Spectroscopy and Electron
Configurations
Light is an electromagnetic wave*.
Wave Characteristics
Frequency (ν): number of waves that pass a point in a given period of time
Total energy is proportional to amplitude and frequency of waves.
Because speed of light (c) is a constant (3 x 108 m/s), wavelength and
frequency of electromagnetic waves are inversely proportional:
E = hν; c = νλ
Color
White light is composed of all colors
which can be separated into a rainbow,
or a spectrum, by passing the light
through a prism.
Each color light has a different
wavelength, and, therefore, frequency.
Amplitude & Wavelength
5
Are there other “Colors”?
The Electromagnetic Spectrum
• Visible light comprises only a small fraction of all
the wavelengths of light – called the
electromagnetic spectrum.
• Short wavelength (high frequency) light has high
energy.
- Gamma ray light has the highest energy.
• Long wavelength (low frequency) light has low
energy.
- Radiowave light has the lowest energy.
Electromagnetic Spectrum
Interactions of light and matter
Emission
Absorption
Transmission
Reflection or Scattering
What types of light spectra can
we observe?
A hot opaque body produces a continuous spectrum, a
complete rainbow of colors without any specific spectral lines.
A hot, transparent gas produces an emission line spectrum
a series of bright spectral lines against a dark background.
A cool, transparent gas in front of a source of a continuous
spectrum produces an absorption line spectrum - a series
of dark spectral lines among the colors of the continuous
spectrum.
Each chemical element produces its
own unique set of spectral lines.
Oxygen spectrum
Neon spectrum
Emission and absorption spectra
are inversely related.
Spectra of Mercury
Identifying Elements with
Flame Tests
Na
K
Li
Ba
Exciting Gas Atoms to Emit Light
with Electrical Energy
Hg
He
H
Analyzing the Hydrogen Emission
Spectrum
Rydberg found the spectrum of hydrogen could be described with an
equation that involved an inverse square of integers.
Bohr Model of Hydrogen Atom
e-
In the Bohr model, electrons:
- have quantized energies.
- have orbits a fixed distance from
the nucleus.
The Bohr Model
Interference: When Waves Interact
Diffraction
2-Slit Interference
Tro, Chemistry: A Molecular Approach
22
Electron Diffraction
If electrons behave like
particles, there should only be
two bright spots on the target.
Electron Diffraction
However, electrons actually present an
interference pattern, demonstrating
they behave like waves.
The Bohr Model
• Integer number of de Broglie wavelengths
must fit in the circumference of orbit.
Electron Transitions
• To transition to a higher
energy state, the electron
must absorb energy equal
to the energy difference
between the final and initial
states.
• Electrons in high energy
states are unstable. They
will transition to lower
energy states and emit light.
26
Principal Energy Levels in Hydrogen
The wavelengths of lines in the
emission spectrum of hydrogen
can be predicted by calculating
the difference in energy
between any two states.
Bohr Model of H Atoms
Hydrogen Energy Transitions
For an electron in energy state
n, there are (n – 1) energy
states to which it can transition.
Therefore, it can generate
(n – 1) lines.
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Chemical Fingerprints
• Every atom, ion, and molecule has a
unique spectral “fingerprint.”
• We can identify the chemicals in gas by
their fingerprints in the spectrum.
• With additional physics, we can figure out
abundances of the chemicals, and much
more.
Other spectroscopy
• Many spectroscopic techniques rely on
these electronic transitions used with
different sources of light.
• Energy can also be absorbed and emitted
in other “modes” including vibration and
rotation, leading to other types of spectra.
OH
Uncertainty Principle
h
x  v 
4
1
 
m
• Heisenberg stated that the product of the
uncertainties in both the position and speed of a
particle was inversely proportional to its mass.
– x = position, Δx = uncertainty in position
– v = velocity, Δv = uncertainty in velocity
– m = mass
• The more accurately you know the position of a
small particle, like an electron, the less you know
about its speed.
– and vice-versa
Wave Function, y
• Calculations show that the size, shape and
orientation in space of an orbital are determined
by three integer terms in the wave function.
– added to quantize the energy of the electron
• These integers are called quantum numbers.
–
–
–
–
principal quantum number, n
angular momentum quantum number, l
magnetic quantum number, ml
spin quantum number, ms
Principal Quantum Number, n
• characterizes the energy of the electron in a particular orbital
– corresponds to Bohr’s energy level
• n can be any integer.
• The larger the value of n, the more energy the orbital has.
• Energies are defined as being negative.
– An electron would have E = 0 when it just escapes the atom.
• The larger the value of n, the larger the orbital.
• As n gets larger, the amount of energy between orbitals gets
smaller.
En = -2.18 x 10-18 J 1
n2
for an electron in H
l = 0, the s orbital
• Each principal energy state has
1 s orbital.
• lowest energy orbital in a
principal energy state
• spherical
• number of nodes = (n – 1)
p orbitals
d orbitals
f orbitals