Chem. 1A Week 10 Discussion Notes Dr. Mack/S12 Electron Configurations: Quantum Numbers: The location of an electron in an atom is defined by a probability relationship that is the solution to the Schrödinger equation. Schrödinger asserted that the wavelike nature of particles requires that matter (electrons in this case) be completely described mathematically by a "Wave Function". A series parameters (variables) arise out of the mathematics that govern these wave functions which we call orbitals. These parameters are known as "quantum numbers" as they describe the electrons quantized behavior in terms of energy and location about the nucleus. "n" defines the "shell" in which an orbital resides, "l" defines the type of orbital (subshell) and "ml" defines the orbital's spatial orientation. l = 0 are designated as "s" orbitals, l = 1 are designated as "p" orbitals l = 2 are designated as "d" orbitals l = 3 are designated as "f" orbitals Within each orbital there may reside two electrons, one "spin up" and one spin down. The spin states are designated by "ms" = ± ½ • Electrons fill into the orbitals by the "Aufbau" principle (bottom up) according to the periodic table. The orbitals are arranged from lowest to highest energy by n + l. • Electrons will fill open degenerate orbitals until forced to pair by "Hund's Rule" • When electrons pair in an orbital, they must have opposite spin to comply with the "Pauli Principle" Each individual orbital can hold a maximum of two electrons. The s−orbital can hold 2 electrons, the set of three p−orbitals can hold 6, the five d−orbitals can hold 10 and the seven f−orbitals can hold 14 total electrons. Filling of Electrons: Starting with hydrogen and moving across the periodic table to helium and down to lithium and across and downward successively, the electrons fill by: 1s1 1s2 2s1 2s2 2p1 2p2…2p6 3s1 3s2 … An electron configuration lists the number of electrons in each set of orbitals: 1s2 2s2 2p6 3s2 … Example: Page 1 of 2
Chem. 1A Week 10 Discussion Notes Dr. Mack/S12 Transition Elements and Beyond: Electrons fill in spdf notation according to the periodic table: Mn: 1s22s22p63s23p64s23d5 Electrons configurations are written by shell. Mn: 1s22s22p63s23p63d54s2 (your text does not do this, however this is the correct form) Electron configurations can be written in a short hand form using "core" notation. The core electrons are replaced by the electron configuration of the nearest lower noble gas. The electron configuration of cations are written by removing electrons from the outermost shell inward. Anions fill from the last open orbital outward. Due to the closeness of the "s" and "d" orbitals, it is experimentally observed that in nd4 and nd9 an "s" electron will promote itself to half fill or fill the d−orbital. Cr: [Ar]3d44s2 becomes Cr: [Ar]3d54s1 Atoms or ions with unpaired electrons are said to be Paramagnetic, Atoms or ions with no unpaired electrons are said to be Diamagnetic. Formation of Ions: Cations are formed as successive electrons are removed from the outermost shells inward. Paramagnetic Diamagnetic Anions conversely fill from the lowest unfilled shell outward. Page 2 of 2