CHEM-3010: Spring 2026

by Yulia Zaikina

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Iowa State University
Subjects
Chemistry
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CC BY 4.0, as LibreTexts states it
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Tweens profiles and older, judged from its title and subjects

Contents

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  1. Front Matter
  2. TitlePage
  3. InfoPage
  4. Table of Contents
  5. Licensing
  6. 1: Introduction to Inorganic Chemistry
  7. 1.1: What is Inorganic Chemistry?
  8. 1.2: Inorganic vs Organic Chemistry
  9. 1.3: History of Inorganic Chemistry
  10. 1.4: Perspectives
  11. 1.5: Practice problems
  12. 2: Atomic Structure
  13. 2.1: Historical Development of Atomic Theory
  14. 2.1.1: The Periodic Table
  15. 2.1.2: Discovery of Subatomic Particles and the Bohr Atom
  16. 2.2: The Schrödinger equation, particle in a box, and atomic wavefunctions
  17. 2.2.1: Particle in a Box
  18. 2.2.2: Quantum Numbers and Atomic Wave Functions
  19. 2.2.3: Aufbau Principle
  20. 2.2.4: Shielding
  21. 2.3: Periodic Properties of Atoms
  22. 2.3.1: Ionization energy
  23. 2.3.2: Electron Affinity
  24. 2.3.3: Covalent and Ionic Radii
  25. 2.4: Problems
  26. 3: Simple Bonding Theory
  27. 3.1: Lewis Electron-Dot Diagrams
  28. 3.1.1: Resonance
  29. 3.1.2: Breaking the octet rule with higher electron counts (hypervalent atoms)
  30. 3.1.3: Formal Charge
  31. 3.1.4: Lewis Fails to Predict Unusual Cases - Boron and Beryllium
  32. 3.2: Valence Shell Electron-Pair Repulsion
  33. 3.2.1: Lone Pair Repulsion
  34. 3.2.2: Multiple Bonds
  35. 3.2.3: Electronegativity and Atomic Size Effects
  36. 3.2.4: Ligand Close Packing
  37. 3.3: Molecular Polarity
  38. 3.4: Hydrogen Bonding
  39. 4: Symmetry and Group Theory
  40. 4.1: Symmetry Elements and Operations
  41. 4.2: Point Groups
  42. 4.2.1: Groups of Low and High Symmetry
  43. 4.2.2: Other Groups
  44. 4.3: Properties and Representations of Groups
  45. 4.3.1: Matrices
  46. 4.3.2: Representations of Point Groups
  47. 4.3.3: Character Tables
  48. 4.4: Examples and Applications of Symmetry
  49. 4.4.1: Chirality
  50. 4.4.2: Molecular Vibrations
  51. 4.P: Problems (under construction)
  52. 5: Molecular Orbitals
  53. 5.1: Formation of Molecular Orbitals from Atomic Orbitals
  54. 5.1.1: Molecular Orbitals from s Orbitals
  55. 5.1.2: Molecular Orbitals from p Orbitals
  56. 5.1.3: Molecular orbitals from d orbitals
  57. 5.1.4: Nonbonding Orbitals and Other Factors
  58. 5.2: Homonuclear Diatomic Molecules
  59. 5.2.1: Molecular Orbitals
  60. 5.2.2: Orbital Mixing
  61. 5.2.3: Diatomic Molecules of the First and Second Periods
  62. 5.2.4: Photoelectron Spectroscopy
  63. 5.3: Heteronuclear Diatomic Molecules
  64. 5.3.1: Orbital ionization energies
  65. 5.3.2: Polar bonds
  66. 5.3.3: Ionic Compounds and Molecular Orbitals
  67. 5.4: Larger (Polyatomic) Molecules
  68. 5.4.1: Bifluoride anion
  69. 5.4.2: Carbon Dioxide
  70. 5.4.3: H₂O
  71. 5.4.4: NH₃
  72. 5.4.5: CO₂ (Revisted with Projection Operators)
  73. 5.4.6: BF₃
  74. 5.P: Problems
  75. 6: Acid-Base and Donor-Acceptor Chemistry
  76. 6.1: Acid-Base Models as Organizing Concepts
  77. 6.2: Lewis Concept and Frontier Orbitals
  78. 6.2.1: 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.
  79. 6.2.2: 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
  80. 6.2.3: The electronic spectra of charge transfer complexes illustrate the impact of frontier orbital interactions on the electronic structure of Lewis acid-base adducts
  81. 6.2.4: Bulky groups weaken the strength of Lewis acids and bases because they introduce steric strain into the resulting acid-base adduct.
  82. 6.2.5: Frustrated Lewis pair chemistry uses Lewis acid and base sites within a molecule that are sterically restricted from forming an adduct with each other.
  83. 6.3: Hard and Soft Acids and Bases
  84. 6.3.1: 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
  85. 7: Chemistry of the Main Group Elements
  86. 7.1: General Trends in Main Group Chemistry
  87. 7.1.1: The Periodic Table is an Organizing Concept in Main Group Chemistry
  88. 7.1.1.1: The metal-nonmetal-metalloid distinction and the metal-nonmetal "line" are useful for thinking about trends in elements' physical properties
  89. 7.1.1.2: There are qualitative differences between the chemistry of the elements in the first two rows and those in the rest of the periodic table
  90. 7.1.2: Electronegativity increases and radius decreases towards the upper left of the periodic table, with electron withdrawing substituents, and with oxidation state
  91. 7.1.3: Ionization energy roughly increases towards the upper left of the periodic table but is also influenced by orbital energy and pairing energy effects
  92. 7.1.4: As may be seen from considering an element's redox diagrams, main group elements (aside from the noble gases) generally are more oxidizing towards the upper left of the periodic table and more reducing towards the lower right of the periodic table
  93. 7.1.4.1: Latimer Diagrams summarize elements' redox properties on a single line
  94. 7.1.4.2: Frost Diagrams show how stable element's redox states are relative to the free element
  95. 7.1.4.3: Pourbaix Diagrams are Redox Phase Diagrams that Summarize the most stable form of an element at a given pH and solution potential
  96. 7.2: What are the main group elements and why should anyone care about them?
  97. 7.3: Group 1, The Alkali Metals
  98. 7.3.1: Alkali Metals' Chemical Properties
  99. 7.4: Hydrogen
  100. 7.4.1: Hydrogen's Chemical Properties
  101. 7.5: Group 2, The Alkaline Earth Metals
  102. 7.5.1: Preparation and General Properties of the Alkaline Earth Elements
  103. 7.5.2: Alkaline Earth Metals' Chemical Properties
  104. 7.6: Group 13 (and a note on the post-transition metals)
  105. 7.6.1: Properties of the Group 13 Elements and Boron Chemistry
  106. 7.6.2: Heavier Elements of Group 13 and the Inert Pair Effect
  107. 7.7: Group 14
  108. 7.7.1: The Group 14 Elements and the Many Allotropes of Carbon
  109. 7.7.2: Inorganic Compounds of the Group 14 Elements
  110. 7.7.3: Chemistry of Carbon (Z=6)
  111. 7.7.4: Chemistry of Silicon (Z=14)
  112. 7.7.4.1: Silicates
  113. 7.7.4.2: Silicon and Group 14 Elements
  114. 7.7.5: Chemistry of Germanium (Z=32)
  115. 7.7.6: Chemistry of Tin (Z=50)
  116. 7.7.7: Chemistry of Lead (Z=82)
  117. 7.7.7.1: Lead Plumbate
  118. 7.8: The Nitrogen Family
  119. 7.8.1: General Properties and Reactions
  120. 7.8.1.1: Nitrogen Group (Group 5) Trends
  121. 7.8.2: Chemistry of Nitrogen (Z=7)
  122. 7.8.3: Chemistry of Phosphorus (Z=15)
  123. 7.8.4: Chemistry of Antimony (Z=51)
  124. 7.8.5: Chemistry of Bismuth (Z=83)
  125. 7.8.6: Chemistry of Moscovium (Z=115)
  126. 7.9: The Oxygen Family (The Chalcogens)
  127. 7.9.1: General Properties and Reactions
  128. 7.9.1.1: Oxygen Group (Group VIA) Trends
  129. 7.9.2: Chemistry of Oxygen (Z=8)
  130. 7.9.2.1: Ozone
  131. 7.9.3: Chemistry of Sulfur (Z=16)
  132. 7.9.4: Chemistry of Selenium (Z=34)
  133. 7.9.5: Chemistry of Tellurium (Z=52)
  134. 7.9.6: Chemistry of Polonium (Z=84)
  135. 7.9.7: Chemistry of Livermorium (Z=116)
  136. 7.10: The Halogens
  137. 7.10.1: Physical Properties of the Halogens
  138. 7.10.1.1: Atomic and Physical Properties of Halogens
  139. 7.10.1.2: General Properties of Halogens
  140. 7.10.1.3: Halogen Group (Group 17) Trends
  141. 7.10.1.4: Physical Properties of the Group 17 Elements
  142. 7.10.2: Chemical Properties of the Halogens
  143. 7.10.2.1: Halide Ions as Reducing Agents
  144. 7.10.2.2: Halogens as Oxidizing Agents
  145. 7.10.2.3: Interhalogens
  146. 7.10.2.4: More Reactions of Halogens
  147. 7.10.2.5: Oxidizing Ability of the Group 17 Elements
  148. 7.10.2.6: Testing for Halide Ions
  149. 7.10.2.7: The Acidity of the Hydrogen Halides
  150. 7.10.3: Chemistry of Fluorine (Z=9)
  151. 7.10.4: Chemistry of Chlorine (Z=17)
  152. 7.10.4.1: The Manufacture of Chlorine
  153. 7.10.5: Chemistry of Bromine (Z=35)
  154. 7.10.6: Chemistry of Iodine (Z=53)
  155. 7.10.7: Chemistry of Astatine (Z=85)
  156. 7.11: The Noble Gases
  157. 7.11.1: Properties of Nobel Gases
  158. 7.11.1.1: Noble Gas (Group 18) Trends
  159. 7.11.2: Reactions of Nobel Gases
  160. 7.11.3: Chemistry of Helium (Z=2)
  161. 7.11.4: Chemistry of Neon (Z=10)
  162. 7.11.5: Chemistry of Argon (Z=18)
  163. 7.11.6: Chemistry of Krypton (Z=36)
  164. 7.11.7: Chemistry of Xenon (Z=54)
  165. 7.11.8: Chemistry of Radon (Z=86)
  166. 8: Introduction to Solid State Chemistry
  167. 8.1: The Nature of Crystalline Solids
  168. 8.1.1: Crystal Structures of Metals
  169. 8.1.2: Ionic and Covalent Solids - Structures
  170. 8.1.2.1: Prelude to Ionic and Covalent Solids - Structures
  171. 8.1.2.2: Close-packing and Interstitial Sites
  172. 8.1.2.3: Structures Related to NaCl and NiAs
  173. 8.1.2.4: Tetrahedral Structures
  174. 8.1.2.5: Layered Structures and Intercalation Reactions
  175. 8.1.2.6: Bonding in TiS₂, MoS₂, and Pyrite Structures
  176. 8.1.2.7: Spinel, Perovskite, and Rutile Structures
  177. 8.1.2.8: Discussion Questions
  178. 8.1.2.9: References
  179. 8.2: Bonding in Metals, Semiconductors and Insulators – Band Structure
  180. 8.2.1: Prelude to Electronic Properties of Materials - Superconductors and Semiconductors
  181. 8.2.2: Periodic Trends- Metals, Semiconductors, and Insulators
  182. 8.2.3: Bonding in Metals
  183. 8.2.4: Conduction in Metals
  184. 8.2.5: Superconductors
  185. 8.2.6: Semiconductors- Band Gaps, Colors, Conductivity and Doping
  186. 8.2.7: Semiconductor p-n Junctions
  187. 8.2.8: Diodes, LEDs and Solar Cells
  188. 8.3: X-rays and X-ray Diffraction
  189. 8.4: The Imperfect Solid State
  190. 8.5: Glasses
  191. 8.6: Phase Equilibria and Phase Diagrams
  192. Back Matter
  193. Index
  194. Glossary
  195. Detailed Licensing
  196. Detailed Licensing