1: Basic Inorganic Concepts
- Page ID
- 428141
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)- 1.1: Introduction to Inorganic Chemistry
- Welcome to Inorganic chemistry; The image shown here was created by graphics designer, John Megahan (https://johnmegahan.com/) in consultation with Anne McNeil (https://mcneilgroup.chem.lsa.umich.edu/diversity/)
- 1.2: Atomic Structure
- 1.2.1: Historical Development of Atomic Theory
- 1.2.1.1: The Periodic Table
- 1.2.1.2: Discovery of Subatomic Particles and the Bohr Atom
- 1.2.2: The Schrödinger equation, particle in a box, and atomic wavefunctions
- 1.2.2.1: Particle in a Box
- 1.2.2.2: Quantum Numbers and Atomic Wave Functions
- 1.2.2.3: Aufbau Principle
- 1.2.2.4: Shielding
- 1.2.3: Periodic Properties of Atoms
- 1.2.3.1: Ionization energy
- 1.2.3.2: Electron Affinity
- 1.2.3.3: Covalent and Ionic Radii
- 1.2.4: Problems (do we want this here?)
- 1.3: Simple Bonding Theory
- 1.3.1: Lewis Electron-Dot Diagrams
- 1.3.1.1: Resonance
- 1.3.1.2: Breaking the octet rule with higher electron counts (hypervalent atoms)
- 1.3.1.3: Formal Charge
- 1.3.1.4: Lewis Fails to Predict Unusual Cases - Boron and Beryllium
- 1.3.2: Valence Shell Electron-Pair Repulsion
- 1.3.2.1: Lone Pair Repulsion
- 1.3.2.2: Multiple Bonds
- 1.3.2.3: Electronegativity and Atomic Size Effects
- 1.3.2.4: Ligand Close Packing
- 1.3.3: Molecular Polarity
- 1.3.4: Hydrogen Bonding
- 1.3.5: Problems (do we want this here?)
- 1.4: Symmetry and Group Theory
- 1.4.1: Symmetry Elements and Operations
- 1.4.2: Point Groups
- 1.4.2.1: Groups of Low and High Symmetry
- 1.4.2.2: Other Groups
- 1.4.3: Properties and Representations of Groups
- 1.4.3.1: Matrices
- 1.4.3.2: Representations of Point Groups
- 1.4.3.3: Character Tables
- 1.4.4: Examples and Applications of Symmetry
- 1.4.4.1: Chirality
- 1.4.4.2: Molecular Vibrations
- 1.4.P: Problems (under construction)
- 1.5: Molecular Orbitals
- Molecular Orbital (MO) Theory is a sophisticated bonding model. It is generally considered to be more powerful than Lewis and Valence Bond Theories for predicting molecular properties, however this power comes at the price of complexity. In its full development, MO Theory requires complex mathematics, though the ideas behind it are simple. Atomic orbitals (AOs) that are localized on individual atoms combine to make molecular orbitals (MOs) that are distributed over the molecule.
- 1.5.1: Formation of Molecular Orbitals from Atomic Orbitals
- 1.5.1.1: Molecular Orbitals from s Orbitals
- 1.5.1.2: Molecular Orbitals from p Orbitals
- 1.5.1.3: Molecular orbitals from d orbitals
- 1.5.1.4: Nonbonding Orbitals and Other Factors
- 1.5.2: Homonuclear Diatomic Molecules
- 1.5.2.1: Molecular Orbitals
- 1.5.2.2: Orbital Mixing
- 1.5.2.3: Diatomic Molecules of the First and Second Periods
- 1.5.2.4: Photoelectron Spectroscopy
- 1.5.3: Heteronuclear Diatomic Molecules
- 1.5.3.1: Orbital ionization energies
- 1.5.3.2: Polar bonds
- 1.5.3.3: Ionic Compounds and Molecular Orbitals
- 1.5.4: Larger (Polyatomic) Molecules
- 1.5.4.1: Bifluoride anion
- 1.5.4.2: Carbon Dioxide
- 1.5.4.3: H₂O
- 1.5.4.4: NH₃
- 1.5.4.5: CO₂ (Revisted with Projection Operators)
- 1.5.4.6: BF₃
- 1.5.P: Problems
- 1.6: Acid-Base and Donor-Acceptor Chemistry
- 1.6.1: Acid-Base Models as Organizing Concepts
- 1.6.2: Arrhenius Concept
- 1.6.3: Brønsted-Lowry Concept
- 1.6.3.1: Brønsted-Lowry Concept
- 1.6.3.2: Rules of Thumb for thinking about the relationship between Molecular Structure and Brønsted Acidity and Basicity*
- 1.6.3.3: The acid-base behavior of binary element hydrides is determined primarily by the element's electronegativity and secondarily by the element-hydrogen bond strength.*
- 1.6.3.4: Brønsted-Lowry Superacids and the Hammett Acidity Function
- 1.6.3.5: Thermodynamics of Solution-Phase Brønsted Acidity and Basicity
- 1.6.3.6: Thermodynamics of Gas Phase Brønsted Acidity and Basicity
- 1.6.3.7: The Acidity of an Oxoacid is Determined by the Electronegativity and Oxidation State of the Oxoacid's Central Atom*
- 1.6.3.8: High Charge-to-Size Ratio Metal Ions Act as Brønsted Acids in Water
- 1.6.3.9: The Solvent System Acid Base Concept
- 1.6.3.10: Acid-Base Chemistry in Amphoteric Solvents and the Solvent Leveling Effect
- 1.6.3.11: Non-nucleophilic Brønsted-Lowry Superbases
- 1.6.4: Lewis Concept and Frontier Orbitals
- 1.6.4.1: The frontier orbital approach considers Lewis acid-base reactions in terms of the donation of electrons from the base's highest occupied orbital into the acid's lowest unoccupied orbital.
- 1.6.4.2: All other things being equal, electron withdrawing groups tend to make Lewis acids stronger and bases weaker while electron donating groups tend to make Lewis bases stronger and acids weaker
- 1.6.4.3: The electronic spectra of charge transfer complexes illustrate the impact of frontier orbital interactions on the electronic structure of Lewis acid-base adducts
- 1.6.4.4: Substances' solution phase Lewis basicity towards a given acid may be estimated using the enthalphy change for dissociation of its adduct with a reference acid of similar hardness.
- 1.6.4.5: In the boron trifluoride affinity scale, the enthalphy change on formation of an adduct between the base and boron trifluoride is taken as a measure of Lewis basicity.
- 1.6.4.6: Lewis base strength may also be estimated by measuring structural or energy changes upon formation of a Lewis acid-base complex, as illustrated by efforts to spectroscopically assess the strengths of halogen bonds
- 1.6.4.7: Bulky groups weaken the strength of Lewis acids and bases because they introduce steric strain into the resulting acid-base adduct.
- 1.6.4.8: Frustrated Lewis pair chemistry uses Lewis acid and base sites within a molecule that are sterically restricted from forming an adduct with each other.
- 1.6.5: Intermolecular Forces
- 1.6.5.1: Host-Guest Chemistry and π-π Stacking Interactions
- 1.6.5.2: Hydrogen bonds may be considered as a special type of Lewis acid-base interaction in which a Lewis acid hydrogen ion is shared between Lewis bases
- 1.6.6: Hard and Soft Acids and Bases
- 1.6.6.1: Quantitative Measures of Hardness, Softness, and Acid-Base Interactions from a Hard Soft Acid-Base Principle Perspective Involve Orbital Energies and/or Apportioning Acid-Base Bonding in Terms of Electrostatic and Covalent Factors
- 1.6.6.2: Hard-Hard and Soft-Soft preferences may be explained and quantified in terms of electrostatic and covalent and electronic stabilization on the stability of Lewis acid-base adducts
- 1.6.7: Problems