Rucete ✏ SAT Chemistry In a Nutshell
3. Bonding
Chemical bonding explains how atoms combine to form stable structures. This chapter explores the types of chemical bonds—ionic, covalent, and metallic—as well as molecular geometry, hybridization, and intermolecular forces like hydrogen bonding and van der Waals interactions.
Why Atoms Form Bonds
- Atoms bond to achieve a stable octet (full outer shell).
- Bond formation releases energy; bond breaking absorbs energy.
- Noble gases don’t bond because they already have full outer shells.
Ionic Bonds
- Form when electronegativity difference ≥ 1.7.
- One atom donates electrons (cation), the other accepts (anion).
- Attraction between opposite charges forms the bond.
- High melting points, form crystal lattices, conduct electricity when dissolved.
Covalent Bonds
- Atoms share electrons to complete outer shells.
- Nonpolar: equal sharing (ΔEN = 0–0.4)
- Polar: unequal sharing (ΔEN = 0.4–1.6), forms dipoles
- Symmetric molecules can be nonpolar even if bonds are polar.
Metallic Bonds
- Valence electrons move freely in a "sea of electrons."
- Accounts for metal conductivity, malleability, and strength.
Intermolecular Forces
- Dipole-Dipole: between polar molecules (+/– attraction)
- London Dispersion: temporary dipoles, weak, in all molecules
- Hydrogen Bonding: H with N, O, or F; strongest IMF, raises boiling points
Double and Triple Bonds
- Double = 2 shared pairs; Triple = 3 shared pairs
- As bond order increases: bond strength ↑, bond length ↓
Resonance Structures
- Some molecules can't be drawn with one Lewis structure.
- Resonance: hybrid of multiple valid structures
- Examples: SO₃, C₆H₆ (benzene)
Molecular Geometry: VSEPR Theory
- Electron pairs repel to minimize repulsion → molecule shape
- Linear: 180° — BeF₂
- Trigonal planar: 120° — BF₃
- Tetrahedral: 109.5° — CH₄
- Trigonal pyramidal: ~107° — NH₃
- Bent: ~104.5° — H₂O
- Octahedral: 90° — SF₆
Hybridization
- Atomic orbitals blend to match observed shapes.
- sp: linear — BeF₂
- sp²: trigonal planar — BF₃
- sp³: tetrahedral — CH₄
- sp³d²: octahedral — SF₆
- Lone pairs distort bond angles — e.g., NH₃, H₂O
Sigma and Pi Bonds
- Sigma (σ): single bonds, head-on overlap
- Pi (π): side-on overlap; appears in double/triple bonds
- Double bond = 1 σ + 1 π; Triple = 1 σ + 2 π
Properties of Ionic Substances
- Conductive when molten/dissolved, not as solids
- High melting/boiling points
- Low vapor pressure, brittle, form electrolytic solutions
Properties of Molecular Crystals and Liquids
- Poor conductors
- Often gases or volatile liquids at room temp
- Low melting/boiling points, soft textures
- Require energy to break apart chemically
In a Nutshell
Bonding explains how atoms achieve stability through electron transfer or sharing. Ionic, covalent, and metallic bonds differ in how electrons are involved, influencing compound properties. Molecular geometry is predicted by VSEPR theory and explained by hybridization. Intermolecular forces like hydrogen bonding and London dispersion explain states of matter and boiling points. Recognizing bond types and structures is key to understanding chemical behavior.