️ Unit 2: Molecular and lonic Compound Structure and Properties

️ Unit 2: Molecular and lonic Compound Structure and Properties After you're done reading this, generate practice quiz questions or a podcast to review everything! These are Turbo's most popular features. Brief Overview: This unit focuses on how atoms bond together to form molecular and ionic compounds. We'll explore the differences between ionic and covalent bonding, and discover how the structure of these compounds determines their unique properties. ⚡️ Bonds Overview Atoms seek a lower‑energy, more stable configuration by transferring or sharing electrons. This process is called bonding. Atoms are most stable with an octet (8 valence electrons). Two main bonding types: ionic (electron transfer) and covalent (electron sharing). 🔗 Ionic Bonds An ionic solid consists of a lattice of oppositely charged ions held together by electrostatic (Coulombic) forces. Formation: A metal cation donates one or more electrons to a non‑metal anion. Example: Na → Na⁺ + e⁻, Cl + e⁻ → Cl⁻, forming NaCl. Lattice Energy: Increases with (1) higher ionic charges (Coulomb’s Law) and (2) smaller ionic radii. MgO (+2 / –2) melts at a higher temperature than NaCl (+1 / –1). LiF (small ions) melts higher than KBr (large ions). Physical properties: Typically solid at room temperature, with high melting/boiling points. Poor electrical conductors in the solid state (electrons are localized). Ionic liquids conduct because ions are mobile in the liquid phase. ️ Metallic Bonds & Alloys Metals are described by the “sea of electrons” model: a lattice of positive metal cores immersed in a delocalized electron cloud. Conductivity: Mobile electrons allow excellent electrical conduction. Malleability & Ductility: Delocalized electrons let metal ions slide past one another without breaking the bond. Alloy formation: Interstitial alloys – small atoms (e.g., C) occupy spaces between larger metal atoms. Substitutional alloys – atoms of similar size replace some host metal atoms. The left diagram shows carbon atoms in the interstices of iron (steel); the right diagram shows zinc atoms substituting for copper in brass. 🧪 Molecular Covalent Bonding Covalent bonds involve the sharing of electron pairs between non‑metal atoms. Each shared pair counts toward the octet of both atoms. Example: Two fluorine atoms each need one electron → form a single covalent bond. Molecules range from diatomic (e.g., F₂) to macromolecules (e.g., glucose, C₆H₁₂O₆). Bond types: Sigma (σ) bond – the first bond formed, located along the internuclear axis. Pi (π) bond – formed from side‑on overlap; appears in double and triple bonds. 📏 Internuclear Distance The bond length is the distance at which the potential energy of two atoms is minimized—where attractive and repulsive forces balance. Too close → strong nuclear repulsion → high energy. Too far → negligible attraction → energy approaches zero. Bond type Designation Bond order Length Energy Single 1σ 1 Longest Least Double 1σ+1π 2 Shorter Intermediate Triple 1σ+2π 3 Shortest Greatest The dashed line marks zero potential energy; the true minimum (most negative) occurs at 74pm. ️ Network Covalent Solids A network solid is a continuous lattice of covalent bonds—effectively one giant molecule. Properties: Very hard, high melting/boiling points, poor electrical conductors (electrons localized). Common examples: Diamond, graphite, silicon dioxide (quartz) – all contain C or Si (four valence electrons) enabling extensive 3‑D bonding. 🔌 Conductivity Overview Phase Ionic Molecular Covalent Network Covalent Metallic Solid No No No Yes Aqueous Yes No N/A Yes Pure water does not conduct; tap water conducts because dissolved ions provide charge carriers. Ionic conductivity depends on concentration and the number of ions produced per formula unit (e.g., 1M CaCl₂ → 3ions, conducts better than 1M NaCl → 2ions). ✏ Lewis Dot Structures Steps to draw a Lewis structure 1. Count valence electrons of all atoms. 2. Adjust for ionic charge (+→subtract, –→add). 3. Sketch the skeletal framework; place a single bond (2e⁻) between each connected pair. 4. Complete octets of outer atoms first. 5. Place remaining electrons on the central atom. 6. If the central atom lacks an octet, convert lone pairs on surrounding atoms into additional bonds (π bonds). Example: Carbonate ion (CO₃ ²⁻) Total valence electrons: C(4)+3×O(6)=22; add 2 for the –2 charge → 24 e⁻. Central C, single bonds to three O → 6 e⁻ used. Distribute remaining 18 e⁻ to complete O octets → each O gets 6 e⁻ (3 lone pairs). Carbon still has only 6 e⁻; convert one O lone pair into a double bond. Liquid Yes No N/A Yes Gas No No No No 🔄 Resonance & Bond Order Resonance: When multiple Lewis structures can represent the same ion/molecule, the true structure is a hybrid. In CO₃ ²⁻, the double bond can be placed on any of the three oxygens, giving three resonance forms. Bond order calculation: Bond order = Sum of bond orders in all resonance forms / Number of resonance forms For each C–O bond in CO₃ ²⁻: (1+2+1) /3=1.33 → all three bonds are equivalent, intermediate between single and double. ⚛ Incomplete and Expanded Octets Incomplete octet (stable with <8 e⁻): Hydrogen (needs 2 e⁻) Helium (2 e⁻, never bonds) Boron (stable with 6 e⁻, e.g., BF₃) Expanded octet (more than 8 e⁻) is possible for elements in period3 or higher (have d‑orbitals). Common for Si, P, S, Cl, and even noble gases (e.g., XeF₄). Maximum of 12 valence electrons (6 bonds). 📐 Formal Charge Formal charge (FC) = (valence electrons)−[(non‑bonding electrons)+½(bonding electrons)] The most plausible Lewis structure minimizes the magnitude and number of formal charges. For neutral molecules, the sum of FCs = 0; for ions, sum = overall charge. Molecule Formal charges (example) CO₂ C=0, each O=0 Image of various polyatomic structures (PCl₅, SF₄, XeF₄) helps visualize how formal charges distribute across central atoms. CO₃ ²⁻ Two O with –1, one O with 0, C=+2 (overall –2) → less favored than the resonance hybrid with all bonds equivalent. 🧭 Molecular Geometry (VSEPR) Electron pairs (bonding+lone) repel; the molecule adopts the geometry that maximizes separation of these pairs. Hybridization & basic shapes: Electron pairs on central atom Hybridization Basic geometry Typical bond angle 2 (0 lone) sp Linear 180° 3 (0 lone) sp² Trigonal planar 120° 3 (1 lone) sp² Bent <120° 4 (0 lone) sp³ Tetrahedral 109.5° 4 (1 lone) sp³ Trigonal pyramidal <109.5° 4 (2 lone) sp³ Bent (V‑shape) <109.5° 5 (0 lone) sp³d Trigonal bipyramidal 120°/90° 5 (1 lone) sp³d See‑saw <120° 5 (2 lone) sp³d T‑shaped <90° 5 (3 lone) sp³d Linear 180° 6 (0 lone) sp³d² Octahedral 90° 6 (1 lone) sp³d² Square pyramidal <90° 6 (2 lone) sp³d² Square planar 90° Number of lone pairs Geometry Examples 0 B PCI B AB PF B B trigonal bipyramidal 2 3 B SF B IF A B B folded square, seesaw, distorted tetrahedron B CIF, ICI, B-A B T-shaped B :-A B linear XeF, Key points: Double/triple bonds count as a single electron‑pair region for geometry, but exert slightly greater repulsion. Lone pairs occupy more space than bonding pairs → bond angles are reduced relative to the ideal geometry. 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