Chemical bond Study Guide

📖 Core Concepts Chemical bond – the attraction that holds atoms/ions together into molecules, crystals, etc. Primary (strong) bonds – covalent, ionic, metallic; involve sharing or full transfer of electrons. Secondary (weak) bonds – dipole‑dipole, London dispersion, hydrogen bonds; act between separate molecules. Electronegativity (χ) – tendency of an atom to pull shared electrons toward itself; larger χ → stronger pull. Bond polarity continuum – bonds range from non‑polar covalent (Δχ ≈ 0‑0.3) → polar covalent (Δχ ≈ 0.3‑1.7) → ionic (Δχ > 1.7). Sigma (σ) vs. Pi (π) bonds – σ: head‑on overlap, cylindrical symmetry; π: side‑on overlap, electron density above/below bond axis. Valence‑bond (VB) vs. Molecular‑orbital (MO) models – VB: localized electron pairs, hybridisation; MO: delocalised electrons over the whole molecule. --- 📌 Must Remember Δχ ≤ 0.3 → non‑polar covalent. 0.3 < Δχ ≤ 1.7 → polar covalent. Δχ > 1.7 → ionic bond. Single bond = 1 σ bond; double bond = 1 σ + 1 π; triple bond = 1 σ + 2 π. Hydrogen bond donor: H attached to N, O, or F. Acceptor: lone pair on N, O, or F. London dispersion present in all molecules; strongest in large, polarizable atoms. Ionic crystals: high melting point, brittle, soluble in polar solvents. Metallic bond properties – luster, conductivity, ductility from delocalised electron “sea.” --- 🔄 Key Processes Predict bond type using electronegativity: Calculate Δχ between two atoms → apply Δχ thresholds above. Determine bond order & strength: Count σ + π bonds → more bonds = stronger & shorter. Identify hydrogen‑bonding capability: Look for H‑X (X = N, O, F) and a nearby lone pair on N, O, or F. Apply VSEPR (quick check): Use electron‑pair repulsion to predict molecular geometry → infer polarity direction. Choose bonding model: Use VB for localized descriptions (hybridisation, resonance). Use MO for delocalised systems (conjugation, magnetic properties). --- 🔍 Key Comparisons Ionic vs. Covalent Ionic: Δχ > 1.7, full electron transfer, non‑directional crystal lattice. Covalent: Δχ ≤ 1.7, electron sharing, directional bonds. Polar vs. Non‑polar Covalent Polar: Δχ 0.3‑1.7, unequal electron cloud → dipole moment. Non‑polar: Δχ ≤ 0.3, equal sharing → no permanent dipole. Sigma vs. Pi Bonds σ: head‑on, present in all single bonds, allows free rotation (unless π present). π: side‑on, restricts rotation, found in double/triple bonds. Weak Forces: Keesom vs. London Keesom: permanent dipole–permanent dipole (requires polar molecules). London: instantaneous dipole–induced dipole (present in all molecules). --- ⚠️ Common Misunderstandings “All ionic bonds are 100 % ionic.” – Real bonds lie on a continuum; many have mixed character. “Hydrogen bonds are covalent.” – They are strong intermolecular forces, not true covalent bonds. “Metallic bonds involve sharing of specific electron pairs.” – Electrons are delocalised across the lattice, not shared pairwise. “π bonds are stronger than σ bonds.” – σ bonds are generally stronger; π bonds are weaker and add to overall bond order. --- 🧠 Mental Models / Intuition Electronegativity difference as a “balance scale”: Small difference → balance (non‑polar); big difference → one side dominates (ionic). Bond‑type continuum as a colour gradient: Dark blue = non‑polar covalent, green = polar covalent, red = ionic. Sigma‑pi layering: Think of a road (σ) with an overpass (π) – the road carries the main traffic, the overpass adds extra capacity but is more fragile. Weak forces as “glue”: Keesom = strong velcro (needs matching poles); London = generic tape (sticks to anything). --- 🚩 Exceptions & Edge Cases Δχ ≈ 1.7 – borderline; bond may show both ionic lattice and covalent character (e.g., AlCl₃). Hydrogen bonding with C–H – generally negligible because C is not electronegative enough. Metallic bonding in alloys – different metals can alter electron‑sea density, affecting conductivity. Large, highly polarizable atoms (e.g., I₂) → strong London forces despite being non‑polar. --- 📍 When to Use Which Choose Δχ rule when quickly classifying bond type in exam‑style questions. Use VSEPR for predicting molecular shape and polarity of small molecules. Apply VB/hybridisation when asked about orbital composition (sp, sp², sp³). Switch to MO theory for conjugated systems, aromaticity, or magnetic property questions. Consider hydrogen‑bond criteria when evaluating boiling point trends or solubility of compounds containing N, O, or F. --- 👀 Patterns to Recognize Increasing bond order → decreasing bond length & increasing bond energy. Presence of N, O, or F attached to H → likely hydrogen bonding → higher boiling point. Molecules with only London forces → low melting/boiling points, non‑polar. Transition metal complexes → ligand field theory (MO) needed, not simple VSEPR. --- 🗂️ Exam Traps Δχ = 1.7 presented as “ionic” – answer may be “polar covalent” if the question emphasises partial ionic character. Hydrogen bond listed as “covalent” – remember it’s an intermolecular force. “Metallic bond strength > covalent bond strength” – false; covalent bonds are typically stronger per bond. Choosing London forces for polar molecules – polar molecules also have Keesom forces; the most significant force is often the dipole‑dipole (Keesom). Confusing resonance with delocalisation in MO theory – resonance is a VB concept; MO delocalisation is a separate model. ---