Chapter 1-7: Electron Configurations, Periodic Trends, and Quantum Rules

Flashcards covering key vocabulary related to electron configurations, quantum rules (Aufbau, Pauli, Hund), periodic trends (atomic size, ionization energy, metallic character), and related concepts from the lecture notes.

Shorthand configuration

A method of writing electron configurations using noble gas symbols in square brackets to represent core electrons, followed by the remaining valence electrons.

Electron configuration exceptions (Cr, Mo, Cu, Ag, Au)

Elements like chromium, molybdenum, copper, silver, and gold exhibit exceptions where an electron from an 's' orbital is promoted to a 'd' orbital to achieve greater stability (e.g., half-filled or fully-filled d orbitals).

Half-filled d orbital (stability)

An electron configuration where the 'd' orbital is exactly half-filled (d5) provides additional stability due to all spins being aligned and lower total energy.

Fully-filled d orbital (stability)

An electron configuration where the 'd' orbital is fully-filled (d10) provides additional stability.

Shielding

The blocking effect of inner-shell electrons that reduces the positive charge felt by outer-shell electrons from the nucleus.

Effective nuclear charge (Z_effective)

The net positive charge experienced by an electron in an atom, calculated as the number of protons minus the number of core electrons.

Atomic size trend (across a period)

Atomic size generally decreases across a period from left to right because the effective nuclear charge increases, pulling valence electrons closer to the nucleus.

Transition elements (Z_effective & radius)

Across transition elements, the effective nuclear charge and atomic radii remain relatively stable because added electrons go into a core 'd' shell, increasing both protons and shielding electrons proportionally.

Aufbau principle

States that electrons fill atomic orbitals of the lowest available energy levels before occupying higher energy levels.

Pauli exclusion principle

States that no two electrons in the same atom can have the same four quantum numbers, implying that if two electrons are in the same orbital, they must have opposite spins.

Hund's rule

States that for degenerate orbitals (orbitals of the same energy), electrons will fill them singly first with parallel spins before pairing up.

Degenerate orbitals

Orbitals that have the same energy level.

Orbital diagram

A visual representation of electron configuration where each orbital is shown as a box and electrons are shown as arrows, illustrating their spin.

Paramagnetic

Describes substances with unpaired electrons that are attracted to a magnetic field.

Diamagnetic

Describes substances where all electrons are paired, leading to spins cancelling out, and are slightly repelled by a magnetic field.

Ground state (electron configuration)

The non-excited state of an atom where electrons fill orbitals from lowest to highest energy according to the Aufbau principle.

Atomic radius trend (down a group)

Atomic size generally increases down a group because new shells of electrons are added, placing outer electrons further from the nucleus.

Ionization energy

The energy required to remove an electron from an atom or ion.

Ionization energy trend

Ionization energy generally increases towards the top right of the periodic table (e.g., towards fluorine) because atoms are smaller and hold onto their electrons more tightly.

Metallic character

Describes the degree to which an element exhibits the properties of a metal (e.g., electrical and thermal conductivity, malleability, shininess).

Metallic character trend

Metallic character generally increases towards the bottom left of the periodic table (e.g., towards francium).

First ionization energy

The energy required to remove the first electron from a neutral atom.

Ionization energy (valence vs. core)

It is much easier to remove valence electrons than core electrons, resulting in a large jump in ionization energy when attempting to remove a core electron.