Download But How is This Related to the Atom?

But How is Quantized Light
Related to the Atom?
Atomic Spectroscopy and the
Bohr Model


Discovery of particle nature of light began to
break down the division that existed in 19thcentury physics between EM radiation (wave
phenomenon) and small particles
During this time, many scientists experimented
with atomic spectroscopy which is the study
of EM radiation absorbed and emitted by
atoms

There was a phenomenon called line spectra
that puzzled scientists
White Light Spectra
White
light is a blend of all visible wavelengths
They can be separated using a prism
Produces a continuous spectrum
Light and the Dilemma of
Atomic Spectral Lines
 Neils
Bohr, a former
student of Rutherford,
studied the spectra
produced when
atoms were excited in
a gas discharge tube
 He observed that
each element
produced its own set
of characteristic lines
called a line spectrum
instead of a
continuous spectrum
The Dilemma of Atomic Spectra


If electrons were randomly situated, as depicted in
Rutherford’s atomic model, then they would be
able to absorb and release energy of random
colors of light
However, as the electrons in hydrogen atoms were
getting excited and then releasing energy, only
four different color bands of visible light were being
emitted: red, bluish-green, and two violet-colored
lines
How is a Line Spectrum
Produced?

In order to produce a line spectrum,
atoms’ electrons must somehow absorb
energy and then give the energy off in
the form of light at a specific wavelength


What is the relationship between energy
and wavelength again?
Do you think we can we map the
electrons by using these energy
relationships from the emission spectrum?

The answer is YES!
Neils Bohr and the Atomic
Model

Neils Bohr was one of the first to see some
connection between the wavelengths an
element emits and its atomic structure

Related Planck’s idea of quantized energies
to Rutherford’s atomic model
A Summary of Bohr Model of the Atom
 The
Bohr model is a
‘planetary’ type model
 The
nucleus is at the
center of the model
 Electrons can only exist
at specific energy levels
(orbits)
 Each
energy level was
assigned a principal
quantum number, n
 Each principal quantum
represents a new ‘orbit’
or layer
Bohr’s Model of the Atom

Bohr suggested that electrons typically
have the lowest energy possible (ground
state), but upon absorbing energy via
heat or electricity:


Electrons jump to a higher energy level,
producing an excited and unstable state
Those electrons can’t stay away from the
nucleus in those high energy levels forever so
electrons would then fall back to a lower
energy level
Wait…Something is Amiss
If
electrons are going from
high-energy state to a lowenergy state, where is all this
extra energy going?
Connecting Energy to the
Atomic Model
 Energy

does not disappear
First Law of Thermodynamics!
 Electrons
re-emit the absorbed energy
as photons of light

Difference in energy would correspond
with a specific wavelength line in the
atomic emission spectrum

Larger the transition the electron makes, the
higher the energy the photon will have
More on Hydrogen Spectral
Lines
 Transitions
to the ground-state (nf = 1) give rise to
spectral lines in the UV region of EM spectrum
 Set of lines is called the Lyman series
 Transitions to the first excited state (nf = 2) give rise
to spectral lines in the visible region of EM
spectrum
 Set of lines is called the Balmer series
 Transitions to the second excited state (nf = 3)
give rise to spectral lines in the IR region of EM
spectrum
 Set of lines is called the Paschen series
Many Electron Atoms
 Recall
that because each element has
different number of electrons and a
slightly different structure, the colors that
are given off by each element are going
to be different

Thus, each element is going to have its own
distinct color when its electrons are excited
(or its own atomic spectra)
hydrogen (H)
mercury (Hg)
neon (Ne)
Shortcomings of the Bohr
Model

Bohr’s model was too simple




Worked well with only hydrogen because H only has one
electron
Could only approximate spectra of other elements with
more than one electron
Bohr also avoided the problem of why the negativelycharged electron would not just fall into the positively
charged nucleus, by simply assuming it does not
happen
Furthermore, there is a problem with describing an
electron merely as a small particle moving in circular
orbits around the nucleus
So there is more to the atomic puzzle…