Download Chapter 7 NotesAA

ATOMIC STRUCTURE

Characterize wave forms of energy in terms of frequency, speed,
wavelength, and amplitude. Understand the mathematical
relationship between frequency, wavelength, and speed of an
electromagnetic wave.

Understand and describe the Bohr model of the atom including
restricted energy levels for electrons and what happens when
electrons transition from one energy level to another.

Gain a basic understanding of the quantum model of the atom and the
significance of the term orbital in describing the behavior of an
electron.

Be familiar with the major contributors to the quantum model of the
atom including Bohr, Heisenberg, Einstein, Planck, deBroglie, and
Schroedinger.

Describe the allowed energy states of electrons using the system of
quantum numbers.
 #1

 #2

– 21, 23, 25, 27, 29, 67, 69
Electromagnetic Radiation, Spectrum, Energies
–47, 49, 51, 53, 55, 57, 59, 61, 63, 73, 87,
Quantum Mechanics and Quantum #’s
Introduction
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Flaw in Rutherford’s Model
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RECALL: Rutherford and Bohr’s Contributions and Models
of the Atom!
FLAW: Electrons are particles that travel in orbits
Problem for 20th century scientists

Model didn’t work for larger atoms, only for small ones
like H2

The nature of light – a side trip
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Light is not matter…………or is it?!?
Wave-Particle Duality?
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Light has properties that are both wave-like and particlelike
Particle-like: packets of energy called photons
 2.
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Electromagnetic Radiation – definition
Energy in the form of a wave resulting from
oscillating electric and magnetic fields
ALL forms of light, both visible and invisible to the
human eye
Lowest to highest energy
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4. Properties of waves
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Wavelength (): distance between any two points in a
wave
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Frequency (): the number of cycles of the wave that
pass through a point in a unit of time

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measured in nm (1*10-9 m)
Measured in s-1 (per seconds) or Hz (hertz)
Speed of light (c): 3.00 x 108 m/s

Baseline: midpoint between peak and trough (0 slope)

Amplitude: height of a wave
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Node: point on wave where amplitude = 0

Standing Wave: wave that has 2 or more points of no
movement (confined to strict limits)

Ex. Guitar String
5. Visible spectrum
Continuous spectrum: components of white light split
into its colors, ROY G BIV

from 400 nm (violet) to 760 nm (red)
 6.
Light equation and units
c = 
UNITS: m/s
a)
b)
Longer wavelength, lower frequency
Shorter wavelength, higher frequency
Wavelength is indirectly proportional to frequency
 Ex7.1
What is the frequency of visible light with a
wavelength of 625nm?

Types of radiation from lowest to highest energy
 1.

Heated objects
View in 1900 was two fold:

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Matter was particulate (having mass and position)
Energy was waves (having no mass or fixed position)
Then came Planck…
Planck studied the relationship between
temperature and the radiation emitted by an object.
Temperature  Energy emitted
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
2.
Planck’s hypothesis and equation
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As temperature increased molecular vibrations increased (frequency
increased)
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He noted that each time the temperature was increased the
frequency of the radiation emitted increased in a regular increment.
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Thus, he proposed that Energy was QUANTIZED.
E = h
h= 6.626 x 10-34 J s
“Continuous”
“Quantized”

3. Ex7.2 Green light has a wavelength of 500nm. What is the energy
in joules of one photon of green light? What is the energy of one mole
of photons of this green light?

4.
Photoelectric effect – Nobel Prize in Physics 1921 to Albert
Einstein

a. Description

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e- are ejected from the surface of a metal once a threshold frequency
is passed
b. Practical Uses

Solar Cells
 e- are promoted from valence
band to the conduction band
in a semiconductor
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c. Conclusion

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Ephoton=hν=hc/λ
E=mc2 or m= E/c2

1.
2.
Does this mean photons have mass?

Not in the same way as other particles but they do have momentum
Energy is quantized. Only exists in quanta or photons.
Electromagnetic radiation exhibits both wave and particulate
properties.
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1.
Continuous vs discontinuous spectra
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Continuous: shows all possible wavelengths (energies)
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Discontinuous: omits 1 or more possible wavelengths
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2.
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Bright line, or atomic emission spectra
Light is broken down into component pieces by a diffraction
grating
Only SPECIFIC ‘s of light are shown for an element
Atomic Fingerprint (distinct for each element)
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3.
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Bohr Model
e- rotate around nucleus in specific orbits
Energy is QUANTIZED (distinct energy levels)
Put in E  e- goes to a higher orbit
Falls back down to lower orbit  gives off E put in
 4.
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Minimum amount of energy (not zero); closest to nucleus
 5.
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Electrons in the excited state
Higher energy (further from nucleus)
 6.
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Energy of electrons in the ground state
Atomic emission spectra caused by
Electrons coming from excited to ground state which
gives off a photon of light (h)
 7.
Energy of electrons is quantized
Explains why only certain wavelengths of light are
shown for elements
 Don’t see a continuous spectrum for hydrogen because
there are not orbits for every possible wavelength

 8.
Flaw in the Bohr Model
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Didn’t work for elements more complex than hydrogen
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Bottom Line: Bohr’s model FAILED!

Fundamental Physics--A charged body (ie. an electron) following a
circular orbit will emit energy and as energy decreases so does the
distance to the nucleus. Thus as the electron orbits it will get
closer and closer to the nucleus and the atom will eventually
collapse.
 1.

Wave Particle Duality – Louis deBroglie
De Broglie wondered if waves exhibit particulate
properties, then do particles exhibit wave properties?
 2.
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Equation and “matter waves”
DeBroglie related mass (m), velocity (v), and wavelength ()
of matter
Most matter has undetectable wavelengths
 Led to the development of the electron microscope

 3.
Basic idea of Quantum Theory
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All matter exhibits wave and particle properties.
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For very small pieces of matter (photons) wave
properties dominate as they are very small particles.
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For large pieces of matter (baseballs) particulate
properties dominate as associated wavelengths are so
small they are basically not observed.
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For intermediate pieces of matter (electrons) a truly
dual nature is displayed.

1.

Heisenberg Uncertainty Principle
There is a fundamental limitation to just how precisely we can
know both the position and momentum of a particle at a given
time. In layman’s terms: The more precisely we know a particles
position the less we know about its momentum and vice versa.
 2.
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Schroedinger Equation’s major conclusion
A precise location for an electron cannot be determined
however regions which describe the probable location for
electrons can be and these regions are known as
orbitals.
 a.
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Mathematical function that describes the location of the
electron at a given moment in time
 b.
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wave functions ()
electron density
Shows most probable regions to find e-
 c.
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3-D shape of the e- probability region (>90% e- density)
2 = probability region
 d.
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orbitals
solutions to equation
Mathematical function that satisfies the equation
Can have multiple solutions = multiple orbitals
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Quantum Numbers: mathematical way to represent the
location of an electron in an atom
Analogy...
State = energy level, n
City = sublevel, l (“L”)
Address = orbital, ml
House Number = spin, ms
 e.
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Principal Quantum Number
“n”
Positive integer
Analogous to Bohr’s energy levels (and periods on the
Periodic Table)
Larger n value, larger atomic size
 f.
Angular Momentum Quantum Number
“l” (small-case L)
 Positive integer
 within an energy level=sublevel
 l = 0 = s, spherical shape
 l = 1 = p, dumbbell shape
 l = 2 = d, clover shape
 l = 3 = f, flower shape
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n=2 has 2 l values (0 and 1 or s and p)
n=4 has 4 l values (0,1,2,3 or s,p,d,f)
 g.
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Magnetic Quantum Number
“ml ”
Integer from - l to +l
orbital
l =0 (s): ml =0; 1orbital
l =1 (p): ml =-1,0,1; 3 orbitals
l =2 (d): ml =-2,-1,0,1,2; 5 orbitals
h. Spin Quantum Number
“ms ”
 +½ or -½
 2 electrons allowed in each orbital
 “Spins” of each electron must be opposite
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 3.
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Correspond to solutions to the Schroedinger equation

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Allowable values for quantum numbers
Cannot find an e- where there is no electron density
Must determine whether or not a set of quantum
numbers is plausible
 4.
Array of quantum numbers
Determine which of the following are Plausible:
n=3
l=2
ml = -1
ms = +½
n=1
l=1
ml = -1
ms = +½
n=5
l=0
ml = -1
ms = -½
n=2
l=1
ml = 0
ms = 0
n=3
l=2
ml = -2
ms = -½
n=4
l=3
ml = 2
ms = -½

See
 1.
 2.
 3.
 4.

back of your Packet!
s orbital
p orbital
d orbital
f orbital