Atomic Structure, Periodic Properties, and Chemical Bonding Vocabulary

Comprehensive vocabulary flashcards covering subatomic particles, atomic models, quantum theory, periodic properties, and chemical bonding principles from lecture notes.

Cathode Rays

Streams of negatively charged particles (electrons) moving from cathode to anode inside a gas discharge tube, independent of the electrode material or gas present.

Charge-to-Mass Ratio of Electron ($e/m_e$)

The ratio of the charge of an electron to its mass, determined by J.J. Thomson by measuring the deflection of cathode rays in electrical and magnetic fields, equal to $1.75 \times 10^{11}\,\text{C/kg}$.

Charge of an Electron ($e$)

The absolute electric charge carried by a single electron, determined by Millikan using the oil drop experiment to be $1.6022 \times 10^{-19}\,\text{C}$.

Mass of an Electron ($m_e$)

The rest mass of an electron, calculated from its charge and charge-to-mass ratio to be $9.10939 \times 10^{-31}\,\text{kg}$ (or $0.00054\,\text{amu}$).

Discovery of Neutrons

Discovered by James Chadwick by bombarding a thin sheet of beryllium with alpha rays, yielding neutral subatomic particles according to ${}_4^9\text{Be} + {}_2^4\text{He} \rightarrow {}_0^1\text{n} + {}_6^{12}\text{C}$.

Atomic Number ($Z$)

The total number of protons present in the nucleus of an atom, which equals the number of electrons in a neutral atom.

Mass Number ($A$)

The sum of the total number of protons and neutrons (nucleons) present in the nucleus of an atom.

Isotopes

Atoms of the same element having the same atomic number ($Z$) but different mass numbers ($A$), such as Protium (${}_1^1\text{H}$), Deuterium (${}_1^2\text{H}$), and Tritium (${}_1^3\text{H}$).

Isobars

Atoms of different chemical elements having the same mass number ($A$) but different atomic numbers ($Z$), such as ${}_6^{14}\text{C}$ and ${}_7^{14}\text{N}$.

Thomson Atomic Theory

Also known as the Plum Pudding Model, it proposes that an atom is spherical in shape with positive charge uniformly distributed and electrons embedded within it.

Rutherford's Alpha-Ray Scattering Experiment

An experiment where Rutherford bombarded thin gold foil with radioactive $\alpha$-particles, detecting deflections using a fluorescent zinc sulphide ($\text{ZnS}$) screen.

Nucleus

The extremely small, dense region at the center of an atom where all positive charge and almost the entire mass of the atom are concentrated.

Threshold Frequency ($\nu_0$)

The minimum frequency of incident light required by a given metal surface to exhibit the photoelectric effect and eject electrons.

Planck's Quantum Theory

Theory stating that electromagnetic energy is absorbed or emitted discontinuously in discrete energy packets called quanta or photons, with energy $E = h\nu$.

Photoelectric Effect

The ejection of electrons from a metal surface when light of suitable frequency hits it, governed by $h\nu = h\nu_0 + \frac{1}{2} m_e v^2$.

Black Body Radiation

The radiation emitted or absorbed by an ideal body that absorbs and emits radiation of all frequencies completely.

Electromagnetic Spectrum

The systematic arrangement of all types of electromagnetic radiation in increasing order of wavelength or decreasing order of frequency.

Rydberg Equation

An equation used to calculate the frequency or wave number of spectral lines in the hydrogen emission spectrum: $\bar{\nu} = \frac{1}{\lambda} = R \left( \frac{1}{n_1^2} - \frac{1}{n_2^2} \right)$, where $R = 1.097 \times 10^7\,\text{m}^{-1}$.

Wavelength ($\lambda$)

The physical distance between two adjacent crests or two adjacent troughs of a wave.

Wavenumber ($\bar{\nu}$)

The reciprocal of wavelength ($\bar{\nu} = \frac{1}{\lambda}$), representing the number of wavelengths per unit length.

Principal Quantum Number ($n$)

Quantum number specifying the main energy level or shell occupied by an electron, dictating overall size and maximum electron capacity ($2n^2$).

Azimuthal Quantum Number ($l$)

Also called angular momentum quantum number, it defines subshells ($s, p, d, f$) and orbital shape, taking integer values from $0$ to $(n-1)$.

Magnetic Quantum Number ($m$)

Quantum number representing the spatial orientation of an orbital relative to an applied magnetic field, taking values from $-l$ to $+l$.

Spin Quantum Number ($s$)

Quantum number describing the intrinsic spin orientation of an electron about its own axis, equal to $+\frac{1}{2}$ or $-\frac{1}{2}$.

de Broglie Relation

Formula expressing the dual wave-particle nature of matter, given by $\lambda = \frac{h}{p} = \frac{h}{m v}$.

Heisenberg's Uncertainty Principle

Principle asserting that it is impossible to simultaneously measure both the position and momentum of a subatomic particle accurately ($\Delta x \cdot \Delta p \ge \frac{h}{4\pi}$).

Orbital

A three-dimensional space surrounding the nucleus where the probability of finding an electron is maximum.

Zeeman Effect

The phenomenon of splitting of spectral lines of an atom when exposed to an external magnetic field.

Stark Effect

The phenomenon of splitting of spectral lines of an atom when exposed to an external electric field.

Aufbau Principle

Rule stating that electrons fill atomic orbitals in order of increasing orbital energy level in the ground state.

Pauli Exclusion Principle

Principle stating that no two electrons in the same atom can have an identical set of all four quantum numbers.

Hund's Rule of Maximum Multiplicity

Rule stating that electron pairing in degenerate orbitals does not take place until every orbital in the subshell is singly occupied.

$(n+l)$ Rule

Rule determining orbital energy: an orbital with a lower $(n+l)$ value possesses lower energy; if $(n+l)$ values are identical, the orbital with lower $n$ has lower energy.

Dobereiner's Law of Triads

States that when elements with similar properties are grouped in sets of three, the atomic mass of the middle element is approximately the average of the first and third.

Newlands' Law of Octaves

States that when elements are arranged in increasing order of atomic mass, every eighth element exhibits properties similar to the first.

Modern Periodic Law

States that physical and chemical properties of elements are periodic functions of their atomic numbers.

Representative Elements

The s-block and p-block elements in the periodic table, excluding the noble gases.

Transition Elements

The d-block elements whose properties represent a transition between metallic s-block elements and non-metallic p-block elements.

Diagonal Relationship

The similarity in chemical behavior between a second-period element and a third-period element in the adjacent group, such as $\text{Li}$ and $\text{Mg}$, or $\text{Be}$ and $\text{Al}$.

Anomalous Properties

The unique behavior displayed by the first element of each main group compared to heavier elements in the same group due to tiny size and high electronegativity.

Atomic Radius

The distance from the center of the atomic nucleus to the outermost electron shell.

Ionization Enthalpy

The minimum energy required to remove an electron from the outermost shell of an isolated neutral gaseous atom in its ground state.

Electron Gain Enthalpy

The energy released or absorbed when an electron is added to an isolated neutral gaseous atom to form a anion.

Electronegativity

The tendency of an atom in a covalent molecule to attract shared electron pairs toward itself.

Isoelectronic Species

Atoms or ions that contain the exact same number of electrons despite having different nuclear charges (such as $\text{Na}^+$, $\text{F}^-$, $\text{Mg}^{2+}$).

Ionic Bond

A chemical bond formed by complete electron transfer resulting in electrostatic attraction between oppositely charged ions.

Covalent Bond

A chemical bond formed by the mutual sharing of valence electron pairs between two atoms.

Polarization (Fajans' Rule)

The attraction and distortion of an anion's electron cloud by a cation, imparting covalent character to an ionic compound.

Lattice Enthalpy

The energy released when one mole of an ionic crystalline compound is formed from its constituent gaseous ions.

VSEPR Repulsion Order

The relative hierarchy of electron pair repulsions surrounding a central atom: $\text{l.p.}-\text{l.p.} > \text{l.p.}-\text{b.p.} > \text{b.p.}-\text{b.p.}$.

Formal Charge

The calculated individual electrical charge assigned to an atom in a Lewis structure, given by $\text{FC} = \{\text{Valence } e^-\} - \{\text{Lone pair } e^-\} - \frac{1}{2}\{\text{Shared } e^-\}$.

Dipole Moment ($\mu$)

A quantitative measure of molecular polarity defined as the product of charge magnitude ($q$) and distance of charge separation ($d$): $\mu = q \times d$.

Resonance

The representation of a single molecule using two or more canonical Lewis structures to explain its chemical properties and bonding.