Ionic bonding: Electron transfer between a metal and a nonmetal creates oppositely charged ions held together by electrostatic attraction.

• Forms extended ionic lattices
• High melting points and electrical conductivity when molten or dissolved

Covalent bonding: Nonmetals share electron pairs to lower potential energy and achieve more stable electron arrangements.

• Can be nonpolar or polar depending on electronegativity difference
• Bond length and bond strength depend on the balance of attractive and repulsive forces

Metallic bonding: Metal cations are attracted to a sea of delocalized valence electrons.

• Explains conductivity, malleability, ductility, and luster

Bond polarity: Determined by electronegativity difference between bonded atoms.

• Nonpolar covalent: electrons shared equally or nearly equally
• Polar covalent: electrons shared unequally, creating partial charges
• Greater electronegativity difference -> larger dipole moment

Molecular polarity: Vector sum of all bond dipoles.

• Symmetric molecules with identical surrounding atoms can be nonpolar even if bonds are polar
• Asymmetric geometry can produce a net dipole and a polar molecule
• Polarity affects intermolecular forces, solubility, and boiling point

Polar covalent bond: Unequal electron sharing due to electronegativity difference

• Dipole moment (mu) points toward more electronegative atom
• Represented by arrow with cross at positive end

Molecular polarity: Vector sum of all bond dipoles

• Symmetrical molecules can be nonpolar despite polar bonds (e.g., $CO_2$)
• Asymmetrical molecules with polar bonds are polar (e.g., $H_2O$)

Energy required to break a chemical bond

• Average bond enthalpies: H-H 436 kJ/mol, C-C 347 kJ/mol, C=C 614 kJ/mol
• Stronger bonds = higher bond energy
• Multiple bonds stronger than single bonds

Bond length: Distance between nuclei of bonded atoms

• Shorter bonds are generally stronger (higher bond energy)
• Triple bonds < double bonds < single bonds (in length)
• Triple bonds > double bonds > single bonds (in strength)

Factors affecting bond length:

• Atomic size (larger atoms -> longer bonds)
• Bond order (higher order -> shorter, stronger bonds)
• Resonance (delocalization shortens and strengthens bonds)

Graph of potential energy vs. internuclear distance

• Equilibrium bond length: distance at minimum potential energy
• Bond energy: depth of potential well
• Steeper repulsion region for shorter distances (nucleus-nucleus repulsion)

For bond formation:

• Reactants (separated atoms) at higher potential energy
• Products (bonded molecule) at lower potential energy
• Energy released = bond energy

For bond breaking:

• Bonded molecule at lower potential energy
• Separated atoms at higher potential energy
• Energy absorbed = bond energy

Number of ions of opposite charge surrounding each ion

• NaCl structure: coordination number = 6 (octahedral)
• CsCl structure: coordination number = 8 (cubic)
• ZnS structure: coordination number = 4 (tetrahedral)

Smallest repeating unit in ionic crystal

• Represents ratio of ions in compound
• NOT discrete molecules

Lattice energy: Energy released when gaseous ions form 1 mole of solid ionic compound

• Proportional to $(Q_1 \times Q_2) / r$ (charge product / distance)
• Higher lattice energy -> higher melting point, greater hardness

Factors affecting lattice energy:

• Ion charges: higher charges -> much higher lattice energy
• Ion sizes: smaller ions -> higher lattice energy
• $MgO$ has much higher lattice energy than $NaCl$ ($Mg^{2+}, O^{2-}$ vs. $Na^+, Cl^-$)

Metallic bonding model where valence electrons are delocalized

• Positive metal ions in "sea" of delocalized electrons
• Explains metallic properties: conductivity, malleability, ductility
• Electrons free to move throughout structure

Substitutional alloys: Solute atoms replace solvent atoms in crystal structure

• Similar atomic sizes required
• Brass (Cu-Zn), Bronze (Cu-Sn), Sterling silver (Ag-Cu)

Interstitial alloys: Solute atoms occupy holes in solvent crystal structure

• Small solute atoms (C, B, N) in larger metal lattice
• Steel (Fe with C), Cast iron

1. Count total valence electrons
2. Arrange atoms (least electronegative usually central)
3. Connect atoms with single bonds (2 electrons each)
4. Complete octets of outer atoms
5. Place remaining electrons on central atom
6. If central atom lacks octet, form multiple bonds

Octet rule: Atoms tend to have 8 valence electrons (full valence shell) Exceptions:

• H, He, Li, Be, B (less than 8)
• Elements in period 3+ (can have expanded octets using d-orbitals)
• Odd-electron molecules (e.g., $NO_2$)

Multiple valid Lewis structures for same molecule

• Differ only in electron position (not atom positions)
• True structure is hybrid (delocalized electrons)
• Bond order between single and double

Example: Ozone ($O_3$) $O=O^+-O^- \leftrightarrow O^- - O^+=O$

$FC = V - N - \frac{B}{2}$

• V = valence electrons (free atom)
• N = nonbonding (lone pair) electrons
• B = bonding electrons

Guidelines:

• Most stable structure has formal charges closest to zero
• Negative formal charge on most electronegative atom

Electron geometry: Arrangement of all electron domains (bonding + lone pairs) Molecular geometry: Arrangement of atoms only (ignores lone pairs)

VSEPR theory: Electron domains repel; arrange to minimize repulsion

• Lone pair-lone pair > lone pair-bond pair > bond pair-bond pair repulsion

Ideal angles affected by:

• Lone pairs (compress bond angles)
• Multiple bonds (slightly larger repulsion)
• Different atoms (electronegativity differences)

Common angles:

• Linear: 180 degrees
• Trigonal planar: 120 degrees
• Tetrahedral: 109.5 degrees
• Bent: < 109.5 degrees (depends on lone pairs)
• Trigonal pyramidal: < 109.5 degrees

Hybridization: Mixing atomic orbitals to form new equivalent orbitals

• sp: linear, 180 degrees (1 s + 1 p orbital)
• sp2: trigonal planar, 120 degrees (1 s + 2 p orbitals)
• sp3: tetrahedral, 109.5 degrees (1 s + 3 p orbitals)
• sp3d: trigonal bipyramidal (1 s + 3 p + 1 d orbital)
• sp3d2: octahedral (1 s + 3 p + 2 d orbitals)