Atomic Structure, Chemical Bonding, and Nomenclature Principles

Foundational Atomic Theory and Discoveries

John Dalton (1808)

Dalton proposed that matter consists of indivisible particles called atoms. His model (sometimes historically referenced as the "fruitcake" model) established key principles:

• Atoms of the same element are identical in mass and properties.
• Atoms combine to form compounds and are not destroyed or created, only rearranged (involving the rearrangement of electrons).

Key Discoveries in Atomic Structure

• Michael Faraday (1833): Experiments proposed the relationship between electricity and atoms (electrolysis).
• J.J. Thomson: Studied electrical conductivity in gases using discharge tubes.
• Eugen Goldstein: Used a cathode ray tube to describe protons.
• Henri Becquerel (1896): Accidentally discovered radioactivity.
• Ernest Rutherford: Developed the nuclear model of the atom using the gold foil experiment.
• James Chadwick: Discovered the neutron.

The Bohr Model and Subatomic Particles

The current atomic theory describes electrons existing in specific orbits or energy levels. If an electron receives energy, it jumps to a higher level (circular orbits).

• Protons: Positively charged, located in the nucleus.
• Neutrons: Not charged (neutral), located in the nucleus.
• Electrons: Negatively charged, located in orbits/energy levels.

Quantum Mechanics and Electron Configuration

Pauli Exclusion Principle

Wolfgang Pauli proposed that no two electrons in the same atom can have the exact same set of four quantum numbers.

Hund's Rule

When filling degenerate orbitals (s, p, d, f sublevels), electrons must first occupy each orbital singly with parallel spins before pairing up.

Structure of the Periodic Table

The Periodic Table is ordered by increasing atomic number (Z) and is divided based on the orbital occupied by the last electron.

Orbital Sublevels and Maximum Capacity

Periods and Groups

• Periods (Rows 1 to 7): Elements in the same row have the same amount of energy levels.
• Groups (Columns): Elements in the same column generally have similar chemical properties.

Key Periodic Trends

• Atomic Radius: A measure of the distance between two equal nuclei. It depends on the ion: Cations have a smaller atomic radius than the parent atom, and Anions have a larger radius.
• Atomic Volume: V = m / d (Mass / Density).
• Ionization Energy: The energy required to extract an electron from a neutral atom.
• Electron Affinity: The energy exchanged when a neutral atom accepts an electron.
• Electronegativity: Measures the tendency of an atom to attract electrons that are linked in a bond. This value is higher for nonmetals than for metals.

Chemical Bonding Types

Ionic Bond

Formed by the transfer of electrons, typically occurring between a metal and a nonmetal. Characteristics include:

• Formation of crystal lattices.
• High boiling and melting points.
• The metal loses electrons (forming a cation) and the nonmetal gains electrons (forming an anion). Both elements achieve a stable electron configuration.

Covalent Bond

Formed by the sharing of electrons between two atoms, typically occurring between two nonmetals. Covalent bonds can be simple, double, triple, polar, or nonpolar.

Inorganic Chemical Nomenclature

Peroxides

Compounds containing two oxygen atoms with a valency of -1 (O₂²⁻). Example: Sodium Peroxide (Na₂O₂).

Hydrides

Compounds formed with Hydrogen.

• Metal Hydrides: Metal + H⁻ (e.g., Potassium Hydride, KH; Calcium Hydride, CaH₂).
• Nonmetal Hydrides: Nonmetal + Hydrogen (e.g., Hydrogen Fluoride, HF; Hydrogen Nitride, H₃N).
• Hydroacids: Nonmetal hydrides dissolved in water (e.g., HF(aq) is Hydrofluoric Acid).

Binary Salts

Formed by a metal and a nonmetal. The nonmetal is named first with the -ide suffix.

• Example: NaBr (Sodium Bromide); FeCl₃ (Iron(III) Chloride).

Hydroxides (Bases)

Formed by a metal and the hydroxide group (OH⁻).

• Example: Ca(OH)₂ (Calcium Hydroxide); KOH (Potassium Hydroxide); NH₄OH (Ammonium Hydroxide).

Oxoacids

Compounds containing Hydrogen, Oxygen, and a Nonmetal (H + O + Nonmetal). Naming depends on the oxidation state of the nonmetal, using suffixes:

• Lower oxidation state: -ous acid (e.g., HNO₂: Nitrous Acid).
• Higher oxidation state: -ic acid (e.g., HNO₃: Nitric Acid).

Oxosalts

Formed by a metal and an oxoanion (Metal + Oxoanion). The suffix of the acid determines the suffix of the salt:

• -ous acid becomes -ite salt (e.g., Na₂SO₃: Sodium Sulfite).
• -ic acid becomes -ate salt (e.g., CaSO₄: Calcium Sulfate).

Ammonium Oxosalts

Salts consisting of the Ammonium ion (NH₄⁺) and an oxoanion.

• Example: (NH₄)₂SO₄ (Ammonium Sulfate); NH₄NO₂ (Ammonium Nitrite).