Electronic Structure of Atom, Periodic Table and Periodicity

A set of vocabulary flashcards summarising fundamental terms and principles related to atomic structure, quantum mechanics, electron configurations, and periodic trends presented in the lecture.

Atomic Number (Z)

Number of protons in the nucleus of an atom; identifies the element.

Mass Number (A)

Total number of protons and neutrons in an atomic nucleus.

Isotope

Atoms of the same element with identical proton numbers but different neutron numbers (different mass numbers).

Dalton’s Atomic Theory

Early theory stating that elements are composed of indestructible, identical atoms that combine in fixed ratios to form compounds.

Proton

Positively charged sub-atomic particle with relative mass 1 amu, located in the nucleus.

Neutron

Neutral sub-atomic particle with relative mass 1 amu, located in the nucleus.

Electron

Negatively charged sub-atomic particle with relative mass 0.0005 amu, occupying orbitals around the nucleus.

Standard Nuclear Notation

Symbolic representation of an atom: ⁴²₁₀Ne (A over Z next to element symbol).

Bohr Model

Planetary model where electrons move in fixed circular orbits with quantised energies around the nucleus.

Ground State

Lowest energy state of an atom; electrons occupy the lowest available energy levels.

Excited State

Higher energy state achieved when electrons absorb energy and jump to higher levels.

Quantised Energy Levels

Discrete energy values allowed for electrons in an atom; no intermediate energies exist.

de Broglie’s Postulate

Matter (e.g., electrons) exhibits wave–particle duality with wavelength λ = h/p.

Heisenberg Uncertainty Principle

It is impossible to know simultaneously and exactly both the position and momentum of a moving particle.

Schrödinger Equation

Wave equation whose solutions (wavefunctions) describe the probability distribution of electrons in atoms.

Orbital

Region in space around the nucleus with >90 % probability of finding an electron.

Principal Quantum Number (n)

Indicates main energy level (shell) and relative size of the orbital; n = 1, 2, 3…

Angular Momentum Quantum Number (l)

Determines subshell and orbital shape; l = 0 … (n − 1).

Magnetic Quantum Number (mₗ)

Specifies orientation of an orbital in space; values −l … 0 … +l.

Electron Spin Quantum Number (mₛ)

Shows spin of an electron, +½ or −½; two opposite spins allowed per orbital.

s Orbital

Spherically shaped orbital; l = 0; one orientation.

p Orbital

Dumbbell-shaped orbital; l = 1; three orientations (pₓ, pᵧ, p_z).

d Orbital

Cloverleaf-shaped orbital; l = 2; five orientations.

f Orbital

Complex-shaped orbital; l = 3; seven orientations.

Aufbau Principle

Electrons fill orbitals of lowest available energy before occupying higher levels.

Hund’s Rule

Electrons singly occupy degenerate orbitals with parallel spins before pairing occurs.

Pauli Exclusion Principle

No two electrons in an atom can have the same set of four quantum numbers; max two electrons per orbital with opposite spins.

Electron Configuration

Arrangement of electrons in orbitals of an atom or ion (e.g., 1s² 2s² 2p⁶).

Condensed Electron Configuration

Electron configuration written using the previous noble gas core in brackets (e.g., [Ne]3s²).

Valence Electrons

Electrons in the outermost (highest n) energy level; responsible for chemical properties.

Core Electrons

Inner electrons not involved in bonding; shield valence electrons from the nucleus.

Effective Nuclear Charge (Z_eff)

Net positive charge experienced by valence electrons: Z_eff = Z − core electrons.

Shielding Effect

Reduction of attractive force between nucleus and valence electrons due to repulsion by inner-shell electrons.

Atomic Radius

Half the distance between nuclei of identical bonded atoms; indicates atomic size.

Ionic Radius

Radius of a cation or anion in a crystal lattice; varies with charge and electron count.

Cation

Positively charged ion formed when an atom loses one or more electrons; smaller than parent atom.

Anion

Negatively charged ion formed when an atom gains electrons; larger than parent atom.

Isoelectronic Species

Atoms/ions having the same number of electrons and identical electron configurations.

Ionisation Energy (IE)

Minimum energy required to remove 1 mol of electrons from 1 mol of gaseous atoms/ions.

First Ionisation Energy

Energy needed to remove the first electron from a neutral gaseous atom.

Successive Ionisation Energies

Energies required to remove additional electrons after the first; each successive IE is larger.

Electron Affinity (EA)

Energy change when 1 mol of electrons is added to 1 mol of gaseous atoms, forming anions.

Electronegativity (EN)

Ability of an atom in a covalent bond to attract the shared electrons toward itself.

Periodicity

Recurring trends in element properties across periods and groups due to repeating valence electron patterns.

Period (Periodic Table)

Horizontal row of elements sharing the same highest principal quantum number for valence electrons.

Group (Periodic Table)

Vertical column of elements with the same number of valence electrons and similar properties.

s Block

Area of periodic table where outer electrons enter s orbitals (Groups 1 & 2 plus He).

p Block

Area where outer electrons enter p orbitals (Groups 13–18).

d Block

Transition metals where outer electrons fill (n − 1)d orbitals.

f Block

Lanthanides and actinides where electrons fill (n − 2)f orbitals.

Transition Elements

d-block metals with partially filled d subshells; exhibit variable oxidation states.

Anomalous Electron Configuration

Irregular arrangement (e.g., Cr [Ar]4s¹3d⁵) that gains stability from half-filled or filled subshells.

Noble Gas Configuration

Stable electron arrangement with filled s and p subshells in outer shell (octet).

Halogens

Group 17 non-metals that typically gain one electron to form 1⁻ anions.

Alkali Metals

Group 1 metals that lose one electron to form 1⁺ cations.

Alkaline Earth Metals

Group 2 metals that lose two electrons to form 2⁺ cations.