Chemical Bonding I: Covalent Bonds, Lewis Structures, VSEPR, and Molecular Polarity

Chapter 5: Chemical Bonding I

Distinguishing Between Polar and Nonpolar Covalent Bonds

Chemical bonds can be classified based on how electrons are shared between atoms. Understanding the difference between polar and nonpolar covalent bonds is essential for predicting molecular properties.

• Nonpolar Covalent Bond: Two atoms share electrons evenly or almost evenly. This typically occurs between atoms of similar electronegativity. Example: The bond in H2 or Cl2.

• Polar Covalent Bond: Two atoms share electrons unevenly due to a difference in electronegativity. The more electronegative atom attracts the shared electrons more strongly. Example: The bond in HCl or HF.

Length and Strength of Covalent Bonds

The properties of covalent bonds depend on the number of shared electron pairs and the atoms involved.

• Double bonds are shorter and stronger than single bonds.

• Triple bonds are shorter and stronger than double bonds.

• The bond length is the internuclear distance at which the potential energy between two atoms is minimized.

Bond Energy: The energy required to break one mole of a covalent bond in the gas phase. Bond Length: The distance between the nuclei of two bonded atoms.

Formula:

$\text{Bond Order} = \frac{\text{Number of bonding electron pairs}}{\text{Number of bonds}}$

Drawing Lewis Electron Dot Structures

Lewis structures represent the arrangement of valence electrons in molecules. They are essential for understanding bonding and molecular geometry.

• Count the total number of valence electrons for all atoms in the molecule.

• Connect atoms with single bonds (pairs of electrons).

• Distribute remaining electrons to satisfy the octet rule (8 electrons around each atom, except H which needs 2).

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

• Consider resonance structures if electrons can be delocalized.

• Assign formal charges to atoms to determine the most stable structure.

Formal Charge Formula:

$\text{Formal Charge} = (\text{Number of valence electrons}) - (\text{Number of bonding electrons}) - (\text{Number of nonbonding electrons})$

Resonance and Resonance Stabilization

Some molecules can be represented by more than one valid Lewis structure. These are called resonance structures, and the true structure is a hybrid of all possible forms.

• Resonance: Delocalization of electrons across multiple atoms increases stability.

• Resonance hybrids have lower energy than any single resonance structure.

Valence-Shell Electron-Pair Repulsion (VSEPR) Theory

VSEPR theory predicts the 3D shape of molecules based on the repulsion between electron pairs around a central atom.

• Electron domains (bonding pairs and lone pairs) arrange themselves to minimize repulsion.

• The number of electron domains determines the electron geometry.

• Lone pairs compress bond angles compared to ideal geometries.

Electron Geometries and Molecular Shapes

The following table summarizes electron geometries, hybridization, and molecular shapes for molecules with 2–6 electron domains:

Electron Domains	Hybridization	Electron Geometry	0 Lone Pairs	1 Lone Pair	2 Lone Pairs	3 Lone Pairs	4 Lone Pairs
2	sp	Linear	Linear
3	sp2	Trigonal Planar	Trigonal Planar	Bent
4	sp3	Tetrahedral	Tetrahedral	Trigonal Pyramidal	Bent
5	sp3d	Trigonal Bipyramidal	Trigonal Bipyramidal	Seesaw	T-shaped	Linear
6	sp3d2	Octahedral	Octahedral	Square Pyramidal	Square Planar

Additional info: Bond angles decrease as the number of lone pairs increases due to greater electron repulsion.

Bond Angles and Molecular Shape

The shape of a molecule and the number of lone pairs affect bond angles.

• In trigonal planar and tetrahedral geometries, more lone pairs compress bond angles.

• Examples:

  • CH4 (tetrahedral, no lone pairs): bond angle ≈ 109.5°

  • NH3 (tetrahedral, 1 lone pair): bond angle ≈ 107°

  • H2O (tetrahedral, 2 lone pairs): bond angle ≈ 104.5°

Polarity: Polar Bonds and Polar Molecules

Molecular polarity depends on both the presence of polar bonds and the geometry of the molecule.

• Polar Bond: A bond between two atoms with a large difference in electronegativity.

• Dipole Moment: A vector representing the polarity of a bond.

• Polar Molecule: A molecule with a significant net dipole moment.

• A molecule may contain polar bonds but still be nonpolar if the dipole moments cancel due to molecular geometry.

Examples of Polar Molecules: H2O, CH2O, CH3Cl, NH3, SO2

Examples of Nonpolar Molecules: CO2, SO3, SiH4, CH4, CCl4, CH3CH3

Key Terms and Definitions

Understanding the following terms is essential for mastering chemical bonding concepts:

• Nonpolar Covalent Bond: Even distribution of charge; electrons shared equally.

• Polar Covalent Bond: Unequal sharing of electrons; results in partial charges.

• Bond Energy: Energy required to break one mole of a bond.

• Octet Rule: Atoms tend to have eight electrons in their valence shell.

• Bond Length: Distance between nuclei of bonded atoms.

• Bonding Pair: Pair of electrons shared between two atoms.

• Lone Pair: Pair of electrons not shared between atoms.

• Electronegativity: Ability of an atom to attract electrons in a bond.

• Resonance: Delocalization of electrons across multiple atoms.

• Resonance Structure: Different valid Lewis structures for the same molecule.

• Resonance Stabilization: Increased stability due to electron delocalization.

• Formal Charge: Calculated charge on an atom in a molecule. $\text{Formal Charge} = (\text{Number of valence electrons}) - (\text{Number of bonding electrons}) - (\text{Number of nonbonding electrons})$

• Bond Order: Number of bonds between atoms (single = 1, double = 2, triple = 3).

• Bond Angle: Angle between two bonds joining three atoms.

• Electron Geometry: Arrangement of electron domains around a central atom.

• Molecular Geometry: 3D arrangement of atoms in a molecule.

• Bond Dipole: Separation of electrical charge in a covalent bond.