JjangJi's All About AP Chemistry Concept Lecture

Lesson 2: Chemical Bonding (1)

Types of Bonding vs. Types of Substances

Metallic Bond: Formed by the delocalization of electrons in a metal lattice. Electrons are free to move, creating a "sea of electrons" that accounts for properties like conductivity and malleability.

Ionic Bond: Formed by electrostatic attraction between oppositely charged ions. Typically occurs between metals and nonmetals, resulting in the formation of ionic compounds with high melting points.

Covalent Bond: Formed by the sharing of electrons between atoms. Covalent bonding leads to the formation of molecules and can result in different types of substances like molecular compounds and network solids.

Where:

Bond Breaking and Bond Forming

Bond Breaking: Endothermic process (requires energy input).
Bond Forming: Exothermic process (releases energy).

Understanding the energy changes involved in chemical reactions is crucial for predicting reaction spontaneity and calculating enthalpy changes.

Types of Crystals

Metallic Crystals: Consist of metal cations surrounded by a sea of delocalized valence electrons. These are good conductors of electricity and heat.
Ionic Crystals: Composed of alternating cations and anions held together by strong electrostatic forces. They are hard, brittle, and have high melting points.
Covalent Crystals (Network Solids): Atoms connected by covalent bonds in a continuous network. Examples include diamond (C), silicon (Si), and silicon dioxide (SiO₂), which are hard and have very high melting points.
Molecular Crystals: Consist of molecules held together by intermolecular forces like London dispersion forces, dipole-dipole interactions, or hydrogen bonds. They generally have low melting points.

Properties of Different Types of Solids

Type of Solid	Properties	Examples
Metallic Solids	Good conductors of heat and electricity Malleable and ductile Variable melting points	Iron (Fe), Copper (Cu), Gold (Au)
Ionic Solids	High melting and boiling points Brittle and hard Conduct electricity when molten or dissolved	Sodium Chloride (NaCl), Magnesium Oxide (MgO)
Network Covalent Solids	Very high melting points Hard and rigid Do not conduct electricity (except graphite)	Diamond (C), Silicon (Si), Silicon Dioxide (SiO₂)
Molecular Solids	Low melting and boiling points Soft and brittle Do not conduct electricity	Ice (H₂O), Dry Ice (CO₂)

Metallic Bond

Free Electrons: Delocalized electrons (valence electrons) that are free to move throughout the metal lattice.

Bond Strength: Can be compared using Coulomb’s Law:

$$ F = k \dfrac{q_1 q_2}{r} $$

Examples:

Li < Mg (Lithium has a weaker metallic bond than magnesium due to fewer delocalized electrons and a larger atomic radius.)
Li > Na (Lithium has a stronger metallic bond than sodium because lithium atoms are smaller, resulting in stronger attractions between nuclei and delocalized electrons.)

Alloys:

Interstitial Alloy: Formed when small atoms fill the interstitial spaces (holes) in a metal lattice (e.g., steel, where carbon atoms fit into iron lattice).
Substitutional Alloy: Formed when atoms of similar radii substitute for metal atoms in the lattice (e.g., brass, where zinc atoms replace some copper atoms).

Ionic Bond

Bond Strength: Compared using Coulomb’s Law:

$$ F = k \dfrac{q_1 q_2}{r} $$

Examples:

LiF > NaCl (Lithium fluoride has stronger ionic bonds than sodium chloride due to smaller ionic radii and higher charge density.)
LiF < MgO (Magnesium oxide has stronger ionic bonds because of higher charges on the ions, Mg²⁺ and O^2-, compared to Li⁺ and F^-.)

Lattice Energy:

The energy required to separate one mole of an ionic solid into its gaseous ions:

$$ XY(s) \rightarrow X^+(g) + Y^-(g) $$

Lattice energy is proportional to bond strength and can be predicted using Coulomb's Law.

Example: The lattice structure of NaCl consists of alternating Na⁺ and Cl^- ions in a cubic arrangement.

Covalent Bond

Understanding covalent bond formation in terms of potential energy.

Bond Polarity vs. Molecular Polarity:

Polar Covalent Bond: Unequal sharing of electrons between atoms with different electronegativities, resulting in a dipole moment.
Nonpolar Covalent Bond: Equal sharing of electrons between atoms with similar electronegativities.

Key Terms:

Polar Molecules: Molecules with an asymmetric arrangement of polar bonds, resulting in a net dipole moment.
Nonpolar Molecules: Molecules with symmetrical arrangements or only nonpolar bonds, resulting in no net dipole moment.

Potential Energy vs. Distance in Bond Formation

Graph of H₂ Bond Formation:

Bond Energy:

The energy required to dissociate the bond.
The net energy for stabilizing the bond.

Bond Length:

The distance between two nuclei at the minimum potential energy.
The distance where attractive and repulsive forces are balanced.

Potential Energy Graph:

At the lowest point in the graph, the bond is at its most stable energy state (approximately 0.74 Å for the H₂ bond).

Concepts:

Bond energy is directly related to bond strength.
As atoms approach each other, repulsive forces increase until they reach an optimal distance where potential energy is minimized.

Molecular Geometry

VSEPR Theory:

Valence Shell Electron Pair Repulsion theory states that electron pairs arrange themselves to minimize repulsion, determining the molecular geometry.

Hybridization: Involves the mixing of atomic orbitals to form new hybrid orbitals:

sp hybridization: Linear geometry.
sp² hybridization: Trigonal planar geometry.
sp³ hybridization: Tetrahedral geometry.

Order of Increasing Repulsion:

Bonding pair – bonding pair < lone pair – bonding pair < lone pair – lone pair.

Examples of Molecular Geometry:

Linear: 180° bond angle (e.g., CO₂).
Trigonal Planar: 120° bond angle (e.g., BF₃).
Tetrahedral: 109.5° bond angle (e.g., CH₄).
Trigonal Pyramidal: Less than 109.5° bond angle due to lone pair (e.g., NH₃).
Bent: Less than 120° or 109.5° bond angle depending on hybridization (e.g., H₂O).

More Complex Geometries:

Trigonal Bipyramidal: 90°, 120° bond angles (e.g., PCl₅).
Octahedral: 90° bond angles (e.g., SF₆).
Square Planar: 90° bond angles, typically with d⁸ metal complexes (e.g., XeF₄).