General Chemistry: Chemical Bonding, Molecular Structure, and Lattice Energy Study Guide

Chemical Bonding and Molecular Structure

Delocalized vs. Localized Electrons

Electrons in molecules can be either delocalized (spread over several atoms) or localized (confined between two atoms). Delocalized electrons are often found in resonance structures and contribute to molecular stability.

• Delocalized electrons: Participate in bonding between more than two atoms, as in benzene or carboxylate ions.

• Localized electrons: Remain between two specific atoms, typical of single covalent bonds.

• Example: The electrons in the double bonds of benzene are delocalized over the six carbon atoms.

Bond Length and Bond Order

The bond order refers to the number of chemical bonds between a pair of atoms. Higher bond order generally means shorter bond length and greater bond strength.

• Bond order: Single (1), double (2), triple (3).

• Bond length: Decreases as bond order increases.

• Example: The N≡N bond in N2 is shorter and stronger than the N=N bond in N2H2.

Resonance Structures

Resonance occurs when more than one valid Lewis structure can be drawn for a molecule. The actual structure is a hybrid of all resonance forms.

• Example: Nitrate ion (NO3-) has three resonance structures.

• Resonance stabilizes molecules by delocalizing electrons.

Octet Rule and Exceptions

Most main-group elements tend to form bonds until they are surrounded by eight valence electrons (octet rule). Exceptions include molecules with odd numbers of electrons, incomplete octets, or expanded octets.

• Example: Boron trifluoride (BF3) has only six electrons around boron.

• Expanded octet: Elements in period 3 or higher (e.g., SF6).

Binary Ionic Compounds

Binary ionic compounds consist of a metal and a nonmetal. The metal loses electrons to become a cation, and the nonmetal gains electrons to become an anion.

• Formula: Metal cation + nonmetal anion (e.g., NaCl, CaO).

• Properties: High melting points, conduct electricity when molten or dissolved.

Chemical Bonds and Energy

The formation of chemical bonds releases energy, while breaking bonds requires energy. The potential energy of bonded atoms is less than that of the separate atoms.

• Lattice energy: The energy required to separate one mole of an ionic solid into gaseous ions.

• High lattice energy: Indicates strong ionic bonds and high melting points.

• Example: Lattice energy of NaCl is higher than that of KCl due to smaller ionic radii.

Molecular Geometry and Hybridization

VSEPR Theory

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shape of molecules based on electron pair repulsion around the central atom.

• Tetrahedral: 4 electron groups (e.g., CH4).

• Trigonal bipyramidal: 5 electron groups (e.g., PCl5).

• Octahedral: 6 electron groups (e.g., SF6).

Hybridization

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

• sp3: Tetrahedral geometry (e.g., NH3).

• sp2: Trigonal planar geometry (e.g., BF3).

• sp: Linear geometry (e.g., BeCl2).

• sp3d: Trigonal bipyramidal (e.g., PCl5).

• sp3d2: Octahedral (e.g., SF6).

Bond Angles

Bond angles depend on the number of electron groups and lone pairs around the central atom.

• Tetrahedral: 109.5°

• Trigonal planar: 120°

• Linear: 180°

• Example: NH3 has a bond angle of approximately 107° due to one lone pair.

Bonding Theories

Sigma and Pi Bonds

Sigma (σ) bonds are single covalent bonds formed by head-on overlap of orbitals. Pi (π) bonds are formed by side-to-side overlap and are present in double and triple bonds.

• Single bond: 1 σ bond

• Double bond: 1 σ + 1 π bond

• Triple bond: 1 σ + 2 π bonds

• Example: In C2H2 (acetylene), the C≡C bond has 1 σ and 2 π bonds.

Molecular Orbital (MO) Theory

MO theory describes the bonding in molecules by combining atomic orbitals to form molecular orbitals. Electrons fill the lowest energy orbitals first.

• Bond order: $\text{Bond order} = \frac{1}{2}[(\text{number of bonding electrons}) - (\text{number of antibonding electrons})]$

• Diamagnetic: All electrons are paired.

• Paramagnetic: At least one unpaired electron.

• Example: CO is diamagnetic; O2 is paramagnetic.

Polarity and Dipole Moments

Polarity of Molecules

A molecule is polar if it has a net dipole moment due to differences in electronegativity and molecular geometry.

• Non-polar: Symmetrical molecules (e.g., CO2).

• Polar: Asymmetrical molecules (e.g., H2O, SO2).

Lattice Energy and Ionic Compounds

Lattice Energy

Lattice energy is the energy released when gaseous ions form an ionic solid. It is a measure of the strength of the ionic bonds in a crystal lattice.

• Factors affecting lattice energy: Ionic charge (higher charge = higher lattice energy), ionic radius (smaller ions = higher lattice energy).

• Order of increasing lattice energy: KCl < NaCl < K2O < CaO (as charge increases and size decreases).

Lewis Structures and Resonance

Lewis Structures

Lewis structures represent the arrangement of electrons in a molecule. They show bonding pairs and lone pairs of electrons.

• Steps: Count valence electrons, draw skeleton, assign electrons to bonds and lone pairs, check octet rule.

• Formal charge: $\text{Formal charge} = \text{valence electrons} - \text{nonbonding electrons} - \frac{1}{2}(\text{bonding electrons})$

• Best resonance structure: Minimizes formal charges and places negative charges on more electronegative atoms.

Resonance Structures

Some molecules have multiple valid Lewis structures, called resonance structures. The actual molecule is a hybrid of these forms.

• Example: O3, NO3-, CO32-

• Number of resonance structures: Sulfite (SO32-) has 3 resonance structures.

Nomenclature and Acid Naming

Naming Ionic and Covalent Compounds

Compounds are named based on their composition and bonding.

• Ionic compounds: Name cation first, then anion (e.g., sodium sulfide, Na2S).

• Covalent compounds: Use prefixes to indicate number of atoms (e.g., dinitrogen monoxide, N2O).

• Acids: Hydrochloric acid (HCl), phosphoric acid (H3PO4), periodic acid (HIO4).

Summary Table: Molecular Geometry, Electron Geometry, and Hybridization

Electron Groups	Electron Geometry	Molecular Geometry	Hybridization	Bond Angle
2	Linear	Linear	sp	180°
3	Trigonal planar	Trigonal planar, Bent	sp2	120°
4	Tetrahedral	Tetrahedral, Trigonal pyramidal, Bent	sp3	109.5°
5	Trigonal bipyramidal	Trigonal bipyramidal, Seesaw, T-shaped, Linear	sp3d	90°, 120°
6	Octahedral	Octahedral, Square pyramidal, Square planar	sp3d2	90°

Practice Applications

• Lewis Structures: Practice drawing for CF3Cl, PF3, CO, and resonance for O3, NO3-, CO32-.

• Bond Order Calculation: For CN-, bond order = 3.

• MO Theory: CO is diamagnetic (all electrons paired).

• Hybridization: Ammonia (NH3) is sp3 hybridized; ICF2- is sp3d.