. For example, in the 2p sublevel, electrons will spread out over the three p orbitals before pairing.
(n+l) Rule: This rule states that the lower the sum of the principal quantum number (n) and the azimuthal quantum number (l) (which corresponds to the sublevel: s=0, p=1, d=2, f=3), the lower the energy of the orbital. If two orbitals have the same (n+l) value, the orbital with the lower n value has lower energy.
Exceptions to the Aufbau Principle
While the Aufbau Principle works for most elements, there are exceptions, especially in transition metals, ions and heavier elements. These exceptions occur because half-filled and fully-filled orbitals provide extra stability.
For example:
- Chromium (Cr): Instead of 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁴, its configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵. The 4s electron moves to the 3d orbital to make it half-filled (more stable).
- Copper (Cu): Instead of 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁹, its configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰. A similar shift occurs to achieve a fully-filled d orbital.
- Ions: When forming ions, electrons are removed from the highest energy level first, not necessarily the last orbital filled according to the Aufbau principle. For example, when Iron (Fe) loses two electrons to form Fe²⁺, it loses the 4s electrons first, not the 3d electrons.
Fe (26): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶
Fe²⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶
Limitations of the Aufbau Principle
- The principle does not account for the repulsion between electrons in multi-electron atoms, which can slightly alter energy levels.
- In very heavy atoms, relativistic effects can change orbital energies, leading to deviations from the expected order.
- For some elements, the energy difference between orbitals is so small that external factors (like chemical bonding) can affect the electron configuration.
