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MIT 5.111SC: Principles of Chemical Science

Contents · 70

  1. 01-The Importance of Chemical PrinciplesReviews course policies and introductory chemistry fundamentals, including significant figures, SI prefixes, forces, energy, atoms, and the nuclear atom.
  2. 01-The Importance of Chemical PrinciplesIntroduction to the chemistry course, its teaching resources, and the instructor’s path to appreciating chemistry and its relevance to students’ futures.
  3. 02-Atomic StructureCovers electron and nucleus discoveries, Rutherford’s gold foil experiment, and why classical mechanics cannot explain atomic structure.
  4. 02-Atomic StructureDiscovery of the electron and nucleus, including cathode-ray experiments and how scientists inferred the properties of subatomic particles.
  5. 03-Wave-Particle Duality of LightWave properties of light, including wavelength, frequency, the electromagnetic spectrum, and interference; photon energy and the photoelectric effect.
  6. 03-Wave-Particle Duality of LightWave properties including amplitude, wavelength, frequency, period, and intensity, plus light’s wave-particle duality and the photoelectric effect.
  7. 04-Wave-Particle Duality of Matter; Schrödinger EquationExplains the photoelectric effect and photon momentum, de Broglie matter waves and electron diffraction, and the Schrödinger equation.
  8. 04-Wave-Particle Duality of Matter; Schrödinger EquationExplains photoelectric-effect thresholds, photon and electron properties, energy units, and a demonstration, then introduces photon momentum and matter waves.
  9. 05-Hydrogen Atom Energy LevelsHydrogen energy levels and ionization energies, Schrödinger equation predictions, and photon emission and absorption spectra, including the Balmer series.
  10. 05-Hydrogen Atom Energy LevelsHydrogen atom energy levels, binding and ionization energies, and experimental validation of Schrödinger equation predictions.
  11. 06-Hydrogen Atom Wavefunctions (Orbitals)Hydrogen-atom orbitals, quantum numbers, orbital shapes and sizes, nodes, electron probability distributions, spin, and the Pauli exclusion principle.
  12. 06-Hydrogen Atom Wavefunctions (Orbitals)Hydrogen-atom orbitals, quantum numbers, and chemistry labels for shells, subshells, and orbital orientations.
  13. 07-Multi-electron AtomsExplains how shielding and effective nuclear charge determine multielectron orbital energies, then covers electron configurations, filling rules, and ion configurations.
  14. 07-Multi-electron AtomsSimilarities and differences between one- and multi-electron orbitals, including orbital shapes, energy levels, effective nuclear charge, and electron shielding.
  15. 08-The Periodic Table and Periodic TrendsExplains periodic trends in ionization energy and electron affinity, including photoelectron spectroscopy, subshell effects, and successive ionization energies.
  16. 08-The Periodic Table and Periodic TrendsIntroduces the periodic table, its organization by electron configuration, and how valence electrons explain element properties and reactivity.
  17. 09-Periodic Table; Ionic and Covalent BondsElectronegativity and atomic and ionic radii trends, isoelectronic species, and the formation and electrostatic energy of ionic bonds.
  18. 09-Periodic Table; Ionic and Covalent BondsExplains electronegativity, its relationship to ionization energy and electron affinity, periodic trends, and its role in pharmaceutical molecules.
  19. 10-Introduction to Lewis StructuresIntroduces Lewis structures, the octet rule, and formal charge, then uses resonance structures to explain electron delocalization in ozone.
  20. 10-Introduction to Lewis StructuresIntroduces Lewis structures, valence and lone-pair electrons, octets, hydrogen’s duet, and rules for arranging atoms and counting electrons.
  21. 11-Lewis Structures: Breakdown of the Octet RuleExplains octet-rule exceptions: radicals, electron-deficient molecules, and expanded valence shells, with examples including O₂, BF₃, PCl₅, and chromate.
  22. 11-Lewis Structures: Breakdown of the Octet RuleOdd-electron molecules and Lewis structures, radical reactivity and roles in biology, and nitric oxide signaling and vasodilation.
  23. 12-The Shapes of Molecules: VSEPR TheoryExplains VSEPR theory for predicting molecular shapes, bond angles, and geometries with and without lone pairs.
  24. 12-The Shapes of Molecules: VSEPR TheoryIntroduces VSEPR theory and steric number to predict molecular shapes from Lewis structures, emphasizing how molecular geometry influences chemical and biological properties.
  25. 13-Molecular Orbital TheoryExplains molecular orbital formation, bonding and antibonding orbitals, electron configurations, and bond order in homonuclear and heteronuclear diatomic molecules.
  26. 13-Molecular Orbital TheoryIntroduces molecular orbital theory and LCAO, explaining how constructive interference of atomic s orbitals forms bonding molecular orbitals.
  27. 14-Valence Bond Theory and HybridizationValence bond theory, sigma and pi bonds, orbital hybridization (sp, sp², sp³), molecular geometry, and hybridization in complex molecules.
  28. 14-Valence Bond Theory and HybridizationValence bond theory explains bonding through paired electrons, while hybridization and electron promotion account for molecular shapes and carbon’s bonding. Defines sigma and pi bonds and their roles in single, double, and triple bonds.
  29. 15-Thermodynamics: Bond and Reaction EnthalpiesBond enthalpies, reaction enthalpies, and methods for calculating reaction enthalpy using bond energies, formation enthalpies, and Hess’s law.
  30. 15-Thermodynamics: Bond and Reaction EnthalpiesIntroduces thermodynamics and kinetics, contrasts spontaneity with reaction rate, and explains bond enthalpy, bond dissociation, and the energy changes of bond breaking and formation.
  31. 16-Thermodynamics: Gibbs Free Energy and EntropyExplains spontaneity through Gibbs free energy, entropy changes, and standard free energies of formation.
  32. 16-Thermodynamics: Gibbs Free Energy and EntropyExplains spontaneity using Gibbs free energy, the roles of enthalpy and entropy, and calculating ΔG with consistent units and significant figures.
  33. 17-Thermodynamics: Now What Happens When You Heat It Up?How temperature affects reaction spontaneity, hydrogen bonding in biological systems, and ATP-coupled reactions.
  34. 17-Thermodynamics: Now What Happens When You Heat It Up?Temperature-dependent spontaneity of sodium bicarbonate decomposition, including the ΔG–temperature relationship and transition temperature.
  35. 18-Introduction to Chemical EquilibriumExplains dynamic equilibrium, reaction quotients and equilibrium constants, the relationship between K and Gibbs free energy, and how concentration changes shift equilibrium.
  36. 18-Introduction to Chemical EquilibriumIntroduces chemical equilibrium as a dynamic balance between forward and reverse reaction rates, with announcements about exams and thermodynamics assignments.
  37. 19-Chemical Equilibrium: Le Châtelier's PrincipleCovers how volume, inert gases, and temperature affect equilibrium, the Van’t Hoff equation, reaction-yield optimization, hemoglobin and oxygen, and significant figures in logarithms.
  38. 19-Chemical Equilibrium: Le Châtelier's PrincipleExplains Le Chatelier’s principle and how changing the volume of a gaseous system affects pressure and chemical equilibrium.
  39. 20-Solubility and Acid-Base EquilibriumMolar solubility, solubility equilibria, solvent effects, pressure and temperature, and the thermodynamics of dissolving; Arrhenius and Brønsted–Lowry acid–base definitions.
  40. 20-Solubility and Acid-Base EquilibriumIntroduces solubility, solutions, solvents and solutes, how water dissolves ionic and molecular substances, and writing solubility product expressions.
  41. 21-Acid-Base Equilibrium: Is MIT Water Safe to Drink?Defines pH, pOH, and acid–base strength, relates conjugate acid–base constants, and solves weak-acid and weak-base equilibrium problems.
  42. 21-Acid-Base Equilibrium: Is MIT Water Safe to Drink?Reviews Lewis and Brønsted–Lowry acid–base definitions, then introduces water ionization, equilibrium constants, pKw, pH, pOH, and acid–base strength.
  43. 22-Acid-Base Equilibrium: Salt Solutions and BuffersExplains how salts affect solution pH, how acid and base buffers resist pH changes, and how to calculate and design buffers using the Henderson–Hasselbalch equation.
  44. 22-Acid-Base Equilibrium: Salt Solutions and BuffersExplains how salt ions affect solution pH, using conjugate acid/base strength to classify salts as acidic, neutral, or basic.
  45. 23-Acid-Base Titrations Part IDescribes strong- and weak-acid/base titration curves, equivalence and buffering points, indicators, and pH calculations before, at, and after equivalence.Strong- and weak-acid/base titration curves, equivalence and buffering points, indicators, and pH calculations throughout a titration.
  46. 23-Acid-Base Titrations Part IStrong-acid and strong-base titration curves, equivalence points, endpoints, and pH indicators, including a red-cabbage dye demonstration.
  47. 24-Acid-Base Titrations Part IIWeak-acid/strong-base titration at and beyond the equivalence point, pKa and pH, oxidation states, redox reactions, and balancing redox equations in acidic solution.
  48. 24-Acid-Base Titrations Part IICalculates the pH at the equivalence point and beyond in a weak-acid–strong-base titration, using conjugate-base hydrolysis and excess strong base.
  49. 25-Oxidation-Reduction and Electrochemical CellsIntroduces oxidation states, redox agents and disproportionation, then explains balancing redox reactions in acidic and basic solutions and electrochemical cells, including electrodes, cell notation, and Faraday’s law.
  50. 25-Oxidation-Reduction and Electrochemical CellsOxidation–reduction reactions, electron transfer, oxidizing and reducing agents, and rules for assigning oxidation numbers.
  51. 26-Chemical and Biological OxidationsConnects cell potential with Gibbs free energy, explains standard reduction potentials and redox strength, and derives the Nernst equation for nonstandard conditions and equilibrium.
  52. 26-Chemical and Biological OxidationsRelates cell potential to Gibbs free energy and spontaneity, and calculates standard cell potentials from electrode reduction potentials. Compares galvanic and electrolytic cells.
  53. 27-Introduction to Transition MetalsTransition-metal biology, coordination complexes, chelation, geometric and optical isomers, d-electron counting, and d-orbital shapes.
  54. 27-Introduction to Transition MetalsIntroduces transition metals, their d-orbital properties and biological roles, with a focus on nickel in H. pylori urease.
  55. 28-Transition Metals: Crystal Field Theory Part IExplains octahedral d-orbital splitting, the spectrochemical series, high- and low-spin electron configurations, crystal field stabilization energy, and the color and magnetic properties of coordination complexes.
  56. 28-Transition Metals: Crystal Field Theory Part IIntroduces crystal field theory as a model for how ligands split metal d-orbital energy levels, explaining transition-metal colors and magnetic properties, with a focus on octahedral complexes.
  57. 29-Transition Metals: Crystal Field Theory Part IILigand type and geometry effects on spectroscopic properties, crystal-field splitting in tetrahedral and square-planar complexes, and nickel enzyme applications.
  58. 29-Transition Metals: Crystal Field Theory Part IIHow ligand strength and coordination geometry affect d-orbital splitting, observed color, and magnetism in nickel complexes.
  59. 30-Kinetics: Rate LawsChemical reaction rates, rate expressions, experimentally determined rate laws, and integrated first-order kinetics, including half-life.
  60. 30-Kinetics: Rate LawsIntroduces chemical kinetics, distinguishing reaction rates from thermodynamic stability and spontaneity. Surveys factors that affect rates and uses the oscillating clock reaction to connect kinetics with equilibrium, acid-base chemistry, and redox.
  61. 31-Nuclear Chemistry and Chemical KineticsCovers radioactive decay rates and half-lives, second-order integrated rate laws, the relationship between rate constants and equilibrium constants, and elementary reaction steps and molecularity.
  62. 31-Nuclear Chemistry and Chemical KineticsRadioactive decay as a first-order process, its decay law and half-life, with applications in medical imaging, nuclear energy, and waste storage.
  63. 32-Kinetics: Reaction MechanismsUses experimental rate laws, steady-state approximations, and fast-equilibrium assumptions to evaluate reaction mechanisms and identify rate-determining steps.
  64. 32-Kinetics: Reaction MechanismsReaction mechanisms, elementary steps, and how to test whether a proposed mechanism matches experimental rate laws, using NO oxidation as an example.
  65. 33-Kinetics and TemperatureExplains the Arrhenius equation and how temperature affects rate constants, activation energy, reaction coordinates, activated complexes, and reaction mechanisms.
  66. 33-Kinetics and TemperatureArrhenius equation, activation energy, and how temperature affects reaction rate constants.
  67. 34-Kinetics: CatalystsExplains catalyst types and how catalysis lowers activation energy, then develops enzyme catalysis, Michaelis–Menten kinetics, and enzyme inhibition.
  68. 34-Kinetics: CatalystsHow catalysts lower activation-energy barriers, speed reaction rates without changing equilibrium, and differ by phase.
  69. 35-Applying Chemical PrinciplesReviews core chemistry concepts through CO₂ conversion, including molecular structure, thermodynamics, redox, acid-base chemistry, catalysis, and enzyme equilibria in microbial acetogenesis.
  70. 35-Applying Chemical PrinciplesApplying core chemistry concepts to CO₂ capture and conversion into biofuels, including Lewis structures, biological pathways, and catalysis.