General Chemistry

From the inside of an atom to why a reaction goes, and how fast

About this course

The first course in chemistry — atomic structure, the periodic law, bonding and molecular shape, the states of matter, thermodynamics, equilibrium, acids and bases, kinetics and electrochemistry.

What you will be able to do

  • Describe electron configurations and use them to explain periodic trends
  • Predict molecular shape and polarity, and build simple molecular orbital diagrams
  • Relate intermolecular forces to the behaviour of gases, liquids and solids
  • Use enthalpy, entropy and free energy to decide whether a reaction is spontaneous
  • Calculate equilibrium, pH, rate and cell potential for a given system

What this course covers

  • Conservation of Mass and Chemical Equations
  • Balancing Equations with Coefficients
  • The Mole as a Unit of Count
  • The Mole–Mass–Particle Conversion Bridge
  • Stoichiometry: the Three-Junction Road
  • Limiting Reagent and Percent Yield
  • The Unified Stoichiometry Diagram and Course Spine
  • Bond formation as energy minimisation
  • The Born-Haber cycle and ionic bonding
  • Covalent bonding and the octet rule
  • Ionic vs covalent as a continuum
  • Formal charge and resonance
  • Bond order, sigma and pi bonds
  • Four named failures of the octet rule
  • Rate laws are measured, not read from equations
  • Integrated rate laws and the half-life diagnostic
  • Radiocarbon dating as a first-order application
  • Mechanisms, elementary steps, and rate-law fingerprinting
  • The activation barrier and the Arrhenius equation
  • What a catalyst is worth: barrier lowering and enzyme kinetics
  • Arrhenius measurement and the independence of barrier from thermodynamics
  • Heat, work, and the first law
  • Enthalpy: definition and convenience
  • Calorimetry: coffee cup vs. bomb calorimeter
  • Hess's law and state functions
  • Standard enthalpies of formation and Hess's law as a calculating machine
  • Enthalpy as necessary but insufficient for predicting direction
  • Applications of the formation-enthalpy table
  • The Four Empirical Gas Laws
  • Absolute Zero from Extrapolation
  • The Ideal Gas Equation as a Working Tool
  • Kinetic Molecular Theory and the Meaning of Temperature
  • Root-Mean-Square Speed and the Maxwell-Boltzmann Distribution
  • Graham's Law of Effusion
  • Real Gases and the Van der Waals Equation
  • VSEPR: electron domains determine geometry
  • The tetrahedral angle as a geometric derivation
  • Lone-pair compression of bond angles
  • Molecular dipole as vector sum of bond dipoles
  • Molecular polarity and the like-dissolves-like rule
  • Xenon tetrafluoride: full VSEPR worked example
  • Entropy as a count of arrangements
  • The second law applied to the universe
  • Gibbs free energy as the second law rearranged
  • Four sign cases and the crossover temperature
  • Thermodynamics vs. kinetics: the Haber process tension
  • Unification and course position of entropy
  • Why weak forces determine physical state
  • London dispersion force
  • Dipole-dipole and hydrogen bonding forces
  • Dispersion, molecular shape, and contact area
  • Procedure for ranking boiling points
  • Hydrogen bonding and the water anomaly
  • Consequences of the hydrogen-bond network in water
  • Classical physics fails to explain atomic stability
  • Rutherford scattering and the nuclear atom
  • Balmer's empirical formula for hydrogen lines
  • Rydberg form and the energy-level ladder
  • Bohr's quantisation assumption and the hydrogen energy ladder
  • Failure of Bohr's model beyond one electron
  • Working procedure and the 1240 eV·nm rule
  • Redox bookkeeping and oxidation numbers
  • Cell voltage as free energy per unit charge
  • The Nernst equation
  • Concentration cells and voltage as distance from equilibrium
  • Kinetics vs thermodynamics in corrosion
  • Corrosion, sacrificial anodes, and electrolysis
  • Course-spanning argument and closing worked example
  • Failure of Lewis structures: oxygen's paramagnetism
  • Formation of bonding and antibonding molecular orbitals
  • Bond order as bonding minus antibonding over two
  • MO diagram ordering and s–p mixing for the second row
  • Filling the second-row diatomics: bond orders and magnetism
  • Effect of adding or removing electrons: the oxygen ion series
  • Delocalization, limits of Lewis structures, and theory replacement
  • Unit cell and atom counting
  • Sphere packing and the Kepler conjecture
  • Bragg's Law and X-ray diffraction
  • Determining Avogadro's number from a crystal
  • Thermodynamic stability vs. kinetic rate
  • Four kinds of solid
  • Phase diagrams, triple point, and critical point
  • De Broglie's wave-particle symmetry and standing-wave quantization
  • Standing waves on a loop: why in-between orbits vanish
  • Why the orbit dissolves: Heisenberg uncertainty
  • Orbitals as probability clouds; the Bohr radius as most-probable radius
  • Four quantum numbers and the shell-capacity formula 2n²
  • N+L filling order and the periodic table row lengths
  • Pauli, Aufbau, and Hund's rule; ionization-energy bumps and exceptions
  • Equilibrium as a dynamic state
  • The equilibrium constant K
  • The reaction quotient Q and the free-energy–mixture equation
  • Why catalysts do not shift equilibrium
  • Disturbing equilibrium: Q-based arithmetic vs. Le Chatelier's principle
  • The bridge equation: ΔG° = −RT ln K
  • Ammonia synthesis: integrating all four concepts
  • Effective Nuclear Charge and Slater's Rules
  • Atomic Size Trends and Ion Size
  • Ionization Energy Trends and Successive Ionizations
  • Electron Affinity and Its Anomalies
  • Electronegativity and Bond Character
  • Unified Reading Rules and Diagnostic Checks
  • Periodic Law as Explained Regularity
  • Three definitions of an acid
  • Water's autoionization and the pH scale
  • Weak acid equilibrium and dilution effect
  • The Henderson-Hasselbalch equation and buffer action
  • The titration curve and its four regions
  • Unifying summary: acid-base as one equilibrium constant
  • Blood as an open carbonic-acid buffer

Syllabus

  1. Atoms and the Periodic Law

    1. L1 · What Chemistry Explains
    2. L2 · Inside the Atom
    3. L3 · Electrons in Atoms
    4. L4 · The Periodic Law

  2. Bonding and Molecular Structure

    1. L5 · Ionic and Covalent Bonding
    2. L6 · Molecular Shape and Polarity
    3. L7 · Molecular Orbital Theory

  3. States of Matter

    1. L8 · Gases
    2. L9 · Intermolecular Forces
    3. L10 · Solids and Phase Change

  4. Energy and Equilibrium

    1. L11 · Thermochemistry
    2. L12 · Entropy and Free Energy
    3. L13 · Chemical Equilibrium
    4. L14 · Acids and Bases

  5. Rates and Electrochemistry

    1. L15 · Chemical Kinetics
    2. L16 · Electrochemistry and Send-off