Periodic Table, Atomic Theory, and Atomic Structure

Vocabulary practice flashcards covering the historical development of atomic theory, the periodic table, electron configurations, quantum principles, and ion formation.

Electron

A subatomic particle bearing a negative electrical charge and a mass of approximately $9.11 \times 10^{-31}\,\text{kg}$, named after the Greek word for amber.

Plum Pudding Model

Atomic model introduced by J. J. Thomson that conceived the atom as a diffuse sphere of positive electrical charge embedded with discrete negative electrons.

Alpha ($\alpha$) Particle

A small, dense, positively charged particle emitted by radioactive isotopes (such as radium), utilised by Ernest Rutherford to probe the internal structure of atoms.

Geiger-Marsden Experiment (1909)

Scattering experiment in which alpha particles were directed at thin gold foil; large deflections proved the atom has a small, concentrated, positively charged core rather than a diffuse mass.

Atomic Nucleus

The dense central core of an atom (derived from the Latin word for nut or kernel) carrying positive charge and comprising virtually all of the atom's total mass.

Atomic Orbital

A mathematical three-dimensional region of space surrounding an atomic nucleus where there is greater than a $90\%$ probability of finding a specific electron.

Chemical Element

A species of atoms or a pure substance formed entirely of atoms that possess the same atomic number (number of protons in the atomic nucleus).

Essential Biological Elements

The eleven chemical elements required by all biological systems, dominated in percentage of human body atoms by hydrogen ($62.8\%$), oxygen ($25.4\%$), and carbon ($9.4\%$).

Biological Selection Criteria for Elements

Living systems select lightweight elements that are naturally abundant in the environment and present in chemical states easily extractable from the Earth's crust, sea, or atmosphere.

Carbon-Based Life Suitability

Carbon forms the backbone of biological life due to its crustal abundance ($0.18\%$) compared to other light elements, and its capacity to form stable, complex covalent bonds with $O$, $H$, $N$, $P$, and $S$.

Proton

A positively charged subatomic particle situated within the nucleus with a mass of $1\,\text{amu}$ ($\sim 1.67 \times 10^{-27}\,\text{kg}$).

Neutron

An electrically neutral subatomic nucleon with a mass of $1\,\text{amu}$ ($\sim 1.67 \times 10^{-27}\,\text{kg}$).

Atomic Number ($Z$)

The absolute quantity of protons contained inside an atom's nucleus, which defines the chemical identity of that element.

Isotopes

Variants of a given chemical element that possess identical atomic numbers ($Z$) but different numbers of neutrons, leading to different atomic masses.

Relative Atomic Mass ($A_r$)

The weighted average mass of atoms of an element across all naturally occurring isotopes, calculated based on each isotope's relative natural abundance.

Valence Shell

The outermost electron shell of an atom located furthest from the nucleus, containing the valence electrons that dictate the atom's chemical behavior.

External Configuration

The shorthand electronic configuration notation that represents solely the valence electrons positioned within the outermost shell (such as $2s^2 2p^2$ for carbon).

Atomic Subshell

A subdivision of an electron shell comprising equivalent orbitals categorized as $s$ (spherical), $p$ (dumbbell), $d$ (four-leaf clover), and $f$ (raspberry-like).

Degenerate Orbitals

Orbitals belonging to the same subshell that have precisely equal amounts of energy.

Aufbau Principle

The rule dictating that electrons systematically occupy orbitals of lower energy before filling higher-energy orbitals.

Pauli Exclusion Principle

The quantum rule establishing that a single orbital can accommodate no more than two electrons, and those two electrons must have opposite spins.

Hund's Rule

The principle stating that degenerate orbitals are each singly occupied with parallel-spin electrons before any orbital receives a second electron.

Periodic Table Blocks

Sections of the periodic table ($s, p, d, f$) designated according to the specific orbital type that accommodates an element's valence electrons.

Electron Excitation

The process by which an electron absorbs a discrete packet of energy, prompting a jump from a lower orbital to an unoccupied higher-energy orbital.

Electron Relaxation

The transition wherein an excited electron drops back to its lower energy ground state, emitting a photon of electromagnetic radiation corresponding to $\Delta E$.

Planck's Equation

Fundamental equation $E = h\nu = \frac{hc}{\lambda}$, showing that photon energy is directly proportional to frequency ($\nu$) and inversely proportional to wavelength ($\lambda$).

Bohr's Frequency Condition

The relationship stating that the frequency of absorbed or emitted radiation is directly proportional to the difference in energy between an excited state and the ground state ($\Delta E = h\nu$).

Ionisation Energy

The energy (measured in joules or $\text{eV}$) required to eject one mole of electrons from one mole of isolated gaseous atoms or molecules (e.g., $Na(g) \rightarrow Na^+(g) + e^-$ at $496\,kJ\,mol^{-1}$).

Periodic Trends in Ionisation Energy

General elemental trends whereby ionisation energy decreases down a group due to increased atomic size and increases from left to right across a period due to increasing effective nuclear charge.

Cation

A positively charged chemical species generated when a neutral atom loses one or more electrons, resulting in a net excess of protons over electrons.

Anion

A negatively charged chemical species generated when a neutral atom captures one or more electrons, creating an excess of electrons over protons.