Chapter 5: Electronic Structure and Periodic Properties

Flashcards covering atomic structure, electron configurations, quantum principles, periodic trends, and ionic bonding energetics from Chapter 5.

Ground State

The lowest energy electronic state of an atom or ion, in which electrons occupy the lowest available energy orbitals.

Excited State

An unstable, higher-energy state produced when an atom absorbs a photon or energy, causing an electron to transfer to an orbital that is larger and higher in energy.

Ionization Energy ($IE$)

The minimum energy required to completely remove one electron from an isolated atom or ion in the gas phase (e.g., $\text{H} \rightarrow \text{H}^+ + e^-$, where $IE_{\text{H}} = 2.18 \times 10^{-18} \,\text{J}$).

Screening (Shielding)

The electrostatic repulsion exerted by inner and same-shell electrons that partially cancels the attractive nuclear charge felt by an outer electron.

Effective Nuclear Charge ($Z_{\text{eff}}$)

The net positive nuclear charge experienced by an electron in a multielectron atom after accounting for the screening effects of other electrons.

Photoelectron Spectroscopy (PES)

A technique that applies the photoelectric effect to gas-phase atoms or molecules by using high-energy light to eject electrons and measure their kinetic energy to determine orbital binding energies.

Mendeleev's Periodic Table

An arrangement of elements in order of increasing atomic mass, grouping elements with similar chemical and physical properties into columns while leaving gaps for undiscovered elements.

Modern Periodic Table

An arrangement of elements in order of increasing atomic number ($Z$), in which elements in the same vertical column share similar chemical and physical properties.

Metals

A class of elements characterized by being shiny, solid at room temperature (with the exception of liquid mercury), malleable, ductile, and good conductors of heat and electricity.

Nonmetals

A class of elements that exist as brittle solids, liquids, or gases, and are nonconductors of heat and electricity.

Metalloids

Elements possessing intermediate physical and chemical properties between metals and nonmetals; they are typically shiny like metals, brittle like nonmetals, and act as semiconductors.

Pauli Exclusion Principle

A quantum mechanical principle stating that no two electrons in an atom can have the same four quantum numbers ($n$, $l$, $m_l$, and $m_s$), limiting any orbital to a maximum of two electrons with opposite spins ($m_s = +\frac{1}{2}$ and $m_s = -\frac{1}{2}$).

Aufbau Principle

The principle stating that electrons are added to atomic orbitals starting at the lowest available energy orbital and filling upward to higher energy levels.

Orbital Filling Order (Diagonal Rule)

The relative energy sequence by which atomic subshells fill with electrons: $1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d \rightarrow 4p \rightarrow 5s \rightarrow 4d \rightarrow 5p \rightarrow 6s \rightarrow 4f \rightarrow 5d \rightarrow 6p \rightarrow 7s \rightarrow 5f \rightarrow 6d \rightarrow 7p$.

Hund's Rule

The rule dictating that the most stable arrangement of electrons within a subshell is that with the maximum number of parallel spins; orbitals of equal energy must each be singly occupied before pairing occurs.

Electron Configuration

A notation showing the distribution of electrons among an atom's or ion's various atomic orbitals, written in the general form $n l^{\text{number of electrons}}$.

Valence Electrons

Electrons residing in the shell with the highest principal quantum number ($n$) that take part in chemical bonding and chemical reactions.

Core Electrons

Inner electrons that possess principal quantum numbers ($n$) lower than the valence shell and do not actively participate in chemical reactions.

Noble Gas Configuration

A compact shorthand notation in which the symbol of the preceding noble gas is enclosed in brackets to represent the core electrons, followed by the remaining valence electrons (e.g., $\text{Na}$ as $[\text{Ne}]\,3s^1$).

Aufbau Exceptions for $\text{Cr}$ and $\text{Cu}$

Anomalies in orbital filling where an electron promotes from $4s$ to $3d$ to achieve the added thermodynamic stability of a half-filled ($\text{Cr}: [\text{Ar}]\,4s^1 3d^5$) or fully filled ($\text{Cu}: [\text{Ar}]\,4s^1 3d^{10}$) $d$-subshell.

Near-Degenerate Orbitals

Orbitals belonging to different subshells that possess nearly identical energy levels, commonly leading to electron configuration irregularities in elements with atomic numbers $Z > 40$.

Isoelectronic

Refers to atoms, cations, or anions that possess the exact same total number of electrons and identical ground-state electron configurations (e.g., $\text{Na}^+$, $\text{Al}^{3+}$, $\text{F}^-$, and $\text{N}^{3-}$ are isoelectronic with $\text{Ne}$).

Paramagnetism

A form of magnetism occurring in substances with one or more unpaired electrons, causing the substance to be attracted into a magnetic field.

Diamagnetism

A form of magnetism occurring in substances where all electrons are paired with opposing spins, causing the substance to be weakly repelled by a magnetic field.

Ferromagnetic

Substances that exhibit unusually strong, permanent paramagnetic properties.

Atomic Radius

The distance between an atom's nucleus and the outermost boundary of its largest electron orbital; decreases across a period and increases down a group.

Cation Radius Trend

The property that a cation is always strictly smaller than the neutral atom from which it forms, due to the loss of outer electrons and decreased electron-electron repulsions.

Anion Radius Trend

The property that an anion is always strictly larger than the neutral atom from which it forms, caused by increased electron-electron repulsions when additional electrons are gained.

First Ionization Energy Irregularity (Groups 2A to 3A)

The decrease in first ionization energy from magnesium ($738 \,\text{kJ/mol}$) to aluminum ($577 \,\text{kJ/mol}$) resulting from the removal of an electron from the higher-energy $3p$ orbital in aluminum rather than the fully filled, well-shielding $3s$ orbital.

First Ionization Energy Irregularity (Groups 5A to 6A)

The decrease in first ionization energy from nitrogen ($1402 \,\text{kJ/mol}$) to oxygen ($1314 \,\text{kJ/mol}$) caused by the removal of a paired electron from the doubly occupied $2p$ orbital in oxygen, which experiences extra electron-electron repulsion.

Successive Ionization Energies

The sequential removal energies ($IE_1$, $IE_2$, $IE_3$, etc.) for electrons in an atom, exhibiting an enormous increase when electron removal shifts from the valence shell to inner core orbitals.

Electron Affinity ($EA$)

The energy change ($\Delta H$) accompanying the addition of an electron to an isolated gas-phase atom (e.g., $\text{Cl}(g) + e^- \rightarrow \text{Cl}^-(g)$, where $\Delta H = -349 \,\text{kJ/mol}$).

Lattice Energy

The energy required to completely dissociate one mole of a crystalline solid ionic compound into its separated gas-phase ions (e.g., $\text{NaCl}(s) \rightarrow \text{Na}^+(g) + \text{Cl}^-(g)$).

Born-Haber Cycle

A thermochemical cycle applying Hess's law to determine the lattice energy of an ionic compound through individual steps: sublimation, bond dissociation, ionization, electron affinity, and lattice formation.