Basic Concepts of Chemical Bonding

Chemical Bonds, Lewis Symbols, and the Octet Rule

• Chemical Bond: A strong attraction between two atoms or ions, responsible for holding substances together.

• Ionic Bond: Electrostatic attraction between ions of opposite charges (e.g., NaCl).

• Covalent Bond: Bond formed by sharing electrons between atoms (e.g., H2O).

• Metallic Bond: Found in metals; involves freely moving electrons within a lattice of metal cations.

• Lewis Symbols: Represent the valence electrons of an atom using dots around the element symbol.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve eight electrons in their valence shell (noble gas configuration).

Ionic Bonding

• Energetics of Ionic Bond Formation: Formation involves ionization energy (energy to remove electrons from metals), electron affinity (energy released when nonmetals gain electrons), and lattice energy (energy released when ions form a solid lattice).

• Electron Configuration of Ions: Groups 1A, 2A, 3A form cations (+1, +2, +3); Groups 5A, 6A, 7A form anions (−3, −2, −1).

• Transition Metals: Lose s electrons first, then d electrons; common charges are +1, +2, +3.

Lewis Structures

• Definition: Diagrams showing the arrangement of valence electrons in a molecule.

• Multiple Bonds: Double bonds (two pairs shared, four electrons), triple bonds (three pairs shared, six electrons).

Bond Polarity and Electronegativity

• Bond Polarity: Describes how electrons are shared between atoms.

• Nonpolar Covalent Bond: Electrons shared equally (e.g., Cl2).

• Polar Covalent Bond: Electrons shared unequally (e.g., HCl).

• Electronegativity: Ability of an atom to attract shared electrons. Increases across a period, decreases down a group. Highest: F > O > Cl > N.

• Bond Type Determination: Large electronegativity difference → ionic bond; small or zero difference → nonpolar covalent; intermediate difference → polar covalent.

• Illustrating Bond Polarity: Use δ+ and δ− or an arrow pointing toward the more electronegative atom.

Drawing Lewis Structures

• 1. Count total valence electrons (add for negative ions, subtract for positive ions).

• 2. Arrange atoms in a skeletal structure.

• 3. Assign octets to outer atoms (except H).

• 4. Place remaining electrons on the central atom.

• 5. If needed, create double/triple bonds to satisfy octet rule.

Formal Charge

• Calculation:

• Best Lewis Structure: Smallest formal charges, negative charges on most electronegative atoms.

Resonance Structures

• When more than one valid Lewis structure exists, the actual structure is a resonance hybrid of all possibilities (e.g., O3).

Exceptions to the Octet Rule

• Odd Electron Molecules: Molecules with an odd number of electrons (e.g., NO, ClO2).

• Incomplete Octet: B and Be can have fewer than 8 electrons (B: 6, Be: 4).

• Expanded Octet: Elements in period 3 or higher (e.g., P, S) can have more than 8 electrons.

Strengths of Covalent Bonds

• Bond Enthalpy: Energy required to break a bond; always positive.

• Bond Enthalpy and Reaction Enthalpy:

• Bond Length: Stronger bonds are shorter; single bonds are longer and weaker than double or triple bonds.

Molecular Geometry and Bonding Theories

Molecular Shapes and the VSEPR Model

• VSEPR Model: Valence Shell Electron Pair Repulsion theory; electron pairs arrange themselves to minimize repulsion.

• Steps to Determine Shape:

  1. Draw the correct Lewis structure.

  2. Count electron domains (atoms + lone pairs) around the central atom.

  3. Assign electron domain geometry:

  Electron Domains	Geometry
  2	Linear
  3	Trigonal planar
  4	Tetrahedral
  5	Trigonal bipyramidal
  6	Octahedral

• If lone pairs are present, molecular shape is determined by ignoring their positions.

Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles

• Lone pairs occupy more space and decrease bond angles between atoms.

• Multiple bonds (double/triple) also affect bond angles due to increased electron density.

Molecules with Expanded Valence Shells

• Atoms in period 3 or higher can have more than 8 electrons, leading to shapes like trigonal bipyramidal and octahedral.

Molecular Shape and Molecular Polarity

• Polarity Determination: A molecule is polar if it contains polar bonds and the bond dipoles do not cancel due to molecular shape.

• Dipole Moment: Zero for nonpolar molecules; nonzero for polar molecules.

• Examples: CO2 is nonpolar (linear, dipoles cancel); OCS is polar (dipoles do not cancel).

• Molecules Always Nonpolar:

  • Linear (A2X)

  • Trigonal planar (A3X)

  • Tetrahedral (A4X)

  • Trigonal bipyramidal (A5X, A3B2X with identical equatorial and axial atoms)

  • Octahedral (A6X, A4B2X with identical atoms)

Covalent Bonding and Orbital Overlap

• Valence Bond Theory: Bonds form from the overlap of atomic orbitals (e.g., H2: s-s overlap; F2: p-p overlap).

• Bond formation lowers potential energy and stabilizes the molecule.

Hybrid Orbitals

• Hybridization: Mixing atomic orbitals to form new hybrid orbitals that explain molecular shapes.

• Number of electron domains determines hybridization:

Multiple Bonds: Sigma and Pi Bonds

• Sigma (σ) Bonds: End-to-end overlap of orbitals; all single bonds are sigma bonds.

• Pi (π) Bonds: Sideways overlap of unhybridized p orbitals; present in double and triple bonds.

• Bond Composition:

  • Single bond: 1 σ

  • Double bond: 1 σ + 1 π

  • Triple bond: 1 σ + 2 π

• Delocalized π Bonding: In molecules with three or more atoms in a plane, π electrons can be delocalized, increasing stability (e.g., benzene).

Gases

Characteristics of Gases

• Gases expand to fill their container.

• Volume changes with pressure (compressible).

• Gases mix freely to form homogeneous mixtures.

• Gas molecules are far apart; most of the volume is empty space.

• Molecules move rapidly and randomly; intermolecular forces are negligible.

• Described by four variables: pressure (P), volume (V), temperature (T), and number of moles (n).

Pressure

• Definition: (force per unit area)

• Units: mm Hg (torr), atmospheres (atm), pascals (Pa), kilopascals (kPa)

• Measurement: Barometer (atmospheric pressure), manometer (gas pressure in a container)

The Gas Laws

• Boyle's Law (P-V): At constant T,

• Charles's Law (V-T): At constant P,

• Avogadro's Law (V-n): At constant T and P,

• Combined Gas Law:

The Ideal Gas Equation

• Equation:

• Variables: P (atm), V (L), n (mol), T (K), R = 0.0821 L·atm/(mol·K)

• Ensure all units are compatible with R before solving.

Further Applications of the Ideal Gas Law

• Gas Density: , where M is molar mass.

• Molar Mass from Density:

• Volumes in Chemical Reactions: At STP (0°C, 1 atm), 1 mol gas = 22.4 L.

Gas Mixtures and Partial Pressures

• Dalton's Law:

• Mole Fraction: , where

• Collecting Gases Over Water: Subtract vapor pressure of water from total pressure to get pressure of collected gas.

Kinetic-Molecular Theory of Gases

• Gas molecules are in constant, random motion.

• Collisions are elastic; no energy is lost.

• Volume of molecules is negligible compared to container volume.

• No intermolecular forces between molecules.

• Root Mean Square Speed: (M in kg/mol)

Applications of Kinetic Theory

• Boyle's Law: At constant T, increasing V decreases collision frequency, lowering P.

• Temperature Increase at Constant V: Molecules move faster, hit walls harder, increasing P.

Molecular Effusion and Diffusion

• Effusion: Escape of gas through a tiny hole into a vacuum.

• Diffusion: Spread of gas throughout space or another gas.

• Graham's Law: Rate of effusion/diffusion is inversely proportional to the square root of molar mass:

• Rates and times are inversely proportional.

• Mean Free Path: Average distance a molecule travels before colliding; decreases as pressure increases.

Example: Calculate the volume occupied by 2.0 mol of an ideal gas at 27°C and 1.00 atm:

Additional info: For more practice, refer to sample exercises in your textbook as indicated in the notes (e.g., Sample Exercises 8.6–8.9, 9.4, 10.1–10.2, 10.9).