chem bonds

Why Atoms Form Bonds: The Octet Rule
Atoms form bonds to achieve a stable configuration with 8 electrons in their valence shell, lowering their energy and increasing stability.Ionic Bond
A bond formed when one atom transfers electrons to another, creating oppositely charged ions that attract. Electronegativity difference: ≥ 2.0. Example: NaCl.

Covalent Bond
A bond formed when two atoms share electrons. Electronegativity difference: < 2.0.

• Polar Covalent Bond: Unequal sharing of electrons, electronegativity difference 0.5–2.0. Example: H₂O.

• Nonpolar Covalent Bond: Equal sharing of electrons, electronegativity difference < 0.5. Example: H₂.

Metallic Bond
A bond where metal atoms share a "sea" of delocalized electrons that move freely throughout the structure. Example: Cu, Al.

Formation of Ionic Bonds
Occurs when a metal donates electrons (becoming a cation) and a nonmetal accepts them (becoming an anion). Example: NaCl.

Formula Unit
The simplest ratio of ions in an ionic compound that balances overall charge. Example: NaCl, MgCl₂.

Coulomb's Law
F=kq1q2r2F = k \frac{{q_1q_2}}{{r^2}}F=kr2q1​q2​​. The force of attraction between oppositely charged ions depends on charge magnitude and distance.

Crystal Lattice Structure
A 3D repeating pattern of ions in an ionic solid. Example: NaCl forms a cubic lattice.

Lattice Energy
The energy required to separate one mole of an ionic compound into gaseous ions. Increases with higher charges and smaller ion sizes.

Lewis Structures of Ionic Compounds
Represent electron transfer from a metal to a nonmetal. Example: Na with 1 dot transfers its electron to Cl with 7 dots.

Formation of a Covalent Bond
Atoms share electrons to fill their valence shells. Example: H + H → H₂.

Bond Length
The distance between two bonded atoms at minimum potential energy. Shorter bonds are stronger (triple < double < single).

Bond Energy
The energy required to break a bond. Stronger bonds have higher bond energy.

Lewis Structures of Molecules
Represent shared electrons as lines between atoms. Example: H₂O has single bonds with lone pairs on oxygen.

Single Bond
A covalent bond with one pair of shared electrons.

Double Bond
A covalent bond with two pairs of shared electrons.

Triple Bond
A covalent bond with three pairs of shared electrons.

Polyatomic Ion
A charged group of covalently bonded atoms. Example: SO42−\text{SO}_4^{2-}SO42−​.

VSEPR Theory
Valence Shell Electron Pair Repulsion theory states electron pairs repel to minimize repulsion, determining molecular shape.

Electron Pair Geometry
The 3D arrangement of regions of electron density (bonding and lone pairs). Example: CH₄ is tetrahedral.

Molecular Geometry
The 3D shape of a molecule determined by atoms only (not lone pairs). Example: Water is bent (104.5°).

Molecular Polarity
Uneven charge distribution in a molecule due to shape and bond polarity. Example: H₂O is polar.

Dipole
A polar bond with partial positive and negative charges.

Net Dipole
The overall molecular polarity resulting from the vector sum of individual bond dipoles.

Comparison of Ionic, Molecular, and Metallic Compounds

Exceptions to the Octet Rule

1. Incomplete Octet:
   Some elements are stable with fewer than 8 valence electrons:

   • Hydrogen (H): Stable with 2 electrons (e.g., H₂, HCl).

   • Beryllium (Be): Stable with 4 electrons (e.g., BeCl₂).

   • Boron (B): Stable with 6 electrons (e.g., BF₃).