CHEM 205 lecture 11

• Excited atoms/ions


  Emit characteristic colours of light
• Bohr's model


  • Could explain experimental observations for H and He: line spectrum, ionization energy
  • Can predict wavelengths of unseen lines in the spectrum
  • Introduced concept of energy quantization in describing atomic structure
• Bohr's model is a : simplification of the atomic structure and does not accurately describe the behavior of electrons in atoms with more than one electron. The idea of planet-like orbits is also incorrect, as electrons do not follow predictable paths around the nucleus. Instead, they exist in a probability cloud where they are most likely to be found.
• Particle-wave duality

  Waves can exhibit particle like properties, and particles can exhibit wave like properties
• Particles


  • Have mass and position can be specified
• Waves


  • Don't have mass and their position cannot be specified
  • Exhibit diffraction = spread around objects
  • Exhibit particulate properties: bundled up together moving as one
  • As objects shrink, their behaviour is less and less like that of particles, but more of that of waves
• Wavelength (w)


  w = h/mv
• Standing waves


  The number of nodes increases as the energy increases
• Orbitals

  Regions of space with high probability of finding an electron of particular energy

  
• wavefunction

  the probability of finding an electron at a given position
• Quantum numbers


  • Principal quantum number (n)
  • Angular momentum quantum number (l)
  • Magnetic quantum number (ml)
• Principal quantum number (n)


  Describes energy level or shell, also the number of shells and subshells
• Angular momentum quantum number (l)


  Describes orbitals shape in subshell

  
• Magnetic quantum number (ml)

  Describes orbitals orientation
• Atoms with 2 or more electrons can have the same principal quantum number: in the same shell = same or very similar energy</b>
• In each shell: subshells of different orbital shape based on angular momentum quantum numbers</b>
• based on magnetic quantum numbers value: In each subshell electrons can be in some orbital or in orbitals of different orientation
• Electron spin


  Electrons have a magnetic moment: they spin
• Electron spin

  An atomic orbital can contain 2 electrons as long as the electrons spins are opposite
• What is known about electrons:
• What is NOT known about electrons:
• Effective nuclear charge (Z*e')

  The average nuclear charge felt by a particular electron cause of shielding by other electrons
• 1. orbitals = spherical orbitals with the nucleus at its center, p-orbitals is dumbbell shapes… making s-orbitals always having a higher nuclear charge than p-orbitals

  
• Aufbau principle


  As protons are added 1 by 1 to the nucleus to build up the various elements, electrons are similarly added to H-like orbitals
• When adding electrons to an atom's configuration, add them 1 at a time, by counting across as atomic number (Z) increases: when Z goes up by 1, the number of electrons goes up by 1
• Elements in the same period/row have the same core electron configuration and valence electrons in the same shell
• Elements in the same block have valence electrons of the same type of orbital
• Valence electrons

  Highest energy in partly unfilled shells, usually with highest principle quantum number
• Fully filled d- and f- subshells are part of core, not valence, subtle changes in shielding result from occupying d-/f- orbitals… once filled they lower in energy than the next s-subshell
• Cannot use periodic table to predict electron configurations for all d- & f- block metals
• Electron configuration of ions

  If a nonmetal and a representative group metal react: electrons from valence orbitals of metal are lost = cations are formed (positive), electrons enter valence of nonmetals = anion are formed (negative)
• When electrons are lost to form cations, the electrons are removed from the shell with the highest "n" value
• Isoelectronic species

  atoms or ions that have the same number of electrons or electron configuration. This means that they have the same electronic structure and therefore the same electron density distribution, even if they have a different number of protons or atomic nuclei.
• To predict electron configurations of transition metal cations: remove valence electrons from subshell with highest n value first
• Diamagnetic


  Atoms repelled slightly by magnets, if all atoms electrons are in the same orbital as another electron with opposite spin
• Paramagnetism:

  Paramagnetism is a property of materials that are attracted to magnetic fields due to the presence of unpaired electrons. When exposed to an external magnetic field, the magnetic moments of the unpaired electrons align with the field, resulting in a net magnetic moment in the direction of the applied field.
• Hund's rule: all orbitals in sublevel are singly occupied before any orbital is doubly occupied AND all electrons in singly occupied orbitals have the same spin = maximized total spin

  
• Solids


  • Spins of many atoms may align (ferromagnetic) or not align (antiferromagnetic)
• 1. orbitals


  Spherical orbitals with the nucleus at its center

  
• 1. orbitals


  Dumbbell shaped orbitals, 3 different orientations… 2 orbital lobes and one orbital node in the nucleus