4.2 Covalent bonding and Lewis structures

Covalent bonds form when atoms share electrons, creating stable molecules. This sharing leads to specific bond properties, like length and angle, which influence molecular structure and behavior.

Lewis structures visually represent electron arrangements in molecules, helping predict their shape and properties. Understanding these concepts is crucial for grasping how atoms interact and form compounds.

Covalent bond formation and characteristics

Covalent bond formation

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• Covalent bonds form when atoms share one or more pairs of electrons
• Typically occurs between nonmetals (carbon, oxygen, nitrogen)
• Shared electrons are attracted to the nuclei of both atoms
• Attraction results in a stable bond between the atoms
• Number of shared electron pairs determines bond order (single, double, triple)

Covalent bond properties

• Covalent bonds are directional with specific bond lengths and angles
• Bond length is the distance between the nuclei of the bonded atoms
• Bond angle is the angle formed by the imaginary lines connecting the nuclei
• Strength of a covalent bond depends on bond order and electronegativity difference
• Higher bond order (double or triple bonds) results in stronger bonds
• Greater electronegativity difference leads to more polar covalent bonds

Lewis structures for molecules and ions

Constructing Lewis structures

• Lewis structures represent the arrangement of atoms and electrons in a molecule or polyatomic ion
• Each atom is represented by its chemical symbol
• Valence electrons are shown as dots surrounding the symbol
• Atoms are connected by lines representing shared electron pairs (covalent bonds)
• Octet rule states that atoms tend to achieve a stable electron configuration with eight valence electrons (except hydrogen, which requires two)
• Lone pairs are unshared electron pairs not involved in bonding

Guidelines for drawing Lewis structures

• Place the least electronegative atom at the center (usually carbon or the least electronegative nonmetal)
• Arrange the remaining atoms around the central atom
• Distribute the available valence electrons as lone pairs or bonding pairs
• Assign formal charges to atoms to determine the most stable arrangement
• Formal charge = [# of valence electrons in free atom] - [# of non-bonding electrons] - (1/2)[# of bonding electrons]
• Aim for the most stable structure with the lowest formal charges and satisfying the octet rule

Molecular geometry and polarity

VSEPR theory and molecular geometry

• VSEPR (Valence Shell Electron Pair Repulsion) theory predicts molecular geometry based on electron domains (bonding and nonbonding) around the central atom
• Electron domains repel each other and arrange to minimize repulsion
• Five basic electron domain geometries: linear (2 domains), trigonal planar (3 domains), tetrahedral (4 domains), trigonal bipyramidal (5 domains), octahedral (6 domains)
• Molecular geometry is determined by the arrangement of atoms, not considering lone pairs
• Lone pairs affect bond angles and overall molecular shape (bent, trigonal pyramidal, seesaw, T-shaped, square pyramidal)

Molecular polarity

• Polarity depends on geometry and electronegativity differences between bonded atoms
• Polar molecules have an uneven charge distribution due to polar covalent bonds and asymmetric geometry (water, ammonia)
• Nonpolar molecules have a symmetric charge distribution, either due to absence of polar covalent bonds (O2, N2) or symmetric arrangement of polar bonds (CO2, BF3)
• Dipole moment is a vector quantity that measures the magnitude and direction of polarity
• Molecules with a net dipole moment are polar, while those with a zero net dipole moment are nonpolar

Resonance and molecular stability

Resonance structures

• Resonance occurs when a molecule or polyatomic ion can be represented by multiple valid Lewis structures
• Actual structure is a hybrid of these resonance structures, with electron density distributed among them
• Resonance structures are not different molecules, but different representations of the same molecule or ion
• Resonance structures are connected by double-headed arrows, indicating they contribute to the overall electronic structure

Resonance and stability

• More resonance structures lead to greater stability due to electron delocalization
• Delocalization lowers the overall energy of the molecule or ion compared to any single resonance structure
• Stability gained through resonance is called resonance energy
• Resonance affects bond lengths, with actual bond lengths intermediate between single and double bonds
• Examples of molecules with significant resonance: benzene, nitrate ion, carbonate ion