Molecular Structure, Bonding, and Properties: A Study Guide

Molecular Representations and Lewis Structures

Translating Between Molecular Representations

Chemists use several models to represent molecules, each highlighting different structural features. Understanding how to translate between these representations is essential for visualizing molecular structure and predicting chemical behavior.

• Space-filling models: Show the relative sizes and spatial arrangement of atoms, emphasizing the molecule's overall shape and volume.

• Ball-and-stick models: Depict atoms as balls and bonds as sticks, clarifying bond angles and connectivity.

• Perspective formulae: Use wedges and dashes to indicate three-dimensional arrangement on paper.

• Lewis structures: Represent valence electrons as dots and bonds as lines, focusing on electron arrangement and connectivity.

Example: Methane (CH4) can be shown as a Lewis structure (central C with four single bonds to H), a ball-and-stick model (tetrahedral geometry), or a space-filling model (compact, nearly spherical shape).

Drawing Lewis Structures

Lewis structures are diagrams that show the bonding between atoms and the lone pairs of electrons in a molecule. They are especially useful for molecules containing C, H, N, O, F, Cl, Br, I, P, and S.

• Step 1: Count total valence electrons for all atoms.

• Step 2: Arrange atoms (least electronegative in the center, H and F are always terminal).

• Step 3: Connect atoms with single bonds, then distribute remaining electrons to complete octets (or duets for H).

• Step 4: Use double or triple bonds if necessary to satisfy the octet rule.

Example: For CO2 (carbon dioxide):

• Total valence electrons: 4 (C) + 2×6 (O) = 16

• Lewis structure: O=C=O (each O double-bonded to C, all atoms have full octets)

Structural Isomers and Alternative Representations

Structural isomers are compounds with the same molecular formula but different connectivity of atoms. Alternative representations (e.g., different drawings of the same molecule) do not change connectivity.

• Example: C4H10 can be n-butane (straight chain) or isobutane (branched chain).

• Distinguishing: Isomers have different physical and chemical properties; alternative representations do not.

Bonding and Molecular Structure

Bond Rotation: Sigma and Pi Bonds

Bond rotation is influenced by the type of bond between atoms.

• Sigma (σ) bonds: Formed by head-on overlap of orbitals; allow free rotation around the bond axis.

• Pi (π) bonds: Formed by side-on overlap; restrict rotation due to electron cloud overlap above and below the bond axis.

Example: Ethane (C2H6) has free rotation around the C–C single bond; ethene (C2H4) cannot rotate around the C=C double bond.

VSEPR Theory: Electron Pair Geometry and Molecular Shape

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on repulsion between electron pairs around a central atom.

• Electron pair geometry: Arrangement of all electron pairs (bonding and lone pairs).

• Molecular shape: Arrangement of only the atoms (ignoring lone pairs).

Example: Ammonia (NH3): Electron pair geometry is tetrahedral; molecular shape is trigonal pyramidal.

Hybridization

Hybridization describes the mixing of atomic orbitals to form new, equivalent hybrid orbitals for bonding.

• sp3 hybridization: Four electron domains (e.g., CH4).

• sp2 hybridization: Three electron domains (e.g., C2H4).

• sp hybridization: Two electron domains (e.g., C2H2).

Converting Between Lewis Structures and 3D Representations

Lewis structures are two-dimensional, but molecules exist in three dimensions. Use VSEPR and hybridization to predict 3D geometry from a Lewis structure, and vice versa.

• Example: Water (H2O) has a bent shape in 3D, even though its Lewis structure appears linear.

Bond Polarity, Molecular Polarity, and Intermolecular Forces

Bond Polarity and Electronegativity

Electronegativity is the ability of an atom to attract shared electrons. The difference in electronegativity between two atoms determines bond polarity.

• Nonpolar covalent bond: Electronegativity difference < 0.5

• Polar covalent bond: Electronegativity difference 0.5–1.7

• Ionic bond: Electronegativity difference > 1.7

Example: H–Cl bond is polar; electrons are drawn toward Cl.

Molecular Polarity

Molecular polarity depends on both bond polarity and molecular shape. A molecule is polar if it has a net dipole moment.

• Example: CO2 has polar bonds but is nonpolar overall due to its linear shape; H2O is polar due to its bent shape.

Intermolecular Forces (IMFs)

IMFs are forces between molecules that affect physical properties.

• London dispersion forces: Present in all molecules; stronger in larger, more polarizable molecules.

• Dipole-dipole interactions: Occur between polar molecules.

• Hydrogen bonding: Strong dipole-dipole interaction when H is bonded to N, O, or F.

Example: Water exhibits hydrogen bonding, leading to high boiling point.

IMFs and Physical Properties

IMFs explain trends in melting and boiling points, solubility, and other properties.

• Stronger IMFs → higher melting/boiling points.

• Hydrogen bonding > dipole-dipole > London dispersion (in general strength).

Example: Methanol (CH3OH) boils at a higher temperature than methane (CH4) due to hydrogen bonding.

Covalent and Ionic Bonding

Covalent vs. Ionic Bonding: The Bonding Continuum

Bonds exist on a continuum from pure covalent (equal sharing of electrons) to ionic (complete transfer of electrons).

• Covalent bond: Electrons shared between atoms (e.g., H2).

• Ionic bond: Electrons transferred from metal to nonmetal (e.g., NaCl).

• Most bonds have some degree of both covalent and ionic character.

Formation of Ions: Metals vs. Nonmetals

Metals tend to lose electrons and form positive ions (cations); nonmetals tend to gain electrons and form negative ions (anions).

• Common cations: Na+, Mg2+, Al3+

• Common anions: Cl−, O2−, N3−

Predicting charge: Main group metals form cations with charge equal to their group number; nonmetals form anions with charge equal to 8 minus their group number.

Structure and Properties of Ionic Compounds

Structure of Ionic Compounds

Ionic compounds form extended lattice structures, maximizing attractive forces between oppositely charged ions.

• Example: NaCl forms a cubic lattice with each Na+ surrounded by six Cl− ions and vice versa.

Properties of Ionic Compounds

• High melting and boiling points: Due to strong electrostatic forces.

• Hard and brittle: Lattice structure resists deformation but shatters if layers shift.

• Electrical conductivity: Conducts electricity when molten or dissolved in water, not as a solid.

Comparing Melting and Boiling Points

The strength of forces (ionic, covalent, or intermolecular) and the amount of energy required to overcome them explain differences in melting and boiling points.

• Ionic compounds > covalent compounds (with only London forces) in melting/boiling points.

• Within covalent compounds, those with hydrogen bonding have higher melting/boiling points than those with only dispersion forces.

Summary Table: Types of Bonding and Properties

Bond Type	Example	Structure	Melting/Boiling Point	Electrical Conductivity
Ionic	NaCl	Lattice	High	Only when molten or in solution
Covalent (polar)	H2O	Molecules	Low to moderate	Non-conductive
Covalent (nonpolar)	CH4	Molecules	Low	Non-conductive
Metallic	Fe	Metal lattice	Variable	Conductive (solid and liquid)