Atomic structure
• Our current picture of the electronic
structure of atoms developed from early
studies of atomic spectra
• A quantum mechanical picture of atoms is
essential if chemical properties are to be
understood
Atomic spectroscopy
Atomic spectra
• Atoms and molecules that have been
“excited” by the addition of energy emit
radiation with characteristic wavelengths,
frequencies and hence energies
– E = h n and c = n l
• This suggests that the electrons in the atoms
are undergoing transitions between states
with well defined energies
– electron energy is quantized
The Bohr model
• Bohr proposed that the electrons in atoms
are confined to distinct energy levels with
well defined angular momentum
– each different energy level had a different
quantum number
• This explains the spectra of simple atoms
– However, it does not work for multielectron
atoms
The Schrodinger equation
• Spectra from multielectron atoms could be
predicted using a wave equation developed by
Schrodinger
– d2y /dx2 + d2y /dy2 + d2y /dz2 + 8p 2m(EV)y /h2 = 0
• The solutions to this equation describe the
allowed states of the atom or molecule
• y 2 is proportional to the probability of
finding an electron at a particular location
Solutions of the equation
• There are an infinite number of solutions to
the equation
– many of the solutions have different energies
– for one electron atoms or ions solutions with
different energies have different principle
quantum numbers n
– the orbital angular momentum of an electron is
determined by the quantum number l
• the component of orbital angular momentum in a
particular direction is determined by ml
Electron spin
• Experiments passing beams of atoms through
magnetic fields suggest that electrons have a
magnetic moment that can be either up or
down relative to an applied field
– the orientation of the magnetic moment is
specified by a spin quantum number ms
• Electron spin is needed to explain spectra
• ms does not naturally come out of the
Schrodinger equation
– does come out of Dirac’s relativistic treatment
Quantum numbers
• Four quantum numbers n, l, ml, ms are
needed to fully specify the state of a single
electron in atom
– n may take integer values 1 -> infinity
– l may take integer values n -1 to 0
– ml may take integer values between -l and +l
– ms may be +1/2 or -1/2
• Not all combinations of the four are possible
Allowed combinations
Orbital designation and the
number of orbitals
s orbitals
p orbitals
d orbitals
Polyelectronic atoms and the
Aufbau principle
• Orbitals are filled according to their relative
energies. Lowest energy first.
• Only two electrons can go in each orbital
• The relative energies of the orbitals is not
fixed they can change giving rise to some
confusion
Relative orbital energies for
polyelectronic atoms
The periodic chart
Violations of the general trend
• Cr and Cu are the two most notable
violations of the Aufbau principle
– Cr 3d54s1
– Cu 3d104s1
• These configurations are adopted to
minimize electron-electron repulsion
Electron configurations of TM ions
• Electrons are not always removed in the
same order that they go in
• All first row TM ions M2+ have a 3dn4s0
electron configuration
– relative orbital energies can vary with atomic
number and oxidation state
Magnetism
• Atom, ions and molecules with unpaired
electrons have a magnetic moment
• Such materials are said to paramagnetic and
are attracted into a magnetic field
– the more unpaired electrons there are the
stronger the attraction
• Compounds with only paired electrons are
diamagnetic