The statement "3d comes before 4s" refers to a nuanced understanding of electron orbital energy levels and filling order, particularly for transition metals. While introductory chemistry often teaches that the 4s orbital fills before the 3d orbital, the reality for transition metals is more complex and involves a shift in relative energies.
Understanding Electron Orbital Filling
The general rule for filling electron orbitals, known as the Aufbau principle, suggests that electrons occupy the lowest energy orbitals first. This rule, when simplified, often leads to the (n+l) rule, where n is the principal quantum number and l is the azimuthal quantum number.
- For 4s: n=4, l=0, so n+l = 4
- For 3d: n=3, l=2, so n+l = 5
Based on this simplified (n+l) rule, 4s (n+l=4) appears to have a lower energy than 3d (n+l=5), suggesting that 4s would fill first. This is indeed observed for elements like Potassium (K) and Calcium (Ca), where their electron configurations are [Ar] 4s¹ and [Ar] 4s², respectively.
The Energetic Shift in Transition Metals
However, this relative energy order changes as we move into the transition metals. For elements starting from Scandium (Sc) and continuing through the rest of the 3d series, the scenario shifts significantly.
- For elements from Scandium (Sc) onwards, the 3d orbitals are actually lower in energy than the 4s orbital. This means that for these elements, electrons will preferentially enter the 3d orbitals.
This energetic reversal is due to the complex interplay of factors such as:
- Effective Nuclear Charge: As the atomic number increases across the transition series, the increasing positive charge of the nucleus pulls on the inner electrons more strongly.
- Electron-Electron Repulsions and Shielding: The presence of electrons in the 3d orbitals causes changes in the shielding experienced by both 3d and 4s electrons. The 3d orbitals, being less penetrating, experience less shielding from inner electrons, leading to a stronger attraction to the nucleus and thus a lower energy. The 4s orbitals, while having a lower principal quantum number, become more diffuse and experience more electron-electron repulsion, pushing their energy higher.
This means that while 4s might initially be lower in energy for the first few elements, once you begin adding electrons to the 3d orbitals, the 3d orbitals contract and become more stable, effectively dropping below 4s in energy.
Implications for Ion Formation and Notation
This energy shift has important consequences:
- Ion Formation: When transition metals form cations (positive ions), electrons are typically lost from the outermost shell first. Despite 3d being lower in energy in the neutral atom for transition metals, the 4s electrons are removed before 3d electrons because the 4s orbital becomes higher in energy and is spatially further out from the nucleus in the ion. For example, when Iron (Fe) forms Fe²⁺, the two 4s electrons are lost first.
- Fe: [Ar] 3d⁶ 4s²
- Fe²⁺: [Ar] 3d⁶ (the 4s electrons are removed)
- Electron Configuration Notation: In many written electron configurations, particularly for transition metals, the 3d orbitals are listed before the 4s orbitals (e.g., [Ar] 3dⁿ 4s²) even if the 4s might have filled initially. This is a convention for grouping orbitals by their principal quantum number for easier readability and to reflect their role as valence electrons.
In summary, while the simplified Aufbau principle suggests 4s fills before 3d, the more accurate picture for transition metals shows that the 3d orbitals become energetically lower than 4s orbitals due to complex electronic interactions, leading to their preferential filling and their critical role in the chemistry of these elements.
