Electronic Structure of Atoms

1
Chapter 6 Electronic Structure of Atoms
2
Electromagnetic Radiation
Low ?
High ?
3
WMHB broadcasts at 89.7 on the FM dial (89.7
MHz). What is the wavelength of the emitted
radio waves?
4
The AM dial ranges from 500-1800 kHz. What
wavelengths does this represent?
A microwave oven (5 X 109 Hz)?
5
Electromagnetic Spectrum
6
Constructive and Destructive Interference
7
Constructive and Destructive Interference
8
Refraction of Light
When white light is directed through a prism, a
continuous spectrum results (due to a change in
the speed of light through a different medium)
9
Atomic Spectra
Classical physics predicts that energetically
excited atoms will emit white light. But
experiment shows otherwise
Potassium
Hydrogen
Helium
Lithium
Sodium
Physicists could not explain why different colors
were observed.
10
Atomic Spectra
11
The Hydrogen Atomic Spectrum
The four visible lines in the atomic spectrum for
hydrogen were explained through an equation by
Johann Balmer in 1885, but the reason behind the
atomic spectra were still unknown. This was later
extended by Rydberg to
RH 1.096776 X 107 m-1 n1 and n2 are integers
and n1ltn2
12
Quantum Theory
If classical physics failed to explain these
phenomena, then a new theory was needed!
13
The Photoelectric Effect

• Heinrich Hertz (1888) discovered that light
  striking the surface of certain metals causes
  ejection of electrons - the photoelectric effect.
• Electrons were only ejected if the frequency of
  the light (? ) was above a certain threshold
  frequency (?o )
• The number of electrons ejected were proportional
  to the intensity of the light
• The velocity (kinetic energy) of the electrons
  ejected was proportional to the frequency of the
  light

14
Max Planck and Discontinuous Energy
Max Planck (1900) theorized that energy is
discontinuous according to the equation
E h?
h 6.626 X 10-34 J-s
If correct, this would partially explain the
observed atomic line spectra and the
photoelectric effect.
15
Planck, Einstein, and Quantum Theory
Albert Einstein extended Plancks theory and
asserted that particle behavior of light is
consistent with the photoelectric effect.
E h?
h 6.626 X 10-34 J-s
Thus, energy from electromagnetic radiation can
only be transmitted in discrete packets or
quanta. A single packet of electromagnetic
radiation is known as a photon - and a photon can
be thought of as a single particle of light with
a specific energy.
16
The energy required to break a carbon-carbon bond
is approximately 360 KJ/mol. What wavelength of
electro-magnetic radiation can supply this amount
of energy?
17
The Bohr Atom

• Niels Bohr hypothesized (1914) that
• Electrons move in circular orbits around the
  nucleus - called orbitals.
• Electrons can only exist in a discrete set of
  allowed orbitals - electrons do not lose energy
  when in these orbitals.
• Only discrete amounts (quanta) of energy can be
  absorbed to change the orbital of an electron.

18
Hydrogen Energy Levels
19
Emission and Absorption Spectroscopy
20
A New Quantum Mechanics
Diffraction of X-rays by a metal foil
Diffraction of electrons by a metal foil
21
Matter-Waves
Louis de Broglie, 1924 (Nobel Prize 1929) All
particles of matter exhibit wavelike properties
according to
22
The Heisenberg Uncertainty Principle
h
Dx Dp
4p
23
Standing Waves - An Introduction
In a standing wave, the nodes (points of zero
intensity) remain at a constant position.
l n 1, 2, 3
24
Standing Waves
Shrödinger suggested that electrons can be
described as standing waves orbiting the nucleus,
and used a mathematical equation to represent
these waves.
25
The Schrödinger Equation Wave Functions
Y (psi), the wave function. Corresponds to a
standing wave within the boundary of the system
being described. The energy of the standing wave
increases as we add more nodes.
26
Probability of Finding an Electron
27
Quantum Numbers
By using three different quantum numbers, we can
describe the wavefunction of any electron in an
atom. n the principle quantum
number Designates electronic shell (distance
from nucleus), n 1, 2, 3 It also
corresponds to the number of radial nodes l
The angular momentum quantum number Designates
the shape of the orbital, l 0, 1, 2(n-1) ml
The magnetic quantum number Describes the
orientation of the orbital ml - l -2, -1,
0, 1, 2 l
28
l 0, the s orbitals
29
l 1, the p orbitals
30
ml for the p orbitals
31
l 2, the d orbitals
32
Orbital Occupancy
Each orbital (described by 3 quantum numbers) can
contain two electrons. No two electrons can have
the same set of quantum numbers. Therefore
33
The Fourth Quantum Number, s
Electrons have a spin of either 1/2 or 1/2.
Electrons that occupy the same orbital must have
opposite spins.
34
Orbital Energies For Single Electron Systems
35
Penetration and Shielding
36
Orbital Energies
37
Electron Configurations
We will use several rules, plus our knowledge
of electron orbitals to identify the electronic
configuration of any atom

• The Aufbau process (building up)
• Electrons will always populate the lowest energy
  state available.

38
Electron Configurations
The Pauli exclusion principle No two electrons
can have all four quantum numbers alike. Hunds
rule Populate all degenerate orbitals (orbitals
of the same energy) with a single electron before
putting a second electron in any orbital. There
are several exceptions to this rule.
39
Orbital Filling
Because of electron shielding, orbital energies
are usually ordered as shown above
40
The Electronic Configuration of Carbon
41
The Electronic Configuration of N Ne
42
Filling the d Orbitals
43
(No Transcript)
44
Electron Configurations and the Periodic Table