Periodic Trends and Electron Configurations Vocabulary

Vocabulary flashcards covering atomic trends, polyelectronic quantum mechanical principles, orbital rules, and exception cases.

Dmitri Mendeleev

Russian scientist who in 1869 published the first periodic table by arranging elements according to their physical and chemical properties, successfully predicting missing elements before atomic numbers were known.

Henry Moseley

Physicist who in 1913 reorganized the periodic table by increasing atomic number ($\text{number of protons}$) rather than by atomic weight or general properties alone.

Heisenberg's Uncertainty Principle

A fundamental quantum principle defined by $\Delta x \times \text{uncertainty in momentum} \text{ } \ge \frac{h}{4\pi}$, showing that position and velocity cannot be simultaneously known with absolute precision.

Three-Body Problem

A mathematical limitation establishing that Schrödinger's equation cannot be solved exactly for any atomic system containing three or more moving, interacting particles, such as polyelectronic atoms.

Electron Correlation Problem

The mathematical inability to calculate exact electrostatic repulsions between electrons because their specific positions relative to one another cannot be precisely determined.

Shielding

The process by which inner or neighboring electrons screen outer electrons from experiencing the full positive attraction of the nuclear charge.

Penetration

The ability of an electron to get closer to the nucleus, which breaks degeneracy within a principal quantum level $n$ such that energies follow $E_{ns} < E_{np} < E_{nd} < E_{nf}$.

Degenerate Orbitals

A set of atomic orbitals that possess identical energy levels, such as the three orbitals in a $p$ subshell or five orbitals in a $d$ subshell.

Pauli Exclusion Principle

The rule stating that no two electrons in an atom can have the exact same set of four quantum numbers ($n$, $l$, $m_l$, $m_s$) in their ground state.

Aufbau Principle

The principle (meaning "to build up" in German) dictating that electrons fill the lowest energy available atomic orbitals first when constructing an atom's ground state electron configuration.

Hund's Rule

The rule specifying that the lowest energy configuration for degenerate orbitals is the one with the maximum number of unpaired electrons, all having parallel spins.

Valence Electrons

The electrons located in the outermost principal quantum level $n$ of a main group atom that are involved in forming chemical bonds.

Core Electrons

All inner electrons in an atom that do not occupy the outermost principal quantum level $n$ and do not directly participate in bonding.

Chromium Electron Configuration Exception

An anomalous ground state configuration given by $[\text{Ar}]\,4s^1\,3d^5$ instead of $[\text{Ar}]\,4s^2\,3d^4$, driven by the stability of two half-filled subshells.

Copper Electron Configuration Exception

An anomalous ground state configuration given by $[\text{Ar}]\,4s^1\,3d^{10}$ instead of $[\text{Ar}]\,4s^2\,3d^9$, driven by the stability of a fully filled $d$ subshell.

Atomic Radius

A measure of the size of an atom, which increases down a group as $n$ increases and generally decreases from left to right across a period as nuclear charge increases.

Ionization Energy

The minimum energy required to completely remove an electron from a gaseous atom or ion, represented by $X_{(g)} \rightarrow X^{+}_{(g)} + e^{-}$.

Electron Affinity

The energy change associated with adding an electron to a gaseous atom, represented by $X_{(g)} + e^{-} \rightarrow X^{-}_{(g)}$.

Fluorine vs. Chlorine Electron Affinity Exception

The anomaly where Fluorine ($-327.8\,kJ\,mol^{-1}$) has a less negative electron affinity than Chlorine ($-348.7\,kJ\,mol^{-1}$) due to electron repulsions in Fluorine's smaller $n=2$ orbital space.

Isoelectronic Ions

A series of distinct ions and atoms that contain the exact same total number of electrons, such as $N^{3-}$, $O^{2-}$, $F^{-}$, $Ne$, $Na^{+}$, $Mg^{2+}$, and $Al^{3+}$.