Lecture 5 - Chemistry at Winthrop University

Lecture 5: Light and the EMR Spectrum
(Ch 4.4-4.9)
Dr Harris
Suggested HW: (Ch 4) 13, 19, 20, 21, 22, 28, 29, 31, 37
• Our present understanding of the electronic structure of atoms has come
from the light that is absorbed and emitted by substances
• For example, what happens when you switch on a
• Neon lights are glass chambers pressurized with Neon or other noble
• When a voltage is applied, the gas is ionized. This ionization causes a
• Why does this phenomena occur?
• What further information can this provide us about the electronic
structure of atoms?
• Over the course of the next two class periods, we will be able to
understand exactly what’s happening, and what it means
EMR: Light and Energy
• The applied voltage cause the electrons to become excited, or “bumped
up” in energy
• When the electron drops back down to its original, lower energy state, the
excess energy is released as light. This is called emission.
• But what exactly is light?
• The light that we see with our eyes is a type of electromagnetic
radiation (EMR)
• When we use the term radiation, we are referring to energy that is
propagated (moves and spreads outward) through space as waves
• Light, like that which emits from a lamp, is comprised of visible waves
• Radio waves from a radio are another type of EMR
• Invisible UV and Infrared rays from the sun are also EMR
Propagation of Waves
• The waves created in water when
an external force is applied are an
example of propagation.
• The energy transferred to the spot
of impact is spread and
transmitted throughout the water.
• EMR propagates through the
universe as oscillating,
perpendicular electric and
magnetic fields
Wavelength and Frequency
• The distance between local
maxima, or crests, is the
wavelength λ (units of
• If we picture these waves
moving across the page, the
number of crests that pass a
given point per second is the
frequency, ν (units of s-1)
• The speed of a wave is given
by the product of ν and λ:
λν = c
c is the speed of light, 3.0 x 108 m/s.
All EMR moves at this speed
through vacuum
Different Types of EMR Have Different
• The electromagnetic spectrum below shows EMR listed by increasing
• Wavelengths vary from the size of an atomic nucleus to the length of
a football field
The Visible Spectrum
(increasing Energy)
• What is the frequency of orange (~650 nm) light?
• A certain type of radiation has a frequency of 1015 s-1. What is the
wavelength, in nm, of this radiation? What kind of radiation is it?
Continuous Spectra
• White light is comprised of all wavelengths of the visible spectrum.
Because the spectrum of white light has no gaps, it is a continuous
• Sunlight, for example, is
continuous over a long range of
wavelengths. The spectrum of
sunlight is shown.
Line (Discontinuous) Spectra
• Light emitted from chemical samples exhibits a discontinuous spectrum.
The radiation consists of spectral lines at particular wavelengths. This type
of spectrum is a line spectrum, or atomic emission spectra
• Sodium burns very brightly and emits an orangishyellow color: Discontinuous spectrum
Max Planck
• The observation of spectral lines indicates that certain elements can
only emit certain wavelengths
• How can this be? Why can’t any element emit at any wavelength?
• Max Planck first began to answer this question with his interpretation
of a phenomena known as blackbody radiation.
Blackbody Radiation And The End Of Classical
• All solid objects, when heated, emit thermal
• Just when an object is hot enough to glow, it
appears red. As you continue to heat the
material, it becomes “white hot”
• Classical physics predicts that
continuous heating would produce
higher and higher frequencies at
increasing intensity
• This means that light bulbs would
give off UV, gamma, X-rays, and so
on. Of course, this doesn’t happen
The Birth Of Quantum Physics
• The failure of Classical Physics to explain blackbody radiation lead to
the creation of Quantum Physics by Planck, Einstein, and others.
• Planck explained blackbody radiation by asserting that light
(radiation) can only be emitted in small, exact amounts called
photons (or quanta)
• He then derived the amount of energy absorbed or released in a
single event is equal to:
En = nhν
where En is the total energy in J, n is the number of photons, and h
is Planck’s constant, 6.626 x 10-34 J•s
• Calculate the energy contained in a single photon of blue light (~400 nm)
• Calculate the energy of contained in 10 photons of green light (~520 nm)
Einstein and the Photon
• Like Planck, Einstein envisioned light as a beam of particles.
• Borrowing from Planck’s theory, he asserted that each photon in the
beam is a little packet of energy E = hν
• Using this theory, Einstein sought to
understand a phenomena that had defied
physics for many years prior… the Photoelectric
The Photoelectric Effect
• The photoelectric effect is the ejection of
electrons from metal surface under illumination
following the absorption of a photon’s energy.
• Photons too low in frequency (energy), no
matter how intense the beam, will not eject an
electron from a metal surface.
• It is not until some minimum frequency
(threshold frequency, νT) is reached, that a
photon is just energetic enough to loosen an
• At energies beyond the threshold energy (ET =
hvT), the electron converts the excess energy of
the photon into kinetic energy and is ejected.
Excess Energy is Converted to Kinetic Energy
• The energy of motion is called kinetic energy
• The kinetic energy of a body of mass is given
Plot of Ek vs. ν for sodium
slope of line = h

• m is the mass in kg, and  is the velocity
(speed) in meters per second (m/s). The
units of energy are Joules (J).
5.51 x 1014 s-1
• Einstein found that as you increase the energy of the incident photon
striking, the velocity of the ejected electron increases proportionally:
 =  − 
• Given that the threshold frequency of copper is 1.076 x 1015 s-1, calculate
the kinetic energy of an electron that will be ejected when a 210 nm
photon strikes the surface?
• What do we know?
νT = 1.076 x 1015 s-1
νphoton =

3.0  108   −1
2.100  10−7 
= 1.428  1015  −1
 =  − 
 = ℎℎ − ℎ = ℎ(ℎ−  )
 = (6.626  10−34 )(3.52  1014  −1 )
 = .   − 
Example Continued.
• From the example on the previous page, calculate the velocity of the
• Mass of electron = 9.109 x

  :  =
Joule =


2 (2.33  10−19  2  −2 )
= 7.15  105 /
(9.109  10 )
Section 2. Wave-Particle Duality
• Planck and Einstein were able to determine that energy transferred
to or from an electron must be quantized.
• However, the question yet to be answered is: What determines the
allowed energies of emission of a given element?
• The physical nature of photons and electrons needed to be
understood before this issue could be addressed
A New Wave of Thought
• Many years prior to Einstein’s photoelectric effect experiment, it had
been proposed that light was comprised of waves
• Thomas Young was the first physicist to propose that light was of
wave-like character, not particle like as proposed by Issac Newton
• To test his hypothesis, Young conducted the ‘slit experiment’
Light As Waves? Young’s Slit Experiment (1799)
• If light were made of only particles, then light passing through a slight of
height X would appear on a screen with the size and shape of the slit
• What Young observed, however, was a series of light and dark fringes
• This was the first indication of the wave-like character of light
Constructive and Destructive Interference
• The observed diffraction pattern of light can be explained by treating light
as waves with certain wavelengths and amplitudes.
• Waves of light that are in phase, can interact, forming a single wave of
larger amplitude. This is called constructive interference (a).
• Waves that are out of phase
will deconstruct (b), yielding
a lower amplitude
(destructive interference).
• Remember:
• wavelength determines
• amplitude dictates
Double Slit Experiment
• To confirm his hypothesis and prove his idea of constructive
interference, Young repeated the experiment using two slits.
• If light were indeed composed of waves, and the fringes due to
constructive interference, then the light fringes should be twice
as bright. The dark ones should be more defined. He was correct.
Young’s sketch of the interference, 1807.
Real Example
Back To the Photoelectric Effect
• In class yesterday, we described the photoelectric effect (Einstein, 1905)
• Electrons are bound to the metal atoms. The energy of this bond is the
threshold energy. In other words, it takes this much energy to ‘loosen’ the
• When photons strike a metal surface, one of three scenarios can occur:
1. The photon has an energy which is less than the threshold energy.
So, the photon is not absorbed and nothing happens.
The photon has EXACTLY enough energy to separate an electron
from the metal atom. However, there is no energy left for the
electron to move. Motion REQUIRES kinetic energy
The photon has EXCESS energy. The excess energy is converted to
kinetic energy, and the electron moves away at some velocity
(speed) v.
Ep = hνp
Electrons bound by energy E= hvT
Compton Scattering
• Einstein’s Photoelectric effect suggested that photons had momentum, a
property of particles
• Compton asserted… “If EMR is made of particles, lets hit something with
• This lead to the discovery of the ‘Compton Scattering’
• X-rays were found to ‘bounce’ off of
electrons at calculated angles, like pool
balls, and with an energy lower than
the initial energy
• This further supported particle-like
What now?
• Young’s slit experiments did not mean that Newton was wrong
about the particle nature of EMR
• Einstein’s and Compton’s work did not prove that Newton was
• What these experiments DID prove, was that physicists had to
develop a new theory that fused both the wave and particle-like
aspects of EMR into a single theory
DeBroglie’s Approach
• DeBroglie combined Einstein’s special theory of relativity with Planck’s
quantum theory to create the DeBroglie relation. In short, he summates
that if waves are particle-like, then particles, and hence, mass, are wavelike.
E = pc (p is momentum, p= m)
E = hν
pc = hν
Einstein (particle like):
Planck (wave like) :
DeBroglie (both) :
pc =
pλ = h
λD = h/p
• The value, λD is the DeBroglie wavelength,
or the wavelength of any mass m with
momentum p.
Louis DeBroglie (1892-1987)
DeBroglie’s Hypothesis Confirmed
• Below are diffraction patterns of Aluminum foil. The left image is formed
by bombarding Al atoms with X-rays. The right image is formed with an
electron beam.
• As shown, both the EMR and electrons behave in the same wave-like
Both exhibit
the wave-like
ability of
• Calculate the DeBroglie wavelength of an electron travelling at
1.00% of the speed of light.
λ = =

6.626  10−34  
9.109  10−31  .   −
= 2.43 x 10−10 
• What is the DeBroglie wavelength of a golf ball which weighs 45.9 g
and is traveling at a velocity of 120 miles per hour?
• First, convert velocity to meters per second
120  5280  .3048 

= 53.6 m/s

λ = =

6.626  10−34  
.0459  53.6  −1
= 2.69 x 10−34 
• DeBroglie wavelength of large objects is negligible
Quantum Condition
• Recall the Bohr model of the atom. Bohr used DeBroglie’s theory to
justify why electrons are restricted to certain orbits around the nucleus.
• As shown above, if the waves of the electron do not match after a
revolution, you will have progressive destructive interference, and the
waves will cancel.
• Thus, the orbits will only be stable if some whole
number of orbits, n, around the nucleus fit the
circumference (2πr) of the orbit.
Quantum Condition
• Therefore:
2 = λ
n = 1,2,3….
• We define n as the principle quantum number. Bohr showed
that an electron in a given orbit can ONLY have the following
n =3
n =2
n =1
• We say that the energy of
the electrons in each level is
• Each orbit represents an
allowed state, or energy level
in which an electron can
• The lowest energy state is called the ground state. When an
electron is transitioned to a higher state, the electron is said to be
excited, or in an excited state.
• Now, we can understand why certain elements can only emit at
certain wavelengths….
• because only certain transitions exist depending on the
circumference of the orbits around the nucleus
• Thus, when atoms absorb energy, electrons move to an excited
state. When they return to the ground state, the atom emits a
photon to release the energy. The energy of the photon is the
difference in energy between the initial and final states:
ℎ =  − 
• What would the wavelength of emitted light be, in nm, if an excited
hydrogen electron in the n=4 state relaxes back to the n=2 state?
ℎ =  − 
= 4 − 2

= .  
(6.626  10−34 )(3.0 108  −1 )
(.4088  10−18 )
= 486 
Atomic Emission Spectra of Hydrogen
There it is!!!
Transitions for a Hydrogen Atom
Emission in
the visible
• The work of Planck, Einstein, DeBroglie and Bohr has provided much
information into the relationship between EMR and electronic
• From the understanding that energies are quantized, and that
photons and electrons are both wave and particle like, the Bohr
model of the atom was able to explain the line spectra of hydrogen
• We now know that emission is the result of transitions from
quantized energy states. Different atoms have different allowed
• The allowed wavelengths of light that can be absorbed and emitted
by an atom give insight into the energy states involved in a given
process in an atom

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