Unit 2: Comprehensive Summary

Elements are pure substances consisting of only one type of atom, represented by unique chemical symbols (e.g., H, O, C). Compounds are pure substances formed when two or more different elements chemically combine in fixed ratios (e.g., $\mathrm{H_2O}$, $\mathrm{CO_2}$). Mixtures are combinations of two or more substances that retain their individual properties and can be separated by physical methods.

The Laws of Chemical Combination govern compound formation:

• Law of Conservation of Mass: Mass is neither created nor destroyed in chemical reactions.
• Law of Definite Proportions: Compounds always contain the same elements in the same mass ratio.
• Law of Multiple Proportions: When elements form multiple compounds, the mass ratios form simple whole number ratios.

Separation techniques include filtration (insoluble solids from liquids), crystallization (soluble solids from solutions), distillation (liquids with different boiling points), and chromatography (substances with different affinities for stationary phases).

Chemical bonds form when atoms interact to achieve stable electronic configurations, typically by achieving a full outer electron shell (octet rule for most atoms, duet for hydrogen).

Ionic Bonding

Forms between metals and non-metals through complete electron transfer. Metals lose electrons to form cations (positive ions), while non-metals gain electrons to form anions (negative ions). The resulting electrostatic attraction creates ionic bonds. Ionic compounds form giant ionic lattices with high melting/boiling points, conduct electricity when molten or in solution, and are often soluble in polar solvents.

Covalent Bonding

Forms between non-metals through electron sharing. Atoms share pairs of electrons to achieve stable configurations. Single bonds involve one shared pair (e.g., H-H), double bonds involve two pairs (e.g., O=O), and triple bonds involve three pairs (e.g., N≡N). Polar covalent bonds occur when electrons are unequally shared due to electronegativity differences, while non-polar covalent bonds have equal sharing.

Metallic Bonding

Occurs in metals where valence electrons are delocalized, forming a "sea of electrons" around positive metal ions. This explains metals' properties: electrical conductivity (mobile electrons), thermal conductivity, malleability and ductility (layers can slide), and metallic luster.

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular geometry based on electron pair repulsion. Electron pairs (bonding and lone pairs) arrange themselves to minimize repulsion, with lone pairs exerting stronger repulsion than bonding pairs.

Common molecular shapes include:

• Linear: 2 electron domains, 180° bond angle (e.g., $\mathrm{CO_2}$, $\mathrm{BeCl_2}$)
• Trigonal Planar: 3 electron domains, 120° bond angles (e.g., $\mathrm{BF_3}$)
• Tetrahedral: 4 electron domains, 109.5° bond angles (e.g., $\mathrm{CH_4}$)
• Trigonal Pyramidal: 4 domains (3 bonding, 1 lone pair), ~107° (e.g., $\mathrm{NH_3}$)
• Bent/Non-linear: 4 domains (2 bonding, 2 lone pairs), ~104.5° (e.g., $\mathrm{H_2O}$)

Intermolecular forces (IMFs) are weak attractive forces between molecules that determine physical properties like boiling points, melting points, and solubility.

Types of IMFs (weakest to strongest):

• London Dispersion Forces (LDFs): Present in all molecules, arise from temporary dipoles. Strength increases with molecular size and electron count.
• Dipole-Dipole Forces: Between polar molecules with permanent dipoles.
• Hydrogen Bonding: Special dipole-dipole interaction when H is bonded to N, O, or F and attracted to lone pairs on another N, O, or F atom.

Solubility follows the "like dissolves like" principle:

• Polar solutes dissolve in polar solvents (strong dipole-dipole interactions or hydrogen bonding)
• Nonpolar solutes dissolve in nonpolar solvents (similar London dispersion forces)
• Polar and nonpolar substances generally don't mix due to energy costs

The macroscopic properties of substances depend on their microscopic structure and bonding type. Four main structural types exist:

1. Simple Molecular Structures

Bonding: Strong covalent bonds within molecules, weak IMFs between molecules.
Properties: Low melting/boiling points, often gases/liquids at room temperature, no electrical conductivity, generally soft.
Examples: $\mathrm{H_2O}$, $\mathrm{CO_2}$, $\mathrm{CH_4}$, $\mathrm{I_2}$

2. Giant Covalent Structures

Bonding: Extensive network of strong covalent bonds throughout the structure.
Properties: Very high melting/boiling points, hard and brittle, mostly non-conductive (except graphite), insoluble.
Examples: Diamond, graphite, silicon dioxide ($\mathrm{SiO_2}$)

3. Giant Ionic Structures

Bonding: Strong electrostatic forces between oppositely charged ions in a lattice.
Properties: High melting/boiling points, hard and brittle, conduct when molten/in solution, soluble in polar solvents.
Examples: $\mathrm{NaCl}$, $\mathrm{MgO}$

4. Giant Metallic Structures

Bonding: Electrostatic attraction between metal cations and delocalized electron sea.
Properties: High melting/boiling points, excellent electrical/thermal conductivity, malleable and ductile.
Examples: Copper (Cu), iron (Fe), aluminum (Al)

The Kinetic Theory of Matter explains the behavior of particles in different states:

• All matter consists of tiny particles in constant, random motion
• Average kinetic energy is proportional to absolute temperature
• Particles in gases move freely with minimal intermolecular forces
• Particles in liquids have some freedom but experience significant intermolecular forces
• Particles in solids vibrate around fixed positions with strong intermolecular forces

States of Matter:

• Solids: Fixed shape and volume, particles vibrate around fixed positions
• Liquids: Fixed volume but take container shape, particles move freely but stay close
• Gases: Fill container volume, particles move freely with high kinetic energy

State Changes occur when energy is added or removed. Impurities generally lower melting points and raise boiling points by disrupting the regular arrangement of particles.