CHE002B: General Chemistry

thermodynamics

• first law of thermodynamics
• entropy
• Gibbs free energy
• work
• ideal gas law
• Hess’s law
• state function
• enthalpy of formation
• enthalpy of reaction
• spontaneous reaction
• non-spontaneous reaction

quantum mechanics

• Heisenberg uncertainty principle
• Rydberg’s formula
• Aufbau principle
• Hund’s rule
• Pauli exclusion principle
• Schrodinger’s equation
• quantum number
  • principal quantum number
  • azimuthal quantum number
  • magnetic quantum number,
  • spin quantum number
• nodes
  • radial node
  • angular node
• orbital subscript
• periodic table
  • first two lanthanides have d=1, for the rest, the d=1 is given to f instead
  • Actinides are really unpredictable.
• self-consistent field
• penetration effect
• periodic trend
• electron affinity
• N cannot accept e- to become N- since e- would have to get into an already occupied orbital, which would make it very unstable, thus it has a low electron affinity. In contrast, O can accept an e- because the increase in effective nuclear charge is apparently enough to overcome the instability
• electron affinity generally decreases down the periodic table

bonding

• ionic compound

electronegativity

• expected $\mathrm{H}-\mathrm{X}$ bond energy = $[(\mathrm{H}-\mathrm{H}\text{ bond energy})(\mathrm{X}-\mathrm{X}\text{ bond energy}){]}^{1/2}$ (geometric mean)
• Due to electronegativity differences between elements, the actual bond energy differs from the expected since electrons are shared unevenly---electrons are more attracted towards the atom with higher electronegativity
  • e.g. if $\mathrm{X}$ has a higher electronegativity, $\mathrm{H}$ will be partially positive and $\mathrm{X}$ will be partially negative
• $\Delta =(\mathrm{H}-\mathrm{X}{)}_{\text{act}}-({\mathrm{H}-\mathrm{X}}_{\text{exp}})$ = the difference between actual bond energy and expected bond energy
• Electronegativity difference: $\mathrm{EN}(\mathrm{X})-\mathrm{EN}(\mathrm{H})=0.102\sqrt{\Delta }$
  • 0.102 is the conversion factor between $\mathrm{kJ}$ and $\mathrm{eV}$ (the latter of which was the original used when the formula was proposed)
  • bond energies in $\Delta$ should be $\mathrm{kJ}/\mathrm{mol}$

electronegativity difference	bond type
zero	covalent
intermediate	polar covalent
large	ionic

dipole moment

• if dipole moments don’t cancel out, the molecule is polar
• generally, symmetric = nonpolar (dipole moment might cancel out), asymmetric = polar
  • CANNOT BE USED AS JUSTIFICATION, TEST IT!
• ionic character