Chemical Bonding

Chemical bonding is the force that holds atoms together in compounds, creating the incredible diversity of materials around us. From the salt in our food to the water we drink, from the metals in our buildings to the DNA in our cells, everything depends on how atoms bond together.

Why Do Atoms Form Bonds?

The Drive for Stability

Energy Considerations

• Lower energy state: Bonded atoms are more stable than isolated atoms
• Energy release: Bond formation releases energy
• Energy input needed: Breaking bonds requires energy
• Minimum energy principle: Systems naturally seek lowest energy

The Octet Rule

• Noble gas configuration: Atoms want 8 electrons in outer shell
• Exception: Hydrogen wants 2 electrons (like helium)
• Methods: Gain, lose, or share electrons
• Driving force: Achieve stable electron configuration

Electronegativity: The Key Factor

Definition

• Electronegativity: Ability to attract electrons in a bond
• Scale: Pauling scale (0.7 to 4.0)
• Trends: Increases across period, up group
• Most electronegative: Fluorine (4.0)

Bond Type Prediction

• Electronegativity difference > 1.7: Ionic bonding
• Difference 0.4-1.7: Polar covalent bonding
• Difference < 0.4: Nonpolar covalent bonding
• Between metals: Metallic bonding

Understanding electronegativity trends

Why electronegativity increases across a period:

• Nuclear charge increases: More protons attract electrons more strongly
• Same electron shell: Electrons aren't much farther from nucleus
• Result: Stronger attraction for bonding electrons

Why electronegativity decreases down a group:

• Larger atoms: Bonding electrons farther from nucleus
• Electron shielding: Inner electrons reduce nuclear attraction
• Result: Weaker attraction for bonding electrons

Electronegativity values (selected elements):

• F: 4.0 (most electronegative)
• O: 3.5, N: 3.0, Cl: 3.0
• C: 2.5, H: 2.1
• Na: 0.9, K: 0.8, Cs: 0.7 (least electronegative)

Ionic Bonding

Formation of Ionic Bonds

Electron Transfer

• Between: Metal and non-metal
• Process: Metal loses electrons, non-metal gains electrons
• Result: Positive and negative ions form
• Attraction: Opposite charges attract (electrostatic force)

Example: Sodium Chloride (NaCl)

• Sodium atom: Na → Na⁺ + e⁻ (loses 1 electron)
• Chlorine atom: Cl + e⁻ → Cl⁻ (gains 1 electron)
• Both achieve: Noble gas configuration
• Na⁺: Same as neon (2,8)
• Cl⁻: Same as argon (2,8,8)

Properties of Ionic Compounds

Physical Properties

• Crystal structure: Regular arrangement of ions
• High melting/boiling points: Strong electrostatic forces
• Brittle: Shifting layers causes like charges to repel
• Hard: Strong ionic bonds resist deformation

Electrical Properties

• Insulators when solid: Ions fixed in crystal lattice
• Conduct when molten: Ions free to move
• Conduct when dissolved: Ions dissociate in solution
• Electrolytes: Solutions conduct electricity

Solubility

• Polar solvents: Often soluble in water
• Hydration: Water molecules surround ions
• Nonpolar solvents: Usually insoluble
• Lattice energy: Affects solubility

Ionic Crystal Structures

Types of Structures

• Rock salt (NaCl): 6:6 coordination
• Cesium chloride (CsCl): 8:8 coordination
• Fluorite (CaF₂): Different sized ions
• Zinc blende (ZnS): Tetrahedral coordination

Factors Affecting Structure

• Size ratio: Relative sizes of cations and anions
• Charge ratio: Ratio of positive to negative charges
• Coordination number: Number of nearest neighbors
• Efficiency: Maximum density packing

Covalent Bonding

Formation of Covalent Bonds

Electron Sharing

• Between: Non-metals (similar electronegativity)
• Process: Atoms share electron pairs
• Result: Both atoms achieve stable configuration
• Bond formation: Overlap of atomic orbitals

Example: Hydrogen Molecule (H₂)

• Two hydrogen atoms: Each has 1 electron
• Sharing: Both electrons shared between nuclei
• Result: Each hydrogen "feels" 2 electrons
• Stability: Both achieve helium configuration

Types of Covalent Bonds

1. Nonpolar Covalent Bonds

• Equal sharing: Electrons shared equally
• Same elements: H₂, Cl₂, O₂, N₂
• Similar electronegativity: Difference < 0.4
• Example: C-H bonds in methane

2. Polar Covalent Bonds

• Unequal sharing: One atom attracts electrons more
• Partial charges: δ⁺ and δ⁻ (delta plus and minus)
• Electronegativity difference: 0.4-1.7
• Example: H-Cl, C-O, N-H bonds

3. Coordinate Covalent Bonds

• Definition: One atom provides both electrons
• Donor atom: Has lone pair of electrons
• Acceptor atom: Has empty orbital
• Example: NH₃ + H⁺ → NH₄⁺

Multiple Bonds

Single Bonds

• One shared pair: Two electrons shared
• Sigma bond (σ): Orbital overlap along bond axis
• Examples: H-H, C-C, C-H
• Rotation: Free rotation around single bonds

Double Bonds

• Two shared pairs: Four electrons shared
• One sigma + one pi (π): Different orbital overlaps
• Examples: C=C, C=O, O=O
• Stronger and shorter: Than single bonds

Triple Bonds

• Three shared pairs: Six electrons shared
• One sigma + two pi: Multiple orbital overlaps
• Examples: N≡N, C≡C, C≡O
• Strongest and shortest: Very stable bonds

Lewis Structures

Drawing Rules

Common Patterns

• Hydrogen: Always forms 1 bond (2 electrons)
• Carbon: Usually forms 4 bonds (8 electrons)
• Nitrogen: Usually forms 3 bonds + 1 lone pair
• Oxygen: Usually forms 2 bonds + 2 lone pairs

Drawing Lewis structures step by step

Example: Water (H₂O) Step 1 - Count valence electrons:

• Oxygen: 6 valence electrons
• Hydrogen: 1 valence electron × 2 = 2
• Total: 6 + 2 = 8 valence electrons

Step 2 - Arrange atoms:

• Oxygen in center (less electronegative than F, more than H)
• Two hydrogens around oxygen

Step 3 - Connect with bonds:

• H-O-H uses 4 electrons (2 bonds)
• Remaining electrons: 8 - 4 = 4 electrons

Step 4 - Complete octets:

• Hydrogen complete with 2 electrons each
• Oxygen needs 8 total: has 4 in bonds + 4 as lone pairs
• Final structure: H-O-H with 2 lone pairs on oxygen

Molecular Geometry

VSEPR Theory

Basic Principle

• VSEPR: Valence Shell Electron Pair Repulsion
• Key idea: Electron pairs repel each other
• Arrangement: Electron pairs arrange to minimize repulsion
• Shape determination: Geometry based on electron pair positions

Electron Pair Types

• Bonding pairs: Electrons in covalent bonds
• Lone pairs: Non-bonding electron pairs
• Multiple bonds: Count as single unit
• Repulsion strength: Lone pair > bonding pair

Common Molecular Geometries

2 Electron Pairs (Linear)

• Angle: 180°
• Example: BeCl₂, CO₂
• Shape: Straight line

3 Electron Pairs

• Trigonal planar: 3 bonding pairs, 120° angles (BF₃)
• Bent: 2 bonding + 1 lone pair, <120° angle (SO₂)

4 Electron Pairs

• Tetrahedral: 4 bonding pairs, 109.5° angles (CH₄)
• Trigonal pyramidal: 3 bonding + 1 lone pair (NH₃)
• Bent: 2 bonding + 2 lone pairs, 104.5° (H₂O)

5 and 6 Electron Pairs

• Trigonal bipyramidal: 5 pairs (PCl₅)
• Octahedral: 6 pairs (SF₆)
• Various shapes: With lone pairs

Molecular Polarity

Determining Polarity

• Bond polarity: Individual bonds polar or nonpolar?
• Molecular geometry: Do polar bonds cancel out?
• Symmetry: Symmetric molecules often nonpolar
• Dipole moment: Net separation of charge

Examples

• CO₂: Polar bonds, linear geometry → nonpolar molecule
• H₂O: Polar bonds, bent geometry → polar molecule
• CCl₄: Polar bonds, tetrahedral → nonpolar molecule
• NH₃: Polar bonds, pyramidal → polar molecule

Metallic Bonding

The Electron Sea Model

Structure

• Metal cations: Positive ions in regular arrangement
• Delocalized electrons: "Sea" of mobile electrons
• Non-directional bonding: Electrons not localized between atoms
• Electrostatic attraction: Cations attracted to electron sea

Electron Mobility

• Free movement: Electrons can move throughout structure
• No specific pairs: Electrons belong to entire structure
• Flexible bonding: Bonds can form and break easily
• Collective behavior: All electrons contribute to bonding

Properties from Metallic Bonding

Electrical Conductivity

• Mobile electrons: Can carry electric current
• Excellent conductors: Silver, copper, gold
• Temperature effect: Conductivity decreases with heat
• Direction independence: Conduct in all directions

Thermal Conductivity

• Heat transfer: Mobile electrons carry thermal energy
• Good thermal conductors: Same metals as electrical
• Applications: Heat sinks, cooking pots

Mechanical Properties

• Malleability: Can be hammered into sheets
• Ductility: Can be drawn into wires
• Mechanism: Layers can slide without breaking bonds
• Non-directional bonding: Allows plastic deformation

Optical Properties

• Metallic luster: Shiny appearance
• Light interaction: Electrons absorb and re-emit light
• Reflectivity: Good reflectors of light
• Opacity: Don't transmit light

Factors Affecting Metallic Bond Strength

Number of Valence Electrons

• More electrons: Stronger bonding
• Example: Mg (2 valence) harder than Na (1 valence)
• Electron density: Higher density = stronger bonds

Size of Metal Atoms

• Smaller atoms: Stronger bonding
• Closer packing: Better orbital overlap
• Nuclear attraction: Closer electrons more strongly held

Intermolecular Forces

Van der Waals Forces

1. London Dispersion Forces

• Present in: All molecules and atoms
• Cause: Temporary electron distribution asymmetry
• Strength: Weakest intermolecular force
• Size dependence: Stronger in larger molecules

2. Dipole-Dipole Forces

• Between: Polar molecules
• Mechanism: Partial charges attract
• Orientation: Positive end to negative end
• Strength: Stronger than London forces

Hydrogen Bonding

Special Case

• Requirements: H bonded to N, O, or F
• Attraction: H attracted to lone pair on another N, O, F
• Strength: Stronger than other dipole-dipole forces
• Examples: Water, ammonia, hydrogen fluoride

Importance

• Water properties: High boiling point, ice structure
• DNA: Holds complementary strands together
• Proteins: Determines secondary and tertiary structure
• Life: Essential for biological processes

Effects on Physical Properties

Boiling and Melting Points

• Stronger forces: Higher melting/boiling points
• Energy needed: To overcome intermolecular attraction
• Trend: Ionic > hydrogen bonding > dipole-dipole > London

Solubility

• "Like dissolves like": Similar intermolecular forces
• Polar in polar: Water dissolves ionic compounds
• Nonpolar in nonpolar: Oil dissolves in oil
• Mixing: Depends on relative force strengths

Comparing bond and intermolecular force strengths

Approximate bond energies (kJ/mol): Intramolecular forces (chemical bonds):

• Ionic bonds: 600-4000 kJ/mol
• Covalent bonds: 150-1000 kJ/mol
• Metallic bonds: 100-800 kJ/mol

Intermolecular forces:

• Hydrogen bonds: 10-50 kJ/mol
• Dipole-dipole: 1-10 kJ/mol
• London forces: 0.1-5 kJ/mol

Significance:

• Chemical bonds hold atoms together in molecules
• Intermolecular forces hold molecules together
• Chemical bonds much stronger than intermolecular forces
• Breaking chemical bonds = chemical reactions
• Overcoming intermolecular forces = phase changes

Bond Properties and Energetics

Bond Strength

Bond Energy

• Definition: Energy needed to break one mole of bonds
• Units: kJ/mol or kcal/mol
• Bond formation: Releases energy (exothermic)
• Bond breaking: Requires energy (endothermic)

Factors Affecting Bond Strength

• Bond order: Single < double < triple
• Atomic size: Smaller atoms form stronger bonds
• Electronegativity difference: Can affect bond strength
• Hybridization: s character increases bond strength

Bond Length

Trends

• Bond order: Single > double > triple (length)
• Atomic size: Larger atoms = longer bonds
• Relationship: Stronger bonds are usually shorter
• Measurement: X-ray crystallography, spectroscopy

Bond Polarity and Dipole Moments

Measuring Polarity

• Dipole moment (μ): Measure of charge separation
• Units: Debye (D)
• Direction: From positive to negative charge
• Zero dipole: Nonpolar bonds and symmetric molecules

Applications of Bonding Theory

Material Properties

Predicting Properties

• Ionic compounds: Hard, brittle, high melting points
• Covalent networks: Very hard (diamond), high melting points
• Molecular compounds: Soft, low melting points
• Metals: Malleable, ductile, conductive

Drug Design

Molecular Recognition

• Shape complementarity: Drug fits receptor site
• Hydrogen bonding: Specific interactions
• Hydrophobic effects: Nonpolar regions interact
• Electrostatic interactions: Charged groups attract

Catalysis

Bonding in Catalysts

• Active sites: Specific bonding arrangements
• Substrate binding: Temporary bonds with reactants
• Transition states: Stabilized by catalyst bonding
• Product release: Weaker bonds allow product departure

Advanced Bonding Concepts

Resonance

Concept

• Multiple structures: More than one valid Lewis structure
• Reality: Actual structure is average of resonance forms
• Delocalization: Electrons spread over multiple atoms
• Stability: Resonance increases stability

Examples

• Benzene: Six resonance structures
• Carbonate ion: Three equivalent structures
• Ozone: Two resonance forms
• Nitrate ion: Three equivalent structures

Hybridization

Concept

• Orbital mixing: Atomic orbitals combine
• Hybrid orbitals: New orbitals with mixed character
• Shape prediction: Explains molecular geometry
• Bond formation: Better overlap with hybrid orbitals

Types

• sp³: Tetrahedral (CH₄)
• sp²: Trigonal planar (BF₃)
• sp: Linear (BeCl₂)
• sp³d, sp³d²: Expanded octets

Key Takeaways

• Atoms form bonds to achieve lower energy and stable electron configurations
• Electronegativity difference determines bond type: ionic, covalent, or metallic
• Ionic bonds involve electron transfer between metals and non-metals
• Covalent bonds involve electron sharing between non-metals
• Metallic bonding involves delocalized electrons in an "electron sea"
• VSEPR theory predicts molecular geometry based on electron pair repulsion
• Molecular polarity depends on bond polarity and molecular geometry
• Intermolecular forces determine physical properties like boiling point
• Bond strength increases with bond order and decreases with atomic size
• Understanding bonding helps predict and explain material properties