Ionic bonds form when electrons are transferred rather than shared, resulting in ions with opposite charges. This electrostatic attraction stabilizes the compound.

Covalent bonds involve the sharing of electron pairs between atoms, allowing them to achieve a stable electron configuration. This method of bonding is common in many molecular compounds.

Ionic bonds typically form between metals and non-metals due to the transfer of electrons. The metal loses electrons while the non-metal gains them, resulting in charged ions.

The octet rule states that atoms tend to acquire a full valence shell of 8 electrons to achieve stability. Sharing electrons through covalent bonding is one way to satisfy this rule.

H₂ is a diatomic molecule where two hydrogen atoms share electrons, forming a covalent bond. Ionic compounds like NaCl involve electron transfer rather than sharing.

Polar covalent bonds occur when electrons are shared unequally, often due to a difference in electronegativity between atoms. This results in partial charges on the atoms, creating a dipole moment.

NaCl is formed when sodium transfers an electron to chlorine, resulting in oppositely charged ions that attract each other. This is a classic example of ionic bonding.

The oxygen atom in H₂O has two bonding pairs and two lone pairs, which results in a tetrahedral electron pair geometry. However, the molecular shape is bent due to the repulsion from the lone pairs.

CO₂ is a linear molecule with polar bonds; however, the dipoles cancel out, resulting in a nonpolar molecule. In contrast, molecules like H₂O have asymmetrical shapes that lead to a net dipole moment.

The octet rule states that atoms are most stable when they have 8 electrons in their valence shell. This is why atoms share, gain, or lose electrons to achieve this stable configuration.

A larger difference in electronegativity between two atoms leads to an unequal sharing of electrons, thereby increasing the ionic character of the bond. This results in a bond with partial charges on the atoms.

Metallic bonding is characterized by a 'sea' of delocalized electrons that move freely among metal cations. This delocalization contributes to the electrical conductivity and malleability of metals.

Lewis structures are used to visually depict the valence electrons and bonding between atoms in molecules. They are a fundamental tool in understanding molecular structure and reactivity.

Resonance occurs when multiple valid Lewis structures can represent the electron distribution in a molecule. This delocalization of electrons often contributes to the overall stability of the molecule.

Double bonds involve the sharing of two pairs of electrons, which shortens the distance between the nuclei and increases bond strength. This makes double bonds generally shorter and stronger compared to single bonds.

Molecular orbital theory explains the bonding in O₂ by combining atomic orbitals into molecular orbitals that extend over both atoms. The resulting bond order of 2 and the presence of unpaired electrons in antibonding orbitals account for O₂'s paramagnetic behavior.

Bond order, which reflects the number of shared electron pairs between atoms, is a key determinant of covalent bond strength. A higher bond order generally results in a shorter and stronger bond.

Elements in the third period and beyond have d-orbitals available, which enable them to accommodate more than 8 electrons in their valence shell. This capacity for an expanded octet explains many exceptions to the traditional octet rule.

Hybridization of atomic orbitals directly influences molecular geometry and bond angles. For example, CH₄ uses sp³ hybridization to form a tetrahedral shape (about 109.5°), while C₂H₄ uses sp² hybridization to form a trigonal planar shape (about 120°).

Resonance stabilization involves the delocalization of electrons across a system of connected atoms, which distributes electron density and enhances stability. This effect is particularly important in conjugated systems where multiple valid structures can be drawn.