Inorganic chemistry Study Guide

📖 Core Concepts Inorganic chemistry – study of non‑carbon‑centric compounds, including organometallics and solid‑state materials. Bonding types – ionic (cations + anions), covalent (shared pairs), polar covalent (partial ionic character). Lewis acid/base – electron‑pair acceptor/donor; HSAB theory refines this by hard/soft character (size, polarizability). Coordination compounds – central metal bound to ligands (donor atoms with lone pairs). Geometry dictated by metal‑ligand electron counting. Main‑group vs transition‑metal chemistry – main‑group: often covalent, can be hypervalent; transition metals: d‑orbitals drive color, magnetism, substitution mechanisms. Electron‑counting – tally valence electrons on the metal to predict structure/reactivity. Crystal‑field theory – splitting of d‑orbitals by ligands explains magnetism and colors of complexes. Born–Haber cycle – thermodynamic pathway to calculate lattice energy of ionic solids. --- 📌 Must Remember HSAB rule: Hard acids prefer hard bases (e.g., \( \mathrm{Na^+}\) ↔ \( \mathrm{F^-}\)); soft acids prefer soft bases (e.g., \( \mathrm{Ag^+}\) ↔ \( \mathrm{I^-}\)). VSEPR shapes: NH₃ → trigonal pyramidal; SO₂ → bent (polar covalent). Crystal‑field example: \([ \mathrm{Fe^{III}(CN)6}]^{3-}\) has one unpaired electron → weak magnetism. Electron‑deficient boranes (e.g., \(\mathrm{B2H6}\)) behave like carbocations → associative reactivity. Water‑exchange rate trend: \([ \mathrm{M(H2O)6}]^{n+}\) rates vary by \(10^{20}\) across the periodic table. Redox steps: oxidative addition (increase oxidation state) ↔ reductive elimination (decrease oxidation state). Key industrial compound: Ammonium nitrate → fertilizer from Haber‑process ammonia. --- 🔄 Key Processes Born–Haber Cycle Sublimation → ionization → bond dissociation → electron affinity → lattice energy → formation enthalpy. Ligand Substitution (Transition Metals) Associative (A): incoming ligand forms a 7‑coordinate intermediate → ligand loss. Dissociative (D): ligand leaves first → 5‑coordinate intermediate → incoming ligand binds. Oxidative Addition / Reductive Elimination Oxidative addition: \(\mathrm{M^n} + \mathrm{X–Y} \rightarrow \mathrm{M^{n+2}(X)(Y)}\). Reductive elimination: reverse, forms a new X–Y bond and reduces metal oxidation state. Electron‑Counting (18‑electron rule) Count: metal valence electrons + electrons donated by each ligand (2 per L‑type, 1 per X‑type). Catalysis Types Homogeneous: organometallic complex in solution (e.g., hydrogenation). Heterogeneous: reaction on solid surface (e.g., catalytic converters). --- 🔍 Key Comparisons Ionic vs Covalent vs Polar Covalent Ionic: full transfer, large lattice energy (e.g., NaCl). Covalent: equal sharing, discrete molecules (e.g., SO₂). Polar covalent: unequal sharing, many oxides/carbonates/halides. Hard vs Soft Acids/Bases Hard: small, low polarizability (e.g., \( \mathrm{Li^+}, \mathrm{O^{2-}}\)). Soft: large, highly polarizable (e.g., \( \mathrm{Pt^{2+}}, \mathrm{S^{2-}}\)). Associative vs Dissociative Substitution Associative: rate ↑ with ligand concentration; common for 4‑coordinate d⁸ metals. Dissociative: rate independent of entering ligand; common for 6‑coordinate high‑spin d³–d⁵ metals. Homogeneous vs Heterogeneous Catalysis Homogeneous: molecular catalyst, easy to study mechanism. Heterogeneous: solid catalyst, surface phenomena dominate. --- ⚠️ Common Misunderstandings “All metal‑carbon bonds are organometallic.” – Only bonds where carbon is directly attached to a metal (M–C–H fragment) count; metal‑bound carbonyls are a subclass. “All transition‑metal complexes are colored.” – Color requires d‑d or charge‑transfer transitions; d⁰ or d¹⁰ complexes can be colorless. “Hypervalent = >8 e⁻ only for main‑group.” – Hypervalency is a main‑group concept; transition metals achieve >8 e⁻ via d‑orbitals, not “hypervalent” terminology. “Lewis acids are always metals.” – Any species that can accept an electron pair (e.g., \(\mathrm{BF3}\)) qualifies, even non‑metals. --- 🧠 Mental Models / Intuition HSAB as “fit‑and‑stick”: Hard ↔ hard = tight, strong, mostly ionic; Soft ↔ soft = looser, more covalent, polarizable. Crystal‑field splitting as “energy ladder”: Strong‑field ligands (CN⁻) push electrons low → low spin; weak‑field (H₂O) keep them high → high spin. Electron‑counting as “budget”: Metal starts with its valence; each ligand adds a fixed “spending” amount. Aim for 18 e⁻ (stable “full‑budget”) for many low‑oxidation‑state complexes. --- 🚩 Exceptions & Edge Cases d⁰ and d¹⁰ complexes (e.g., \(\mathrm{TiCl4}, \mathrm{Zn(CN)2}\)) do not follow crystal‑field predictions for magnetism or color. Lanthanides/actinides often ignore crystal‑field effects; f‑orbitals are shielded. Electron‑deficient boranes violate octet rule but are stabilized by 3‑center‑2‑electron bonds. HSAB soft–soft interactions can be kinetically sluggish despite thermodynamic favorability (e.g., \(\mathrm{Ag^+}\) with \(\mathrm{I^-}\)). --- 📍 When to Use Which Choose HSAB analysis when predicting acid–base selectivity in inorganic syntheses. Apply crystal‑field theory for transition‑metal complexes with visible colors or magnetic data. Use electron‑counting to assess stability of organometallic catalysts (aim for 18 e⁻). Employ Born–Haber cycle for lattice‑energy calculations of ionic solids (e.g., NaCl, \(\mathrm{YBa2Cu3O7}\)). Select VSEPR for main‑group molecular geometry; switch to crystal‑field for transition‑metal geometry. --- 👀 Patterns to Recognize Color ↔ strong‑field ligand → low‑spin, fewer unpaired electrons. Rapid water exchange in \([ \mathrm{M(H2O)6}]^{n+}\) indicates a soft, high‑spin metal (e.g., Cu²⁺). Hypervalent main‑group species often feature electron‑deficient bonding (e.g., \(\mathrm{B2H6}\)). Ligand‑induced acidity: coordinated amines become more acidic than free amines (e.g., \([ \mathrm{Co(NH3)6}]^{3+}\)). --- 🗂️ Exam Traps Mistaking a Lewis acid for a Brønsted acid – Lewis acids accept electrons; Brønsted acids donate protons. Assuming all 18‑electron complexes are inert – some 18‑e⁻ species are highly reactive (e.g., \(\mathrm{Fe(CO)5}\) under photolysis). Confusing “ionic” with “polar covalent” – many metal oxides have significant covalent character despite being labeled “ionic”. Over‑applying VSEPR to transition‑metal complexes – geometry is governed by d‑orbital splitting, not just electron pair repulsion. Neglecting soft‑soft kinetic barriers – a thermodynamically favored soft‑soft interaction may be slow, leading to “no reaction” in the lab. ---