Chemical Bonding

The Nature of Atomic Bonding and Energy Changes

• Motivation for Bonding: Atoms form chemical bonds to achieve the same electron configuration as noble gases. Noble gases have full valence electron orbitals, which makes them stable and less reactive compared to other elements.

• The Mechanism of Bond Formation: When two hydrogen atoms approach one another, three distinct forces come into play:     - Repulsive Forces: These occur between the electrons of both atoms and between the nuclei of both atoms.     - Attractive Forces: These occur between the nucleus of one atom and the electrons of the adjacent atom.     - Energy Balance: As these forces interact, the energy of the system changes. A bond is formally established when the attractive forces are equal to the repulsive forces at the lowest possible potential energy for the system.

• Bond Length: This is defined as the specific distance between the nuclei of two adjacent atoms at the point when they bond.

• Bond Energy: This refers to the quantity of energy that must be introduced into the system to break a bond that has already formed.

• Bond Strength: The strength of a bond is influenced by several factors:     - Bond Length: Generally, shorter bond lengths correlate with stronger bonds.     - Atomic Size: The size of the bonded atoms impacts strength.     - Bond Order: The number of bonds between atoms; a higher number of bonds (e.g., triple vs. single) generally lead to a stronger bond.

Representation of Bonds and Lewis Notation

• Lewis Notation: This system uses dots or crosses to represent valence electrons. These electrons are categorized into two types:     - Lone Pairs: Valence electrons not involved in bonding.     - Bonding Pairs: Electrons shared between atoms.

• Couper Structures: These representations use a single solid line to represent a bonded pair of electrons.

• Rules for Constructing Lewis Diagrams:     - Central Atom Selection: Choose the atom with the lowest electronegativity as the central atom.     - Valence Electron Calculation: Determine the total number of valence electrons for all atoms involved.     - Electron Pairs: Divide the total number of valence electrons by $2$ to find the number of electron pairs.     - Assignment: Assign these pairs to form bonds or as terminal lone pairs.     - Bond Optimization: If necessary, convert lone pairs to double or triple bonds to satisfy valency requirements.

Chemical Bond Classifications and Electronegativity

• Intramolecular Bond: A chemical bond occurring between atoms within a single molecule.

• Covalent Bond: A bond characterized by the sharing of at least one pair of electrons between two atoms.

• Electronegativity: A measure of the specific tendency of an atom to attract a bonding pair of electrons toward itself.

• Polarity of Covalent Bonds:     - Non-polar Covalent: Electrons are shared equally due to a zero difference in electronegativity.     - Polar Covalent: Electrons are shared unequally due to a difference in electronegativity between the atoms.

• Electronegativity Difference ($\Delta En$) and Bond Types:     - Non-polar Covalent: $\Delta En = 0$     - Weak Polar Covalent: $\Delta En$ ranges from $0.1$ to $1$     - Strong Polar Covalent: $\Delta En$ ranges from $1.1$ to $2$     - Ionic Bond: \Delta En > 2.0

• Ionic Bonding: This involves the complete transfer of electrons from one atom to another to form cations (positive ions) and anions (negative ions). These ions are held together by electrostatic attraction in a crystal lattice. This typically occurs between metals and non-metals.

• Metallic Bonding: Defined as the attraction between positive kernels and a "sea" of delocalised electrons.

• Dative Covalent Bond: This is defined as the overlap of a lone pair of electrons from one atom with an empty orbital of another atom.

Valence Shell Electron-Pair Repulsion (VSEPR) Theory

• VSEPR Definition: A model used to determine molecular shape based on the repulsive behavior of electron pairs, which seek to maximize the angles between them to minimize repulsion.

• Electron Pair Distribution: While electron pairs distribute evenly, lone pairs occupy more space than bonded pairs.

• Hierarchy of Repulsion: The strength of repulsion follows the specific order: \text{lone pair-lone pair} > \text{lone pair-bonded pair} > \text{bonded pair-bonded pair}.

• Common Molecular Geometries:     - Linear: Represented as $AX$ or $AX_2$     - Bent or Angular: Represented as $AX_2E_2$     - Trigonal Planar: Represented as $AX_3$     - Trigonal Pyramidal: Represented as $AX_3E$     - Tetrahedral: Represented as $AX_4$     - Trigonal Bipyramidal: Represented as $AX_5$     - Octahedral: Represented as $AX_6$

• Molecular Polarity:     - Non-polar Molecule: Electrons are equally dispersed across a molecule that is symmetrical.     - Polar Molecule (Dipole): Electrons are distributed unevenly, creating distinct $\delta+$ (partial positive) and $\delta-$ (partial negative) ends.

Intermolecular Forces (IMF)

• IMF Definition: Forces of attraction that exist between molecules or between atoms of noble gases. These are significantly weaker than intramolecular bonds and are the primary determinants of physical properties like melting and boiling points.

• Van der Waals Forces:     - London Forces (Induced Dipole Forces): These occur between all atoms and molecules. Because electrons are in constant motion, momentary uneven distributions of charge occur, which induce dipoles in neighboring particles.     - Dipole-Dipole Forces: Attractions between the permanent dipoles of polar molecules.

• Hydrogen Bond Forces: A specific and strong case of dipole-dipole forces. They occur when a hydrogen atom is covalently bonded to a highly electronegative atom (Nitrogen, Oxygen, or Fluorine) and is attracted to the lone pair of a neighboring molecule. Factors contributing to its strength include:     - Large electronegativity differences.     - High electron density in $N$, $O$, or $F$, resulting in powerful dipoles.     - Small atomic radii, which allows molecules to approach each other very closely.

• Factors Influencing IMF Strength:     - Molecular Size: Larger molecules have more electrons, making it easier to establish temporary dipoles.     - Molecular Shape: Straight-chain molecules provide a larger contact surface area than branched chains, leading to stronger attractions.     - Polarity: A more polar molecule results in stronger overall IMF.     - Bonding Sites: A higher number of hydrogen bonding sites increases the total force strength.

Physical Properties and Substance Structures

• Melting and Boiling Points: These increase as the strength of the intermolecular forces increases.

• Solubility: Governed by the principle of "like dissolving like." Solubility occurs when the solute and solvent possess similar intermolecular force strengths.

• Simple Molecules: Examples include $O_2$, $H_2$, and $H_2O$. These consist of a fixed, specific number of atoms held together by weak intermolecular forces.

• Giant or Macromolecules: These consist of an unknown or unlimited number of atoms held together by strong interatomic forces throughout the solid. Types include ionic crystals, metallic crystals, and covalent network structures.

• Specific Covalent Network Structures:     - Diamond: A giant covalent structure where each carbon atom forms four single bonds in a bonded tetrahedral arrangement. It is extremely hard, has a melting point of approximately $4000 \, ^\circ\text{C}$, does not conduct electricity, and is insoluble.     - Graphite: Features a layered structure where carbon atoms form only three bonds. This leaves delocalised electrons that allow graphite to conduct electricity. Van der Waals dispersion forces exist between the sheets, making it soft, slippery, and less dense than diamond.     - Silicon Dioxide ($SiO_2$): A giant covalent structure similar to diamond. Each silicon atom is bridged by oxygen atoms. It is hard, has a high melting point of around $1700 \, ^\circ\text{C}$, and is both insoluble and non-conductive.

• Giant Ionic Structures (Sodium Chloride, $NaCl$):     - Lattice Structure: Consists of alternating ions held by strong electrostatic attractions.     - Property Determinants: Melting and boiling points depend on the magnitude of the charges and the size of the ions.     - Electrical Conductivity: Solid $NaCl$ does NOT conduct electricity because the ions are fixed in place. However, it undergoes electrolysis in molten or aqueous states because the ions are then free to move and discharge.     - Brittleness: Ionic crystals are brittle because mechanical stress can shift layers of ions. This movement brings ions of like charges side-by-side, causing the crystal to repel itself and shatter.