The periodic table organizes elements into s-, p-, d-, and f-blocks, each corresponding to the type of atomic orbital being filled with electrons. A common observation is that d-block elements, typically associated with the third principal electron shell (n=3), unexpectedly begin in the fourth row, or period 4, of the periodic table. This arrangement raises a question about the order in which electrons fill atomic orbitals.
Building Blocks of Atoms
Electrons occupy specific regions around an atom’s nucleus called electron shells, designated by principal quantum numbers (n = 1, 2, 3, etc.). Higher numbers indicate shells further from the nucleus with higher energy. Each principal shell can hold a maximum number of electrons.
Within these shells are subshells (s, p, d, f). Each subshell contains a specific number of orbitals, which are three-dimensional probability distributions where electrons are likely to be found. An s subshell has one orbital (2 electrons), a p subshell has three (6 electrons), a d subshell has five (10 electrons), and an f subshell has seven (14 electrons). Electrons fill these orbitals from lower to higher energy.
The Unexpected Energy Order
Electrons fill atomic orbitals following the Aufbau principle, which states that electrons occupy the lowest available energy levels first. This principle suggests a filling order based on increasing principal quantum number (n). However, a notable deviation occurs with the 4s and 3d orbitals. Despite the 4s orbital belonging to the fourth principal shell (n=4) and the 3d orbital belonging to the third (n=3), the 4s orbital is generally filled before the 3d orbital for neutral atoms.
This energy difference results from electron-electron repulsion and nuclear shielding. While 3d orbitals are spatially closer to the nucleus than 4s orbitals, the 4s orbital’s spherical shape allows its electrons to penetrate closer to the nucleus. This penetration leads to a stronger effective nuclear charge and less shielding for 4s electrons, making the 4s orbital slightly lower in energy than the 3d orbital in neutral atoms.
As a result of this energy ordering, the 4s orbital fills with two electrons before the 3d orbitals. Once the 4s subshell is complete, electrons then enter the 3d orbitals. This sequential filling marks the beginning of the d-block elements in the fourth row, spanning from scandium (Sc) to zinc (Zn). D-block elements represent a transition where inner (n-1)d orbitals are filled after the outermost ns orbital.
Properties of D-Block Elements
D-block elements, also known as transition metals, are positioned in the middle of the periodic table, between the s-block and p-block. Their properties arise from having partially filled d-orbitals, which enables several distinct behaviors.
They display variable oxidation states, forming ions with multiple charges (e.g., iron as Fe²⁺ and Fe³⁺). This occurs because the small energy difference between their outermost s and inner d electrons allows both to participate in chemical bonding.
Many d-block elements also form colored compounds. Partially filled d-orbitals allow electrons to absorb specific wavelengths of visible light and transition to higher energy d-orbitals. When these excited electrons return to lower levels, they emit light in the visible spectrum, creating characteristic colors.
Transition metals are also known for their catalytic activity, accelerating chemical reactions without being consumed. This capability arises from their multiple oxidation states and ability to form temporary bonds with reactants. They facilitate industrial processes by providing alternative, lower-energy reaction pathways, such as iron in ammonia production.