Chemical Bonds, Electron-Dot Structures, VSEPR, Valence Bond Theory, and Intermolecular Forces

Chemical Bonds

Definition and Types of Chemical Bonds

Chemical bonds are forces that hold groups of atoms together, allowing them to function as a unit. The electrons in atoms form these connections. A bond forms if the energy of the aggregate is lower than that of the separated atoms. The energy required to break a chemical bond is called the bond energy.

• Octet Rule: Main-group elements tend to undergo reactions that leave them with eight outer-shell electrons, achieving a noble-gas electron configuration.

• Main Types of Chemical Bonds:

  • Ionic Bonds: Formed by the transfer of electrons from one atom to another, resulting in electrostatic attraction between cations and anions.

  • Covalent Bonds: Formed by the sharing of electrons between two nonmetal atoms.

  • Metallic Bonds: Involve a 'sea' of delocalized electrons around metal cations.

Chemical Bonds – Metallic Bonding

Properties and Electron-Sea Model

• Occurs in metals.

• High thermal and electrical conductivity.

• Great strength, malleability, ductility, and luster.

• Electron-Sea Model: Metal cations are immersed in a sea of delocalized electrons free to move throughout the crystal, explaining conductivity and malleability.

Chemical Bonds – Ionic Bonding

Formation and Characteristics

• One atom transfers an electron to another, creating an electrostatic attraction between positively charged cations and negatively charged anions.

• Typically forms between metals (low ionization energy) and non-metals (high electron affinity).

• Example: $\mathrm{Na} + \mathrm{Cl} \rightarrow \mathrm{Na}^+ + \mathrm{Cl}^-$

Chemical Bonds – Covalent Bonding

Formation and Bond Polarity

• Involves sharing of electrons between two atoms, usually nonmetals.

• Bond Polarity: Arises from differences in electronegativity (EN). The greater the EN difference, the more polar the bond.

• Polar Covalent Bond: Unequal sharing of electrons.

• Nonpolar Covalent Bond: Equal sharing of electrons.

Chemical Bonds – Comparing Ionic and Covalent Compounds

Physical Properties and Structure

	NaCl (Ionic)	HCl (Covalent)
Physical Appearance	White solid	Colorless gas
Melting Point	801 °C	-115 °C
Boiling Point	1465 °C	-84.9 °C

• Ionic compounds are high-melting solids due to strong ionic bonds in a lattice structure.

• Covalent molecular compounds are low-melting solids, liquids, or gases, consisting of discrete molecules.

• Covalent Network Solids: Atoms are held together by covalent bonds in a giant three-dimensional array (e.g., diamond, quartz).

Electron-Dot (Lewis) Structures

Drawing Lewis Structures

• Valence electrons are distributed as bonding pairs (shared) and non-bonding pairs (lone pairs).

• Steps:

  1. Sum valence electrons from all atoms (add for negative charge, subtract for positive charge).

  2. Identify the central atom (usually the least electronegative).

  3. Place one bond (two electrons) between each pair of atoms.

  4. Complete octets for atoms bonded to the central atom.

  5. Place leftover electrons on the central atom.

  6. If not enough electrons for octet, try multiple bonds.

Exceptions to the Octet Rule: Some elements (e.g., P, S, Cl) can accommodate more than eight electrons (expanded octet), especially in period 3 or higher.

Resonance Structures

Concept and Representation

• Resonance structures are alternative Lewis structures for the same molecule, representing delocalized electrons.

• The actual structure is a resonance hybrid, a weighted average of all contributors.

• Example: Ozone ($\mathrm{O}_3$) has two resonance structures with different locations for the double bond.

Formal Charge

Definition and Calculation

• Formal Charge (FC): Difference between the number of valence electrons in the free atom and the number assigned in the Lewis structure.

$\text{Formal charge} = \text{(Valence electrons in free atom)} - \text{(Nonbonding electrons)} - \frac{1}{2}\text{(Bonding electrons)}$

• The most stable structure has the lowest formal charges and places negative charges on the most electronegative atoms.

The VSEPR Model: Valence Shell Electron Pair Repulsion

Predicting Molecular Geometry

• Pairs of electrons repel each other; the shape of a molecule is determined by minimizing these repulsions.

• Steps:

  1. Draw the Lewis structure.

  2. Determine the electron domain geometry (arrangement of electron pairs around the central atom).

  3. Determine the molecular geometry (arrangement of atoms, not including lone pairs).

• Repulsion order: lone pair-lone pair > lone pair-bonding pair > bonding pair-bonding pair.

Number of Charge Clouds	Electron Geometry	Example
2	Linear	CO2
3	Trigonal planar	BF3
4	Tetrahedral	CH4
5	Trigonal bipyramidal	PCl5
6	Octahedral	SF6

Valence Bond Theory

Bond Formation and Orbital Overlap

• Covalent bonds are formed by the overlap of atomic orbitals, each containing one electron of opposite spin.

• Each bonded atom maintains its own atomic orbitals, but the electron pair is shared by both atoms.

• The greater the orbital overlap, the stronger the bond.

• Sigma (σ) bonds: Formed by head-on overlap (s-s, s-p, or p-p).

• Pi (π) bonds: Formed by side-on overlap of p orbitals.

Hybridization of Atomic Orbitals

Concept and Types

• Atomic orbitals (s and p) can combine mathematically to form new, equivalent hybrid orbitals for bonding.

• Types of Hybridization:

  • sp3: Tetrahedral geometry, bond angle 109.5° (e.g., CH4).

  • sp2: Trigonal planar geometry, bond angle 120° (e.g., C2H4).

  • sp: Linear geometry, bond angle 180° (e.g., C2H2).

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

  • sp3d2: Octahedral geometry (e.g., SF6).

• Hybrid orbitals also accommodate lone pairs (e.g., sp3 in NH3 and H2O).

Intermolecular Forces

Types and Relative Strengths

• Intermolecular forces are attractive forces between molecules, much weaker than intramolecular (covalent/ionic) bonds.

• They determine physical properties such as boiling and melting points.

Force	Strength (kJ/mol)	Characteristics
Ion-dipole	Moderate (10–50)	Occurs between ions and polar solvents
Dipole-dipole	Weak (3–4)	Occurs between polar molecules
London dispersion	Weak (1–10)	Occurs between all molecules; strength depends on size, polarizability
Hydrogen bond	Moderate (10–40)	Occurs between molecules with O—H, N—H, or F—H bonds

• Ion-Dipole Forces: Attraction between an ion and the partial charges on a polar molecule.

• Dipole-Dipole Forces: Attraction between permanent dipoles of polar molecules.

• London Dispersion Forces: Temporary attractive forces due to instantaneous dipoles in all molecules, especially significant in large, polarizable atoms/molecules.

• Hydrogen Bonding: Special dipole-dipole interaction between H bonded to N, O, or F and a lone pair on another N, O, or F atom.

Polar Covalent Bonds and Dipole Moments

Electronegativity and Molecular Polarity

• Electronegativity (EN): Tendency of an atom to attract electrons in a bond. Increases across a period, decreases down a group.

• Bond Dipole: Arises when two atoms with different EN form a bond, resulting in partial charges (δ+ and δ−).

• Molecular Dipole Moment: Vector sum of all bond dipoles in a molecule. If they do not cancel, the molecule is polar.

• Examples:

  • Ammonia (NH3) and water (H2O) are polar (net dipole moment).

  • Carbon dioxide (CO2) and tetrachloromethane (CCl4) are nonpolar (bond dipoles cancel).

How to Determine Molecular Polarity:

1. Draw the Lewis structure.

2. If all regions of electron density are identical, the molecule is nonpolar.

3. If not, the molecule is polar.

Summary Table: Comparison of Intermolecular Forces

Type of Force	Relative Strength	Occurs Between
Ion-dipole	Moderate	Ions and polar molecules
Dipole-dipole	Weak	Polar molecules
London dispersion	Weak	All molecules
Hydrogen bond	Moderate	H bonded to N, O, or F

Additional info: These notes are based on lecture slides and textbook references (McMurry), and cover foundational concepts in chemical bonding, molecular structure, and intermolecular forces, suitable for General Chemistry exam preparation.