Covalent Bonding, Lewis Structures, Molecular Geometry, Polarity, and Intermolecular Forces

Covalent Bonding and Lewis Structures

Covalent Bonding

Covalent bonding involves the sharing of electron pairs between atoms, resulting in the formation of molecules. This type of bond is based on the attractive force between positively charged nuclei and shared electrons.

• Covalent bond: Formed when two atoms share one or more pairs of electrons.

• Each shared pair constitutes a single covalent bond.

• Atoms achieve a stable electron configuration (octet) through sharing.

• Example: Two hydrogen atoms (H2) share electrons to form a single bond.

Lewis Structures

Lewis structures are diagrams that represent the arrangement of valence electrons among atoms in a molecule. They help visualize bonding and lone pairs.

• A single line represents a shared pair of electrons (a single bond).

• Dots represent lone pairs (non-bonding electrons).

• Lewis symbols are used to determine how many electrons each atom needs to achieve a stable octet.

• Atoms in group IVA (carbon family) have 4 valence electrons and typically form 4 bonds.

• Atoms in group VIIA (halogens) have 7 valence electrons and typically form 1 bond.

• Atoms can make up to 4 bonds to achieve a stable octet.

How Many Bonds to Fill the Octet?

The number of bonds an atom forms depends on the number of valence electrons it has and how many more it needs to reach 8 (the octet rule).

Element	Group	Valence Electrons	Bonds Needed
H	IA	1	1
C	IVA	4	4
N	VA	5	3
O	VIA	6	2
F	VIIA	7	1

Lewis Structure Examples

• Water (H2O): Oxygen (group VIA) has 6 valence electrons and forms 2 bonds with hydrogen atoms to complete its octet.

• Molecular Oxygen (O2): Each oxygen atom forms a double bond to satisfy the octet rule.

Rules for Lewis Structures

• Each atom must have a full octet (except for H, which needs 2 electrons).

• The total number of electrons in the structure must equal the sum of valence electrons for all atoms.

• If the molecule has a charge, adjust the total number of electrons accordingly.

• Atoms should have the expected number of bonds based on their group number.

• Some atoms (like S, P) can exceed the octet in certain cases (expanded octet).

Lewis Structure of CO32− (Carbonate Ion)

• Carbon (group IVA): 4 valence electrons

• Oxygen (group VIA): 6 valence electrons × 3 = 18

• Polyatomic ion: add 2 electrons for the 2− charge

• Total: 4 + 18 + 2 = 24 electrons

• Arrange electrons to satisfy octet for each atom, using double bonds as needed.

Lewis Structure of PO43− (Phosphate Ion)

• Total electrons: 5 (P) + 4 × 6 (O) + 3 (charge) = 32

• Distribute electrons to satisfy octet for each atom.

Lewis Structure of SO2 (Sulfur Dioxide)

• Total electrons: 6 (S) + 2 × 6 (O) = 18

• Arrange electrons to satisfy octet, using double bonds and lone pairs as needed.

Molecular Shape and VSEPR Theory

VSEPR: Valence Shell Electron Pair Repulsion

The shape of a molecule is determined by the repulsion between electron pairs (bonded and lone pairs) around a central atom. Electron pairs arrange themselves to minimize repulsion, resulting in specific molecular geometries.

• Double and triple bonds are treated as a single electron group.

Central Atom Geometry

Bond Groups	Bond Angle	Geometry	Example
2	180°	Linear	CO2
3	120°	Trigonal Planar	SO3
4	109.5°	Tetrahedral	CH4

Bond Polarity and Molecular Polarity

Definitions

• Polar bond: Uneven charge distribution between atoms.

• Nonpolar bond: Even charge distribution.

• Dipole: A molecule with a positive and negative end due to uneven electron distribution.

Bond Polarity (Unequal Sharing)

• Determined by the difference in electronegativity between atoms.

• Nonpolar covalent: difference < 0.5

• Polar covalent: difference 0.5–1.9

• Ionic: difference > 1.9

• Example: CH4 (0.4, nonpolar), NH3 (0.9, polar), NaCl (2.1, ionic)

Dipole

• Uneven charge distribution across a bond or molecule.

• Delta (δ) sign indicates partial positive or negative charge.

Molecular Polarity

• Depends on both bond polarity and molecular geometry.

• Symmetric molecules (e.g., CO2) are nonpolar because dipoles cancel out.

• Asymmetric molecules (e.g., H2O, NH3) are polar.

Intermolecular Forces and States of Matter

Intermolecular Forces (IMFs)

IMFs are forces that hold molecules together in the solid, liquid, and gaseous states. The strength of these forces determines the physical properties of substances.

• Solid state: Particles are closely packed and held in fixed positions.

• Liquid state: Particles are close but can move past each other.

• Gaseous state: Particles are far apart and move freely.

Definitions

• Endothermic: Absorbs heat (energy required for the process).

• Exothermic: Releases heat (energy released by the process).

• Enthalpy (ΔH): Amount of heat energy added to or released by the process.

Changing States of Matter

• Endothermic: Melting, vaporization, sublimation (require energy input).

• Exothermic: Freezing, condensation, deposition (release energy).

Types of Intermolecular Forces

• Dipole-dipole interactions: Attraction between polar molecules.

• London dispersion forces: Temporary dipoles in all molecules, stronger in larger atoms/molecules.

• Hydrogen bonding: Special strong dipole-dipole interaction involving H bonded to N, O, or F.

Summary Table: Intermolecular Forces

Type	Strength	Occurs Between
London Dispersion	Weakest	All molecules
Dipole-Dipole	Intermediate	Polar molecules
Hydrogen Bonding	Strongest (of IMFs)	H with N, O, or F

Summary

• Lewis structures help visualize bonding and lone pairs in molecules.

• VSEPR theory predicts molecular shapes based on electron pair repulsion.

• Bond and molecular polarity are determined by electronegativity differences and molecular geometry.

• Intermolecular forces influence physical properties and changes of state.