d orbital

Meaning: A d orbital is one of the orbitals in a subshell with l = 2. The d subshell first appears at n = 3, so there is no 1d or 2d. It has five orbitals (ml = −2, −1, 0, +1, +2) of equal energy in an isolated atom, holding up to 10 electrons. Four have a four-lobed cloverleaf shape; the fifth (dz²) has two lobes and a ring. Transition metals are the elements filling the (n−1)d subshell.

Equation: Number of orbitals = 2l + 1 = 5, and maximum electrons = 2(2l + 1) = 10, where l is the angular momentum quantum number and ml is the magnetic quantum number (both dimensionless).

On the MCAT: You write transition metal configurations, such as Fe: [Ar] 4s² 3d⁶. For cations, remove the 4s electrons first, so Fe²⁺ is [Ar] 3d⁶. Expect the Cr and Cu exceptions, [Ar] 4s¹ 3d⁵ and [Ar] 4s¹ 3d¹⁰, which are usually explained by the stability of half-filled and filled d subshells. Counting unpaired d electrons (Hund's rule) predicts paramagnetism. In coordination complexes, ligands split the d orbitals into two energy sets, and electron jumps between them often absorb visible light, giving colored compounds.

Common confusion: Students assume the highest-energy electrons are the last ones filled, so they remove 3d electrons first when forming a cation. Although 4s fills before 3d, once 3d is occupied the 4s electrons are the outermost and are lost first.