General Chemistry: Atomic Structure, Periodicity, and Chemical Bonding Study Notes

Atomic Structure and Periodicity

Subatomic Particles and Atomic Number

The structure of the atom is defined by the presence of subatomic particles: protons, neutrons, and electrons. The atomic number (Z) is the number of protons in the nucleus and determines the identity of the element.

• Proton: Positively charged particle found in the nucleus.

• Neutron: Neutral particle found in the nucleus.

• Electron: Negatively charged particle found in orbitals around the nucleus.

• Atomic Number (Z): Number of protons in the nucleus.

• Mass Number (A): Sum of protons and neutrons in the nucleus.

Example: Carbon has Z = 6, meaning 6 protons.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers.

• Isotope: Atoms with the same number of protons but different numbers of neutrons.

• Notation: $^{A}_{Z} ext{X}$, where X is the element symbol.

Example: $^{12}_{6} ext{C}$ and $^{13}_{6} ext{C}$ are isotopes of carbon.

Periodic Table Organization

The periodic table arranges elements by increasing atomic number and groups elements with similar chemical properties into columns.

• Groups: Vertical columns, elements share similar properties.

• Periods: Horizontal rows, elements have increasing atomic number.

• Metals, Nonmetals, Metalloids: Classification based on physical and chemical properties.

Electronic Structure of Atoms

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom's orbitals. The Aufbau principle, Pauli exclusion principle, and Hund's rule govern electron filling.

• Aufbau Principle: Electrons fill lowest energy orbitals first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.

• Hund's Rule: Electrons occupy degenerate orbitals singly before pairing.

Example: Electron configuration of oxygen: $1s^2 2s^2 2p^4$

Quantum Numbers

Quantum numbers describe the properties of atomic orbitals and the electrons in them.

• Principal quantum number (n): Energy level (shell).

• Angular momentum quantum number (l): Subshell (s, p, d, f).

• Magnetic quantum number (ml): Orientation of orbital.

• Spin quantum number (ms): Electron spin direction (+1/2 or -1/2).

Chemical Bonding

Ionic and Covalent Bonds

Chemical bonds form when atoms share or transfer electrons to achieve stable electron configurations.

• Ionic Bond: Formed by transfer of electrons from a metal to a nonmetal, resulting in oppositely charged ions.

• Covalent Bond: Formed by sharing of electrons between nonmetals.

Example: NaCl is an ionic compound; H2O is a covalent compound.

Lewis Structures and Electron-Dot Diagrams

Lewis structures represent the arrangement of valence electrons in molecules and help predict molecular shape and reactivity.

• Steps to Draw Lewis Structures:

  1. Count total valence electrons.

  2. Arrange atoms and connect with single bonds.

  3. Distribute remaining electrons to complete octets.

  4. Use double or triple bonds if necessary.

Example: Lewis structure of CO2: O=C=O

Bond Polarity and Electronegativity

Bond polarity arises from differences in electronegativity between bonded atoms.

• Electronegativity: Tendency of an atom to attract electrons in a bond.

• Polar Covalent Bond: Unequal sharing of electrons due to electronegativity difference.

• Nonpolar Covalent Bond: Equal sharing of electrons.

Example: H–F is polar; H–H is nonpolar.

Periodic Trends

Atomic Radius, Ionization Energy, and Electron Affinity

Periodic trends describe how atomic properties change across periods and groups.

• Atomic Radius: Decreases across a period, increases down a group.

• Ionization Energy: Increases across a period, decreases down a group.

• Electron Affinity: Generally increases across a period.

Trend	Across Period	Down Group
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electron Affinity	Increases	Varies

Additional info:

• Some content inferred from context and standard General Chemistry curriculum, as handwriting was partially unclear.

• Topics covered align with Ch.2 (Atoms, Molecules & Ions), Ch.5 (Periodicity & Electronic Structure of Atoms), Ch.6 (Ionic Compounds: Periodic Trends and Bonding Theory), and Ch.7 (Covalent Bonding and Electron-Dot Structures).