Atomic Structure, Periodic Properties, and Chemical Bonding

Comprehensive vocabulary flashcards covering atomic models, subatomic particles, quantum numbers, periodic trends, chemical bonding, and Lewis structures based on the lecture notes.

Atom (Democritus)

A fundamental entity proposed in the 5th century BC, originating from the Greek word atomos meaning indivisible, conceptualized as tiny, indivisible particles that compose all matter.

Atomic Concept (Ruđer Bošković)

A model defining the atom as a dimensionless material point that serves as a center of forces, interacting with other particles via attractive and repulsive forces.

Dalton's Atomic Theory

A theory defining atoms as tiny, indivisible spheres where atoms of the same element share equal mass and properties, different elements have differing masses and properties, compounds form by combining atoms in definite ratios, and chemical reactions involve rearranging atoms without creating or destroying them.

Plum Pudding Model

An atomic model proposed by J.J. Thomson following his discovery of the electron, depicting the atom as a positively charged sphere embedded with negatively charged electrons.

Nuclear Model of the Atom

A model proposed by Ernest Rutherford following his discovery of the atomic nucleus, stating that an atom consists of a small, positively charged central nucleus surrounded by electrons, with the majority of the atom consisting of empty space.

Neutron

An electrically neutral subatomic particle discovered by James Chadwick located within the atomic nucleus, which together with protons constitutes almost all of an atom's mass.

Electrical Neutrality of an Atom

The condition where an atom has a net charge of zero because it contains an equal number of positively charged protons and negatively charged electrons.

Nucleus-to-Atom Diameter Ratio

The structural size relationship wherein the diameter of the atomic nucleus is approximately $100\,000$ times ($10^5$ times) smaller than the diameter of the atom.

Atomic Number ($Z$)

The fundamental property of an atom defined as the total number of protons in its nucleus.

Nucleon Number ($A$)

The total number of nucleons in an atomic nucleus, equal to the sum of the number of protons and the number of neutrons.

Chlorine-35 ($^{35}_{17}\text{Cl}$)

A neutral isotope of chlorine possessing a nucleon number $A = 35$ and an atomic number $Z = 17$, containing $17$ protons, $18$ neutrons ($35 - 17 = 18$), and $17$ electrons.

Atomic Mass Distribution

The physical concentration of almost the entire mass of an atom within its nucleus, owing to the mass of protons and neutrons while the mass of electrons remains negligible.

Ion

An electrically charged particle produced when an atom or group of atoms loses electrons (forming a positive cation) or gains electrons (forming a negative anion).

Isoelectronic Particles

Distinct atoms or ions belonging to different chemical elements that share the exact same number of electrons.

Nuclide

A specific atomic species uniquely identified and defined by its exact number of protons and neutrons.

Isotopes

Atoms of the same chemical element that possess identical numbers of protons but differing numbers of neutrons.

Isobars

Atoms of different chemical elements that share the same nucleon number ($A$).

Protium ($^1\text{H}$)

The isotope of hydrogen comprising $1$ proton, $0$ neutrons, and $1$ electron.

Deuterium ($^2\text{H}$)

The isotope of hydrogen comprising $1$ proton, $1$ neutron, and $1$ electron.

Tritium ($^3\text{H}$)

The isotope of hydrogen comprising $1$ proton, $2$ neutrons, and $1$ electron.

Bohr's Atomic Model

An atomic model proposed by Niels Bohr that accounts for the line spectrum of the hydrogen atom by postulating that electrons can possess only discrete, quantized energy levels.

Stationary State

A specific quantum state in an atom where an electron maintains a defined, constant energy and does not emit radiation.

Energy Absorption and Emission

Processes where an electron absorbs energy to transition to a higher energy level, or transitions to a lower energy level while emitting energy in the form of a photon.

Quantum Energy Equation

The mathematical formula determining radiation energy: $E = h \cdot \nu$ or $E = \frac{h \cdot c}{\lambda}$, where $E$ is energy, $h$ is Planck's constant, $\nu$ is frequency, $c$ is the speed of light, and $\lambda$ is wavelength.

Principal Quantum Number ($n$)

The quantum number designating the main energy level of an electron in an atom.

Orbital Quantum Number ($l$)

The quantum number that specifies the geometrical shape of an atomic orbital.

Magnetic Quantum Number ($m$)

The quantum number that describes the spatial orientation of an atomic orbital.

Spin Quantum Number ($m_s$)

The quantum number defining the intrinsic spin of an electron, taking values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

Electron Configuration

The systematic arrangement of an atom's electrons across its energy levels, sublevels, and atomic orbitals.

Diagonal Rule

The principle defining the order of filling atomic orbitals from lower to higher energy: $1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d \rightarrow 4p \rightarrow 5s \rightarrow 4d \rightarrow 5p \rightarrow 6s$.

Groups (Periodic Table)

Vertical columns in the periodic table containing elements with similar chemical properties due to having similar configurations of valence electrons.

Periods (Periodic Table)

Horizontal rows in the periodic table whose row number corresponds to the total number of occupied electron shells for main group elements.

Main Groups

The eight primary columns of the periodic table: alkali metals, alkaline earth metals, boron group, carbon group, nitrogen group, oxygen group (chalcogens), halogens, and noble gases.

Periodic Properties

Elemental characteristics that vary predictably across the periodic table, including atomic radius, ionization energy, electron affinity, electronegativity, and metallic character.

Atomic Radius

A dimension indicating the size of an atom, which increases descending down a group and decreases moving from left to right across a period.

Ionization Energy

The amount of energy required to remove an electron from an isolated atom in the gaseous state, which increases for each successive electron due to stronger attraction from the positive nucleus.

Electron Affinity

The energy change accompanying the addition of an electron to an isolated atom in the gaseous state, generally increasing toward the upper right of the periodic table, subject to exceptions.

Electronegativity

The tendency of an atom to attract a shared electron pair within a chemical bond, which increases upward and to the right in the periodic table with values ranging from approximately $0.7$ (cesium, $\text{Cs}$) to $4.0$ (fluorine, $\text{F}$).

Ionic Bond

A chemical linkage formed typically between metals and nonmetals via the complete transfer of electrons, generating oppositely charged ions held together by electrostatic attractive forces.

Covalent Bond

A chemical linkage formed predominantly between nonmetal atoms through the mutual sharing of one or more electron pairs to attain a stable electron configuration.

Metallic Bond

A type of chemical bond occurring between metal atoms characterized by delocalized valence electrons shared across the lattice.

Formula Unit

The lowest whole-number ratio of cations to anions in an ionic compound that yields an electrically neutral species, such as a $1:1$ ratio in $\text{NaCl}$.

Redox Reaction

A chemical process characterized by electron transfer, wherein oxidation and reduction occur simultaneously.

Oxidation and Reduction

Coupled processes where oxidation involves the loss of electrons and reduction involves the gain of electrons.

Oxidizing Agent and Reducing Agent

Chemical roles where the oxidizing agent accepts electrons and undergoes reduction, whereas the reducing agent donates electrons and undergoes oxidation.

Properties of Ionic Crystals

Physical traits including solid and brittle structures, high melting and boiling points, electrical insulation in solid form, and electrical conductivity when molten or dissolved in water.

Ionic vs. Molecular Crystal Melting Points

The thermal distinction where ionic crystals exhibit high melting points due to strong electrostatic forces between oppositely charged ions, whereas molecular crystals have lower melting points due to weaker intermolecular forces.

Electrical Conductivity in Melts

The phenomenon where molten ionic compounds conduct electricity because their ions are free to move and transport charge, while molten molecular compounds generally do not conduct due to the absence of mobile ions or electrons.

Diatomic Covalent Multiplicities

Homonuclear molecules exemplifying bond types: single covalent bond in $\text{H}_2$ ($\text{H}-\text{H}$), double covalent bond in $\text{O}_2$ ($\text{O}=\text{O}$), and triple covalent bond in $\text{N}_2$ ($\text{N}\equiv\text{N}$).

Valence Electrons

Electrons located in the outermost electron shell of an atom that participate in chemical bonding through transfer (ionic bonding) or sharing (covalent bonding).

Bonding in Aluminum Oxide ($\text{Al}_2\text{O}_3$)

An ionic compound where two $\text{Al}$ atoms each donate $3$ electrons to become $2\text{Al}^{3+}$ ($6$ electrons donated in total) and three $\text{O}$ atoms each accept $2$ electrons to become $3\text{O}^{2-}$, yielding the ionic formula $2\text{Al}^{3+} + 3\text{O}^{2-}$.

Lewis Structure of Ammonia ($\text{NH}_3$)

A covalent molecule in which a central nitrogen atom with $5$ valence electrons shares pairs with three hydrogen atoms (each contributing $1$ electron) to form three single covalent bonds and retain one non-bonding lone pair.

Lewis Structure of Methane ($\text{CH}_4$)

A covalent molecule in which a central carbon atom with $4$ valence electrons forms four single covalent bonds by sharing electrons with four separate hydrogen atoms.