Atomic Theory, Chemical Bonding, and Molecular Structure

Comprehensive vocabulary flashcards covering atomic theory development, quantum mechanics, electron configuration rules, chemical bonding types, Lewis structures, orbital hybridization, VSEPR geometries, molecular polarity, and intermolecular forces.

Dalton's Atomic Theory (1808)

The atomic model postulating that all matter is composed of very small particles called atoms, which cannot be subdivided or changed into one another.

Cathode Rays

A stream of radiation observed by J.J. Thomson in cathode ray tubes that is common to all matter, deflected by magnetic fields, carries a negative charge, and is composed of electrons.

Plum Pudding Model

J.J. Thomson's model of the atom that attempted to explain atomic neutrality by depicting negatively charged electrons embedded within a positively charged dough or pudding.

Rutherford's Gold Foil Experiment

An experiment where alpha particles were directed at thin gold foil; most particles passed straight through, some were deflected at small or large angles, and a very few bounced straight back.

Rutherford's Scattering Inferences

The deduction that the atom is mostly empty space occupied by electrons, with a dense, positively charged center (the nucleus) responsible for deflecting alpha particles.

Limitations of Rutherford's Model

The model did not adequately describe the location of electrons and conflicted with classical laws of physics, which dictated that orbiting electrons should constantly emit radiation and lose energy.

Continuous Spectrum

An emission spectrum containing all possible frequencies of electromagnetic radiation ($\text{EMR}$) emitted simultaneously, visible as a continuous, blurred rainbow of colour (such as sunlight).

Line Emission Spectrum

A spectrum produced when an electric current excites a specific chemical element, emitting a characteristic pattern of distinct lines of specific colours.

Bohr Atomic Orbits

Fixed circular paths around the nucleus in which electrons move, where an electron possesses a specific, fixed quantum of energy without radiating energy.

Quantized Energy

The principle that an electron in a specific orbit possesses a discrete, fixed quantity of energy at all times, preventing continuous energy loss.

Excited State

An unstable, high-energy state reached when an electron absorbs sufficient energy and jumps from its original orbit to a higher orbit.

Ground State

The lowest, most stable energy level that an electron can occupy within an atom.

Photon Emission (Bohr Model)

The release of energy as a quantum of light when an unstable excited electron drops back to its ground state, exactly equaling the energy difference between the two states.

Wave-Particle Duality

The principle that electrons and other moving particles exhibit both wave-like and particle-like properties, evidenced by de Broglie's wavelength and Thomson's electron interference patterns.

Heisenberg Uncertainty Principle

The principle stating that it is impossible to simultaneously measure both the precise position and momentum of an electron, disproving Bohr's well-defined circular orbits.

Orbital (Electron Cloud)

A three-dimensional probability distribution surrounding the nucleus that indicates the probable location of finding an electron rather than a fixed circular path.

Schrödinger Wave Function

A mathematical formula developed by Erwin Schrödinger whose wave functions determine the probability of finding an electron at a given position in space and time.

Aufbau Principle

The rule dictating that electrons must fill and occupy the atomic orbital of the lowest available energy level first before occupying higher levels.

Pauli Exclusion Principle

The quantum rule stating that no more than $2$ electrons can occupy a single atomic orbital, and they must have opposite spins.

Hund's Rule

The principle stating that when electrons occupy orbitals of equal energy (degenerate orbitals), one electron enters each orbital with parallel spin before any orbital receives a second electron.

Principal Quantum Number ($n$)

The quantum number designating the main energy level or electron shell, where the maximum number of electrons permitted in a shell is calculated as $2n^2$.

Azimuth Quantum Number ($l$)

The quantum number representing orbital shape ($s, p, d, f$), with allowed integer values ranging from $0$ to $n - 1$.

Magnetic Quantum Number ($m_l$)

The quantum number representing the spatial orientation of an orbital around the nucleus, with allowed integer values ranging from $-l$ to $+l$.

Spin Quantum Number ($m_s$)

The quantum number indicating the spin orientation of an electron on its axis, taking a value of either $+\frac{1}{2}$ (up) or $-\frac{1}{2}$ (down).

Wolframite

The mineral ore from which the element tungsten is recovered, giving rise to its chemical symbol $W$.

Ionic Bonding

A type of chemical bond formed between a metal and a non-metal when one atom transfers one or more electrons to another atom to achieve stability.

Covalent Bonding

A type of chemical bond formed when two atoms attain stability by sharing electron pairs between their nuclei.

Electronegativity

A measure of how strongly an atom in a chemical bond attracts and pulls shared electrons toward itself.

Metallic Bonding

A bonding model described as a "sea of electrons," where stationary atomic nuclei share delocalized outer valence electrons that move freely throughout the metal.

Brittleness of Ionic Solids

The mechanical property of ionic crystals that causes them to shatter when struck because shifting lattice layers force ions of like charge adjacent to each other, resulting in repulsive forces.

Expanded Octet

A bonding situation where a central atom accommodates more than $8$ valence electrons, restricted to elements in the third row or lower of the periodic table due to accessible $d$ orbitals.

Coordinate Covalent Bond

A covalent bond in which one participating atom donates both of the shared bonding electrons (such as in the formation of $NH_4^+$).

Resonance Structures

Two or more valid Lewis structures for a single molecule that maintain the same relative nuclear positions but differ in the distribution of bonding electrons and lone pairs.

VSEPR Theory

Valence Shell Electron Pair Repulsion Theory; predicts three-dimensional molecular geometries based on the concept that negatively charged electron pairs repel and remain as far apart as possible.

VSEPR Pair

A unit of electron repulsion around a central atom; every bond (single, double, or triple) and every lone pair counts as one VSEPR pair.

Sigma ($\sigma$) Bond

A covalent bond formed by the in-phase constructive interference and direct end-to-end overlap of atomic orbitals along the internuclear axis.

Hybrid Orbitals

New atomic orbitals formed by mixing an atom's simple $s$, $p$, and sometimes $d$ orbitals to enable bonding geometries such as tetrahedra.

sp³ Hybridization

The mixing of one $s$ orbital and three $p$ orbitals to produce four equivalent hybrid orbitals arranged in a tetrahedral geometry with bond angles of $109.5^\circ$.

sp² Hybridization

The mixing of one $s$ orbital and two $p$ orbitals to produce three equivalent hybrid orbitals oriented in a trigonal planar arrangement with bond angles of $120^\circ$.

sp Hybridization

The mixing of one $s$ orbital and one $p$ orbital to produce two linear hybrid orbitals oriented at an angle of $180^\circ$.

Pi ($\pi$) Bond

A bond formed by the sideways overlap of unhybridized, parallel $p$ orbitals above and below the internuclear axis, occurring in double and triple bonds alongside a sigma bond.

SNAP Rule

A molecular polarity guideline stating that Symmetrical molecules are Nonpolar, whereas Asymmetrical molecules are Polar.

Dipole-Dipole Interaction

An intermolecular force of attraction between polar molecules where the positive dipole of one molecule aligns with the negative dipole of an adjacent molecule.

Hydrogen Bonding

An exceptionally strong type of dipole-dipole attraction occurring between molecules containing hydrogen bonded directly to highly electronegative nitrogen, oxygen, or fluorine ($N, O, F$).

Ion-Dipole Interaction

The electrostatic attraction between an ion and the oppositely charged pole of a polar solvent molecule, responsible for the dissolution of ionic compounds like $NaCl$ in water.

Induced Intermolecular Forces

Temporary attractions (such as dipole-induced dipole or ion-induced dipole) that occur when an ion or polar molecule alters the electron cloud density of an adjacent non-polar molecule, making it momentarily polar.

London Dispersion Forces

Weak intermolecular forces present in all molecules, particularly non-polar ones, caused by momentary uneven distributions of electrons creating instantaneous dipoles that induce charge in neighboring molecules.

Trigonal Bipyramidal Geometry ($AX_5$)

A molecular shape characterized by $5$ bonding electron pairs around a central atom with no lone pairs, exhibiting bond angles of $90^\circ$ (axial) and $120^\circ$ (equatorial).

Seesaw Molecular Shape ($AX_4E$)

A molecular shape derived from a trigonal bipyramidal electron group arrangement having $4$ bonding groups and $1$ lone pair, with bond angles of $<90^\circ$ (axial) and $<120^\circ$ (equatorial).

T-shaped Molecular Shape ($AX_3E_2$)

A molecular geometry derived from a trigonal bipyramidal electron group arrangement featuring $3$ bonding groups and $2$ lone pairs, exhibiting bond angles of $<90^\circ$.

Octahedral Geometry ($AX_6$)

A molecular geometry characterized by $6$ bonding electron pairs surrounding a central atom with bond angles of exactly $90^\circ$.

Square Planar Molecular Shape ($AX_4E_2$)

A molecular geometry derived from an octahedral electron group arrangement with $4$ bonding groups and $2$ lone pairs positioned opposite each other, displaying bond angles of $90^\circ$.