Bonding, Molecular Structure, and VSEPR Theory: Study Notes for General Chemistry

Chemical Bonding and Molecular Structure

Types of Chemical Bonds

Chemical bonds are the attractive forces that hold atoms together in compounds. The two main types are ionic and covalent bonds, each with distinct properties and formation mechanisms.

• Ionic Bonds: Formed by the transfer of electrons from metals to nonmetals, resulting in the formation of oppositely charged ions that attract each other.

• Covalent Bonds: Formed by the sharing of electrons between nonmetal atoms. Covalent bonds can be polar or nonpolar depending on the electronegativity difference between the atoms.

• Metallic Bonds: Involve a 'sea' of delocalized electrons shared among a lattice of metal atoms.

Example: Sodium chloride (NaCl) is an ionic compound, while water (H2O) is a covalent compound.

Electronegativity and Bond Polarity

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

• Nonpolar Covalent Bond: Electrons are shared equally (e.g., H2).

• Polar Covalent Bond: Electrons are shared unequally, creating partial charges (e.g., HCl).

• Ionic Bond: Electrons are transferred, resulting in full charges (e.g., NaCl).

Equation: Electronegativity difference ($\Delta EN$) can be used to estimate bond type:

$\Delta EN < 0.4$ → Nonpolar covalent $0.4 \leq \Delta EN < 2.0$ → Polar covalent $\Delta EN \geq 2.0$ → Ionic

Lewis Structures

Lewis structures are diagrams that show the arrangement of valence electrons among atoms in a molecule. They help predict molecular shape, bond order, and the presence of lone pairs.

• Count total valence electrons for all atoms.

• Arrange atoms with the least electronegative atom in the center (except hydrogen).

• Distribute electrons to satisfy the octet rule (or duet for hydrogen).

• Use double or triple bonds if necessary to complete octets.

Example: The Lewis structure for CO2 is O=C=O.

Formal Charge

Formal charge is used to determine the most stable Lewis structure for a molecule. It is calculated as:

$\text{Formal charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2} \times \text{Bonding electrons}$

The best Lewis structure has the smallest formal charges and places negative charges on the most electronegative atoms.

Resonance Structures

Some molecules cannot be represented by a single Lewis structure. Resonance structures are multiple valid Lewis structures for the same molecule, differing only in the placement of electrons.

• Actual structure is a resonance hybrid, an average of all resonance forms.

• Example: The carbonate ion (CO32-) has three resonance structures.

Molecular Geometry and VSEPR Theory

VSEPR Theory (Valence Shell Electron Pair Repulsion)

VSEPR theory predicts the shape of a molecule based on the repulsion between electron pairs around a central atom. Electron pairs (bonding and lone pairs) arrange themselves as far apart as possible to minimize repulsion.

• Draw the Lewis structure.

• Count the number of electron domains (bonding pairs + lone pairs) around the central atom.

• Assign molecular geometry based on the number of domains.

Common Geometries:

Electron Domains	Geometry	Bond Angle
2	Linear	180°
3	Trigonal planar	120°
4	Tetrahedral	109.5°
5	Trigonal bipyramidal	90°, 120°
6	Octahedral	90°

Bond Angles and Molecular Shape

The presence of lone pairs affects bond angles, making them smaller than the ideal values. For example, in water (H2O), the bond angle is 104.5° due to two lone pairs on oxygen.

Polarity of Molecules

Molecular polarity depends on both bond polarity and molecular geometry. A molecule is polar if it has a net dipole moment (asymmetrical charge distribution).

• Nonpolar molecules: Symmetrical geometry cancels dipoles (e.g., CO2).

• Polar molecules: Asymmetrical geometry results in a net dipole (e.g., H2O).

Naming Compounds

Ionic Compounds

• Name the cation (metal) first, then the anion (nonmetal) with '-ide' ending.

• For transition metals, indicate the charge with Roman numerals (e.g., FeCl3: iron(III) chloride).

Covalent Compounds

• Use prefixes to indicate the number of each atom (e.g., CO2: carbon dioxide).

Acids

• Binary acids: 'Hydro-' prefix + root + '-ic acid' (e.g., HCl: hydrochloric acid).

• Oxoacids: Based on polyatomic ion name (e.g., H2SO4: sulfuric acid).

Practice Problems and Applications

Sample Table: Molecular Properties

Formula	Lewis Structure	VSEPR & Bond Angle(s)	Polarity	Name
CO32-	O=C-O- (resonance)	Trigonal planar, 120°	Non-polar	Carbonate
HCO3-	See image	Trigonal planar & bent, 120° and 104.5°	Polar	Hydrogen carbonate
CH3NH2	See image	Tetrahedral & trigonal pyramidal, 109.5° and 107°	Polar	Methylamine
CNCl	See image	Linear, 180°	Polar	Cyanogen chloride
PO33-	See image	Trigonal pyramidal, 107°	Polar	Phosphite
C2H6	See image	Tetrahedral, 109.5°	Non-polar	Ethane
H2O2	See image	Bent, 104.5°	Polar	Hydrogen peroxide
SO2Cl2	See image	Tetrahedral, 109.5°	Non-polar	Thionyl chloride

Additional info: Lewis structures and VSEPR assignments are based on standard conventions. See textbook or lecture slides for detailed drawings.

Summary Table: VSEPR Shapes and Bond Angles

Shape	Bond Angle	Example
Linear	180°	CO2
Trigonal planar	120°	BF3
Tetrahedral	109.5°	CH4
Trigonal pyramidal	107°	NH3
Bent	104.5°	H2O

Key Concepts and Skills

• Distinguish between ionic, covalent, and metallic bonds.

• Draw and interpret Lewis structures for molecules and polyatomic ions.

• Assign formal charges and identify resonance structures.

• Predict molecular shapes using VSEPR theory.

• Determine bond angles and molecular polarity.

• Name ionic and covalent compounds, including acids.

• Relate molecular structure to physical properties such as melting point and conductivity.