Rucete ✏ AP Chemistry In a Nutshell
4. Covalent Compounds, Formulas, and Structures
This chapter covers the formation and representation of covalent molecules, including Lewis structures, bond types, molecular shapes, bond polarity, resonance, and hybridization.
Covalent Bonding and Lewis Structures
• Covalent bonds form when atoms share electrons to attain noble gas configurations.
• Lewis structures represent valence electrons as dots around element symbols.
• The octet rule drives most covalent bonding: atoms aim for eight electrons in their valence shell.
Drawing Lewis Structures
• Determine the skeleton: central atom is usually carbon or least electronegative element (never H).
• Count total valence electrons (adjust for ion charge).
• Connect atoms using bonding pairs (single bonds).
• Distribute remaining electrons to complete outer atom octets.
• Place leftover electrons on the central atom; if central atom lacks octet, convert lone pairs into double or triple bonds.
• Atoms in period 3 or higher can exceed the octet rule.
Electron-Deficient and Expanded Octets
• Boron often forms electron-deficient compounds (e.g., BF₃).
• Phosphorus and sulfur can expand their octets (e.g., PCl₅, SF₆).
Multiple Bonds and Resonance
• Double and triple bonds form when more electron pairs are shared between atoms.
• Resonance structures exist when multiple valid Lewis structures differ only in electron placement.
• The true structure is a blend of all resonance forms (delocalized electrons).
Lewis Structures of Ions
• Add electrons for anions; subtract for cations.
• Enclose the structure in brackets and indicate the net charge.
Odd Electron Molecules (Free Radicals)
• Some molecules have an odd number of electrons (e.g., NO₂) and cannot form full octets for all atoms.
• These are highly reactive and often pair up (dimerize).
Formal Charge
• Used to evaluate the reasonableness of a Lewis structure.
• Formal charge = (valence electrons) – (nonbonding electrons + ½ bonding electrons).
• Best structures have minimal formal charges and assign negative charges to more electronegative atoms.
Bond Polarity and Electronegativity
• Electronegativity (EN) measures an atom’s ability to attract electrons.
• EN increases from lower left to upper right of the periodic table.
• A bond is polar if ΔEN > 0; if ΔEN ≥ 1.7, it is considered ionic.
• δ+ and δ− are used to show partial charges in polar bonds.
Dipole Moment
• Measures bond polarity as the product of charge and distance: μ = q × r.
• Expressed in debye units (D).
• Vectors representing dipoles can cancel in symmetrical molecules.
Molecular Polarity
• Determined by both bond polarity and molecular shape.
• Nonpolar molecules can contain polar bonds if the dipoles cancel (e.g., CO₂).
• Polar molecules have net dipole moments (e.g., H₂O, NH₃).
Valence Shell Electron Pair Repulsion (VSEPR) Theory
• Electron groups around a central atom repel each other and determine molecular geometry.
• Lone pairs take up more space than bonding pairs and can distort bond angles.
• Common geometries:
• Linear (2 groups): 180°
• Trigonal planar (3 groups): 120°
• Tetrahedral (4 groups): 109.5°
• Trigonal bipyramidal (5 groups): 90°, 120°
• Octahedral (6 groups): 90°
Effects of Lone Pairs
• Lone pairs reduce bond angles due to increased repulsion.
• Example: H₂O is bent, not linear, due to two lone pairs on oxygen.
Hybridization
• Atomic orbitals mix to form hybrid orbitals used in bonding.
• Number of electron groups determines hybridization:
• 2 groups: sp
• 3 groups: sp²
• 4 groups: sp³
• 5 groups: sp³d
• 6 groups: sp³d²
• Hybrid orbitals explain observed shapes of molecules (e.g., CH₄ = tetrahedral = sp³).
Sigma (σ) and Pi (π) Bonds
• Sigma bonds: head-on overlap of orbitals, present in all single bonds.
• Pi bonds: side-by-side overlap, present in double (1 π) and triple bonds (2 π).
• Multiple bonds are stronger and shorter than single bonds.
Molecular Orbitals (MO Theory)
• Atomic orbitals combine to form bonding and antibonding molecular orbitals.
• Electrons fill lower-energy bonding orbitals first.
• Bond order = ½(number of bonding electrons – number of antibonding electrons).
• Higher bond order → stronger and shorter bond.
In a Nutshell
Covalent compounds form through electron sharing, and their structures are described using Lewis diagrams, VSEPR theory, and hybridization. Molecular polarity depends on bond types and shapes, while sigma and pi bonds explain bond strength and length. Understanding these models allows us to predict structure, shape, and reactivity of molecules.