Chemistry 115

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Chemistry 115 Lecture 17 Outline Chapter
7 Wave behavior of matter The Uncertainty
Principle Intro to Quantum Mechanics Recitati
on The Bohr atom versus the Quantum
Mechanical model
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Quantum Theory and the Electronic Structure of
Atoms
Think of atoms as the building blocks of
molecules
The electrons associated with each atomic nucleus
form an orbital structure that leads to chemical
bonds
In Chapter 7, we will learn where that orbital
structure comes from
Quantum Mechanics
Quantum mechanical nature of electrons
Developed in 1920s -1930s by Einstein,
Schrödinger, Heisenberg, Bohr others. QM
gives a much deeper thorough explanation of
chemistry, bonding, and periodic properties. We
will cover this subject by first asking ourselves
the following questions, which we will then
answer over the next 2-3 lectures
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• 7 Questions on Atomic Structure (Ch 7)
• 1) Why does light interfere with itself, causing
  constructive destructive interferences?
• Why can some light excite an electron from a
  metal surface some cannot?
• Why does the light from gases consist of discrete
  lines at certain colors or wavelengths?
• Why can we not describe the electrons motion
  around the nucleus, such as we can for planets?
• Why do different atoms absorb or emit light of
  different energies or wavelengths?
• Why are some atoms so reactive, and others
  unreactive?
• How does quantum mechanics help explain the
  structure of the periodic chart?

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The Wave Nature of Light (Q1) Light is a form
of electromagnetic radiation. Radiation carries
energy through space. Electromagnetic radiation
is characterized by its wave nature. All waves
have a characteristic wavelength, ?, and
amplitude A. The frequency, n, is a wave is the
number of cycles which pass a point in one
second. The units of n are Hertz (1 Hz 1
s-1). The speed of a wave is given by its
frequency multiplied by its wavelength.
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Travelling wave
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Shorter wavelength Higher Frequency
3.0 x 108 m s-1
Amplitude Brightness
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Frequency and wavelength
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Long wavelengths yield low frequencies
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The electromagnetic spectrum
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Two properties of light
Shown in
Property
Two-slit interference, refraction
Wave-like behavior
Particle-like behavior
Wavepacket Localization
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Wave-like property of light shown in two slit
interference
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Refraction Wave-like behavior
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Different behaviors of waves and particles
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Light Particles (Photons) and Quantized energy
(Q2) Planck energy can only be absorbed or
released from atoms in certain amounts. These
amounts are called quanta. The relationship
between energy and frequency is E hn
where h is Plancks constant (6.626 ? 10-34
J-s).
To understand quantization consider the notes
produced by a violin (continuous) and a piano
(quantized) A violin can produce any note by
placing the fingers at an appropriate spot on the
bridge. A piano can only produce notes
corresponding to the keys on the keyboard.
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The Photoelectric Effect (Q2) The photoelectric
effect provides evidence for the particle nature
of light. It also provides evidence for
quantization. If light shines on the surface of
a metal, there is a point at which electrons are
ejected from the metal. The electrons will only
be ejected once the threshold frequency is
reached. Below the threshold frequency no
electrons are ejected. Above the threshold
frequency, the number of electrons ejected depend
on the intensity of the light. Einstein assumed
that light traveled in energy packets called
photons. The energy of one photon, E hn.
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Localized interactions show the photon concept
Light Absorption
Photoelectric effect
Localized, instantaneous interaction of light
with metal surface
Threshold frequency. Not a property of A or t
Number of e- proportional to A
Wavepacket or photon arising from interaction of
light with the metal surface (absorption) must
carry a specific amount of energy
That energy must be determined by the frequency
E hn
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Localization Particle-like behavior
Wavepacket
Constructive interference formed when traveling
waves of different l overlap in space
(at a point of light absorption or emission)
Photon
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The wave - particle duality of light
Light has wave properties
refraction, diffraction
continuous
Light has particle properties
photons in the photoelectric effectand blackbody
radiation
discrete in time and spacepackets of energy
We will see a similar wave - particle duality for
matter
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Bohr model for the hydrogen atom H e-
(Q3)
We cannot see the way the electron binds to the
proton
But we can see light emitted from a gas of
hydrogen atoms
Bohr was the first to connect the pattern of
colors to a model
(like the concept necessary to explain the
emission of a blackbody)
Lines (emission of light) associated with H atom
energy levels
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Bohrs Model Rutherford assumed the electrons
orbited the nucleus analogous to planets around
the sun. Bohr noted the line spectra of certain
elements and assumed the electrons were confined
to specific energy states. These were called
orbits. Colors from excited gases arise because
electrons move between energy states in the
atom. Since the energy states are quantized, the
light emitted from excited atoms must be
quantized and appear as line spectra.
Bohr showed that
where n is the principal quantum number (i.e. n
1,2,3, and nothing else), and RH is the Rydberg
constant 2.18 ? 10-18 J 13.8 eV electron
volt.
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The first orbit in the Bohr model has n 1 and
is closest to the nucleus. The furthest orbit in
the Bohr model has n ? ? and corresponds to E ?
0. Electrons in the Bohr model can only move
between orbits by absorbing and emitting energy
in quanta (E hv). The amount of energy absorbed
or emitted on moving between states is given by
DE RH
?E Ef El hv When ni gt nf energy is
emitted. When nf gt ni energy is absorbed.
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The Bohr model explanation of the three series of
spectral lines