Electronic Structure and Periodic Properties of Elements

Comprehensive vocabulary flashcards focusing on electronic structure, the periodic table's organization, periodic trends, and the definitions of ions and chemical bonds based on lecture notes.

Electronic Configuration

A description of how electrons are distributed among the various atomic orbitals, symbolized by the number of the principal quantum shell ($n$), the letter designating the orbital type ($l$), and a superscript number for the electrons in that subshell.

Aufbau Principle

A rule stating that electrons in a ground state atom occupy the lowest energy state first, meaning lower energy orbitals are filled before higher energy ones.

Hund’s Rule

A principle stating that electrons fill degenerate orbitals (those with the same energy) until all are half-filled before any pairing of electrons occurs.

Pauli Exclusion Principle

A rule stating that individual orbitals can hold a maximum of two electrons which must have opposite spins ($+\frac{1}{2}$ or $-\frac{1}{2}$), meaning no two electrons in an atom can share the same four quantum numbers.

Degenerate Orbitals

Orbitals that possess the same energy level.

Orbital Diagrams

Pictorial representations of electron configurations that show individual orbitals and the pairing arrangement of electrons.

Valence Electrons

Electrons that occupy the outermost shell orbital(s) with the highest principal quantum number ($n$) and determine most of an atom's chemical behavior.

Core Electrons

Electrons occupying the inner shell orbitals, corresponding to the electron configuration of the preceding noble gas.

Periods

The horizontal rows on the periodic table; the period number indicates the number of principal shells in the atoms of those elements.

Groups

The vertical columns on the periodic table; the group number identifies the number of valence electrons, and elements within a group share similar chemical properties.

Main Group Elements

Also known as representative elements, these are elements where the last electron added enters an $s$ or a $p$ orbital in the outermost shell.

Transition Elements

Metallic elements where the last electron added enters a $d$ orbital; their valence electrons include the $ns$ and $(n - 1)d$ electrons.

Inner Transition Elements

Metallic elements where the last electron added occupies an $f$ orbital; this includes the lanthanide series and the actinide series.

Cation

A positively charged ion formed by the loss of one or more electrons from the valence shell.

Anion

A negatively charged ion formed by the gain of one or more electrons into the valence shell.

Covalent Radius

Defined as one half the distance between the nuclei of two identical atoms joined by a covalent bond.

Effective Nuclear Charge ($Z_{eff}$)

The actual nuclear charge felt by an electron, represented by the formula $Z_{eff} = Z - \text{shielding}$, where shielding is caused by inner electrons reducing the nucleus's attraction.

Isoelectronic

A term describing atoms and ions that possess the same electron configuration.

First Ionization Energy ($IE_1$)

The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state.

Electron Affinity ($EA$)

The energy change associated with adding an electron to a gaseous atom to form an anion; a negative value indicates energy is released.

Periodic Law

The principle stating that the properties of elements are periodic functions of their atomic numbers.

Ionic Bonds

Chemical bonds formed by the complete transfer of electrons, typically between metals (cations) and non-metals (anions), stabilized by electrostatic attraction.

Covalent Bonds

Chemical bonds formed by the sharing of electrons between non-metals.

Monoatomic Ions

Ions that contain only one atom, such as $Na^+$ or $Cl^-$.

Polyatomic Ions

Ions consisting of more than one atom bonded together carrying an overall net charge, such as $NO_3^-$ or $PO_4^{3-}$.

Oxyanions

Polyatomic ions that contain one or more oxygen atoms.