CHEM 205 lecture 12

• Z*


  Effective nuclear charge = net positive charge experienced by an electron in a multi-electron atom

  
• Effective nuclear charge


  • The higher the Z*, the stronger the pull of an atom has on electrons of its own and on others

  
• Below 5th row, changes in size are very small cause energy of high energy orbitals are very similar to one another

  
• Cations


  Fewer electrons, smaller than parent atom, less total shielding = electrons are held more tightly

  
• Anions


  More electrons, larger than parent atoms, more total shielding = electrons are held more loosely

  
• Ionization energy


  Quantity of energy required to remove an electron from a gas-phase atom or ion
• Electron affinity and electron attachment enthalpy


  Enthalpy change = type of energy in thermoregulation that is changed when one electron is added to a gas-phase atom
• Noble gas configurations


  • Difficult to change:

  1. Add another electron: Energetically unfavourable cause higher energy levels are closer together, smaller repulsion: larger valence orbitals = weaker interaction with effective nuclear charge

  2. Remove an electron: Removing a core electron requires lots of energy

  
• Main group metals

  • Form cations with noble gas electron configuration, higher charged ions are unfavourable because: you need to remove electrons from full octet, quantum number for the lower energy level is very high = strong attraction to nucleus, requires high energy to remove electrons

  
• Transition metals


  • More complicated, common ions arise from *not usually noble gas configurations but still stable*: emptying the s subshell, emptying or half-emptying the d subshell
• Valence bond theory


  "Localized electron model": Bonding electron pairs localized between pairs of atoms, lone pair electrons on single atoms. Mainly used for: Visual pictures of molecular structure and bonding, Molecules in ground state

  
• Molecular orbital theory


  "Delocalized electron model": Bonding electrons spread out over entire molecule in molecular orbitals. Particularly useful for: Quantitative views, Excited states, Paramagnetic compounds with even number of electrons

  
• Valence bond theory: Orbital overlap model of bonding

  2 overlapping orbitals + 2 electrons = one bond called a sigma bond (σ-bond)

  
• Valence bond theory: Hybridization of atomic orbitals


  Atoms respond as needed to make the newly forming molecules have lowest energy possible: As bonds begin to form, central atoms own valence shells orbitals mix together = new formed hybrid orbitals, overlap with outer atoms to permit electron pairs to be as far away as possible in the molecule. Any orbitals from an atom's valence shell can be used.

  
• Identifying hybrid orbitals involved in bonding


  Draw Lewis structure

  2. One hybrid orbital is required per electron pair (treat multiple bonds as single)

  3. If mix 2 atomic orbitals = 2 hybrid orbitals form, etc...

  4. Hybrid orbitals are named after their parent atomic orbitals

  5. Start with the s orbital + as many p and d orbitals as needed to make enough hybrid orbitals to hold all the σ-bonding pairs and lone pairs around the central atom

  
• If some atomic orbitals are not involved in hybridization, one p orbital is left over in unhybridized atomic orbital containing an electron, it overlaps the p orbital on the neighbouring element atom = forming a pi-bond

  
• Sigma (σ) bond


  Centers along the internuclear axis: electron pair lies between nuclei = a regular single bond, orbitals overlap end-on: efficient overlap = strong bond, σ-bonds can rotate without breaking
• Pi (π) bond


  Occupies the space above and below internuclear axis, orbitals overlap side-on: less efficient overlap = weaker than σ, π-bond must also exist between atoms, π-bond cannot rotate without breaking, they are rigid

  
• Resonance


  1. orbitals on many atoms overlap at once, electrons are delocalized over many atoms