Ch. 4 “Electron Configurations Quantum Mechanics Made Simple!
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- Ch. 4 Electron Configurations Quantum Mechanics Made
Simple!
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- In chapter 3, we began our historical journey though the
development of atomic theory. Rutherfords Nuclear Atom was more
useful than Daltons or Thomsons models because it was able to
explain the results of the alpha particle scattering experiment. As
more evidence was accumulated, it, too, was replaced by a better
model!
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- An atom consists of a nucleus nucleus (of protons and neutrons)
(of protons and neutrons) electrons in space outside the nucleus.
electrons in space outside the nucleus. What we know so far Nucleus
Electron cloud
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- Much of our understanding of how electrons behave in atoms
comes from studies of how light interacts with matter. As you know,
light travels through space & is a form of radiant energy. This
is what makes you feel warm as you stand in sunlight! How light
travels through space has been a major source of debate for
centuries!
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- 1600s, Isaac Newton suggested that light was made of tiny
particles. Newton used a glass prism to refract (bend) sunlight
(white light) into a continuous spectrum. Continuous Spectrum a
complete array of colors from red to violet. (a rainbow!) ROYGBIV
(or VIBGYOR) This process is called Diffraction passing white light
through a diffraction grating to produce a continuous
spectrum.
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- 1600s, Christian Huygens (Dutch) suggested that light consists
of waves (rather than particles) - Wave Model of Light He thought
light travels away from its source the way water waves travel away
from a stone dropped in a pond. This Wave Model of Light Survived
into the 1900s!
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- In the early 1900s scientists were still using cathode ray
tubes to study light ! When they passed electricity through gases,
the electrons in the gas atoms would absorb the extra energy. The
atom is then said to be excited! However, the electrons dont keep
this extra energy for long. They immediately give it back off in
the form of Electromagnetic Radiation Energy that travels through
space as waves.
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- 12 Light (E-M Radiation) All types travel at light speed (c)
3.00x10 8 m/s All types have wave characteristics (wavelength,
frequency) wavelength ( lambda) - distance between successive peaks
(m) Frequency ( nu) - # cycles passing a given point each second
(1/s or Hz One cycle (Frequency is # of cycles per second)
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- Electromagnetic Radiation covers a broad spectrum: Radio waves
~ 10 3 m Microwaves ~ 10 -3 m Infrared light ~ 10 -5 m Visible
light ~ 10 -6 m Ultraviolet light ~ 10 -8 m X-rays ~ 10 -10 m Gamma
rays ~ 10 -12 m red750 nm orange yellow green blue indigo Violet400
nm (decreasing (only long wave on list!) Types of Electromagnetic
Radiation
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- Link to FCC Radio Frequency Chart
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- Because all EM radiation moves at the same speed, wavelength
and frequency are inversely proportional: c What is the wavelength
of radiation whose frequency is 6.24 x l0 14 sec -1 ? c 3 x 10 8
m/s c 3 x 10 8 m/s 6.24 x 10 14 s = = = 4.81 x 10 -7 m Speed of
light! Is this visible light? If so, what color? 4.81 x 10-7m x 10
9 nm 1 m = 481 nmYES! Blue
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- 2. what is the frequency of radiation whose wavelength is 2.20
x l0 -6 nm? (1 m = 10 9 nm) 2.20 x l0 -6 nm x 1 m 10 9 nm 10 9 nm c
= 3 x 10 8 m/s 2.20 x l0 -15 m = 1.36 x 10 23 s -1 Is this visible
light? If so, what color? No! Gamma or cosmic radiation = 2.20 x 10
-15 m
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- Remember - When heat or electricity is passed through a gas,
the electrons in the gas atoms absorb the extra energy. The atom is
then said to be excited! But, the electrons dont keep this extra
energy for long. They immediately give it back off in the form of
Electromagnetic Radiation (visible light) One way to demonstrate
the emission of light from excited atoms is by using a Flame
Test.
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- Flame Tests strontiumsodiumlithiumpotassiumcopper Many elements
give off characteristic light which can be used to help identify
them.
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- Flame Tests You heat a metallic salt & it burns with a
colored flame! You heat a metallic salt & it burns with a
colored flame! This is the characteristic glow of the excited metal
ions! This is the characteristic glow of the excited metal
ions!
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- Fireworks
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- Copyright 2007 Pearson Benjamin Cummings. All rights
reserved.
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- Flame Emission Spectra methane gas wooden splintstrontium
ioncopper ionsodium ion calcium ion
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- Neon signs Bent up cathode ray tube! NOT!!!
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- The Electric Pickle Excited atoms can emit light. Excited atoms
can emit light. Here the solution in a pickle is excited
electrically. The Na + ions in the pickle juice give off light
characteristic of that element. Here the solution in a pickle is
excited electrically. The Na + ions in the pickle juice give off
light characteristic of that element.
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- Bright-Line Spectra Passing the light from excited atoms
through a prism does something different - Passing the light from
excited atoms through a prism does something different - The
spectrum contains lines of only a few colors or wavelengths.
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- Bright-Line Emission Spectrum ground state excited state ENERGY
IN PHOTON OUT 656 nm486 nm 410 nm 434 nm Wavelength (nm) Prism
Slits
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- Each element has a unique bright-line spectrum. Each element
has a unique bright-line spectrum. i.e. an elements fingerprint
Helium This is how we know what stars are made of!
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- Spectrum of White Light
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- Spectrum of Excited Hydrogen Gas
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- Emission Spectrum of Hydrogen 1 nm = 1 x 10 -9 m = a billionth
of a meter 410 nm434 nm486 nm656 nm 1 nm = 1 x 10-9 m = a billionth
of a meter
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- Continuous and Line Spectra 4000 A o 5000 6000 7000 light Na H
Ca Hg 400 450 500 550 600 650 700 750 nm Visible spectrum (nm)
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- At the beginning of the 20 th century the accepted theory of
light was still the wave model. (light & other forms of
electromagnetic radiation travel as waves) Scientists found that
only a certain minimum energy could excite atoms & get them to
emit light. So they knew that energy had to be related to the
fundamental properties of frequency & wavelength
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- The temperature of a Pahoehoe lava flow can be estimated by
observing its color. The result agrees well with the measured
temperatures of lava flows at about 1,000 to 1,200 C.Pahoehoe
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- 1900, Max Planck (Germany) accurately predicted how the
spectrum of radiation emitted by an object changes with its
temperature. Max Planck The color (wavelength) of light depends on
the temperature Red hot objects are cooler than white hot
objects
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- He named each small chunk of energy a quantum (meaning fixed
amount) A quantum is the smallest unit of energy Although small,
quanta are significant amounts of energy on the atomic level.
Planck suggested that the energy absorbed or emitted by an object
is restricted to pieces of particular size.
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- Planck said that the energy of a Planck said that the energy of
a light is directly proportional to light is directly proportional
to its frequency its frequency E = h E:energy (J, joules) h:Plancks
constant (6.6262 10 -34 Js) :frequency (Hz) :frequency (Hz) really
small! E = h c (so inversely proportional to wavelength!) c = so
c/
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- Long Wavelength = Low Frequency = Low ENERGY Short Wavelength =
High Frequency = High ENERGY Wavelength Table
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- Small wavelength Large frequency Large energy Large wavelength
Small frequency Small energy
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- Quantum Theory GIVEN: E = ? = 4.57 10 14 1/s h = 6.6262 10 -34
J s WORK: E = h E = ( 6.6262 10 -34 J s ) ( 4.57 10 14 1/s ) E =
3.03 10 -19 J Example: Find the energy of a red photon with a
frequency of 4.57 10 14 1/s. Example: Find the energy of a red
photon with a frequency of 4.57 10 14 1/s.
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- Examples: 1. If a certain light has 7.18 x l0J of energy, what
is the frequency of this light? 1. If a certain light has 7.18 x l0
-19 J of energy, what is the frequency of this light? A: 1.08X10 15
s -1 or Hz b. what is the wavelength of this light? A: 2.78X10 -7 m
2. If the frequency of a certain light is 3.8 x l0 14 Hz, what is
the energy of this light? A: 2.5X10 -19 J 3. The energy of a
certain light is 3.9 x l0 -19 J. What is the wavelength of this
light? Is it visible? A: 510 nm Yes visible light.
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- What if the energy of a car was quantized? The car would only
be able to move at certain speeds! Lets say a cars fundamental
quantum of energy was equal to 10 mph. If it had 7 quanta, how fast
would it be going? Yeppers! 70 mph If it had 3 quanta? 30 mph The
car can gain or lose energy only in multiples of its fundamental
quantum 10 No gradual acceleration or deceleration! It couldnt go
25 mph or 67 mph
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- So, why arent we aware of quantum effects in the world around
us? Remember the size of Plancks Constant? It is very small (10 -34
) To us, energy seem continuous because the quanta are too small to
be noticed. However, for atoms, which are also very small, quanta
are of tremendous significance!
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- Albert Einstein saw Plancks idea of quantized energy as a new
way to think about light. In 1905, Einstein used Plancks equation
to explain another puzzling phenomenon The Photoelectric
Effect.
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- The Photoelectric Effect refers to the emission of electrons
from a metal when light shines on the metal. The wave theory of
light (early 1900) could not explain this phenomenon. For a given
metal, no electrons were emitted if the lights frequency was below
a certain minimum regardless of how long the light was shone. Light
was known to be a form of energy, capable of knocking loose an
electron from a metal. But the wave theory of light predicted that
light of any frequency could supply enough energy to eject an
electron. Scientists couldnt explain why the light had to be of a
minimum FREQUENCY in order for the photoelectric effect to occur.
The wave theory of light (early 1900) could not explain this
phenomenon. For a given metal, no electrons were emitted if the
lights frequency was below a certain minimum regardless of how long
the light was shone. Light was known to be a form of energy,
capable of knocking loose an electron from a metal. But the wave
theory of light predicted that light of any frequency could supply
enough energy to eject an electron. Scientists couldnt explain why
the light had to be of a minimum FREQUENCY in order for the
photoelectric effect to occur.
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- Solar Calculator Solar Panel
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- Albert Einstein expanded on Plancks theory by explaining that
electromagnetic radiation has a dual wave-particle nature. While
light exhibits many wavelike properties, it can also be thought of
as a stream of particles. Albert Einstein expanded on Plancks
theory by explaining that electromagnetic radiation has a dual
wave-particle nature. While light exhibits many wavelike
properties, it can also be thought of as a stream of particles.
Each particle of light carries a quantum of energy directly
proportional to the frequency. Einstein called these particles
photons. A PHOTON is a packet of light carrying a quantum of
energy.
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- Photon Is light a wave or a particle? Macroscopically it
behaves as a wave! On the atomic level, we observe particle
properties! Seen on the door to a light-wave lab: "Do not look into
laser with remaining good eye."
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- Dual Nature of Light light exhibits wave properties &
particle properties
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- Einstein explained the photoelectric effect by proposing that
electromagnetic radiation is absorbed by matter only in whole
numbers of photons. In order for an electron to be ejected from a
metal surface, the electron must be struck by a single photon
possessing at least the minimum energy (E photon = hv) required to
knock the electron loose, this minimum energy corresponds to a
minimum frequency. If a photons frequency is below the minimum,
then the electron remains bound to the metal surface. Electrons in
different metals are bound more or less tightly, so different
metals require different minimum frequencies to exhibit the
photoelectric effect.
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- Photoelectric Effect No electrons are emitted Electrons are
emitted Metal plate Bright red light infrared rays or Dim blue
light ultraviolet rays or
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- 57 Quantized Energy and Photons Phenomena not explained by wave
nature of light: 1) Black-body radiation light coming from a heated
object (Planck) 2) Photoelectric effect electrons emitted from
light illuminated surface (Einstein) 3) Emission Spectra light from
electronically excited gas atoms emission spectrum (top),
absorption spectrum (bottom)
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- 1913 Niels Bohr studied under Rutherford at Victoria University
in Manchester. Bohr refined Rutherford's idea by adding that the
electrons were in orbits. Rather like planets orbiting the sun.
With each orbit only able to contain a set number of electrons.
Neils Bohr
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- Bohrs Model of Hydrogen Neils Bohr incorporated Plancks quantum
theory to explain bright-line spectra. Bohr said the absorptions
and emissions of light by hydrogen corresponded to energy changes
within the atom. The fact that only certain frequencies are
absorbed or emitted by an atom tells us that only certain energy
changes are possible in an atom. Niels Bohr (1885-1962) (1913)
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- Bohrs Bohrs Planetary Model of the Atom electrons exist only in
orbits with specific amounts of energy called energy levels
Therefore electrons can only gain or lose certain amounts of energy
(quanta) The orbit closest to The nucleus is the most stable &
lowest In energy.
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- Bohrs Planetary Model of the Atom Nucleus Electron Orbit Energy
Levels Niels Bohr &Albert Einstein
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- The lowest energy state of an atom is its. The lowest energy
state of an atom is its ground state. A state in which an atom has
a higher amount of energy is an excited state. When an excited atom
returns to its ground state, it gives off photons of energy
(light!)
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- e-e- e-e- Ground state Excited state Electrons can only be at
specific energy levels, NOT between levels. Electrons can jump to a
higher energy level when the atom absorbs energy. When the electron
drops back down to a lower level, it gives the extra energy off as
light. Electrons cant stop between energy levels so the jumps
involve definite amounts of energy. (amount of energy ~ light
color!)
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- An excited lithium atom emitting a photon of red light as it
drops to a lower energy state. Photon of red light emitted Li atom
in lower energy state Excited Li atom Energy
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- Electron Energy Levels nucleus 1 st energy level 2 nd energy
level 3 rd energy level Energy absorbed Energy lost
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- Bohr Model of Atom Increasing energy of orbits n = 1 n = 2 n =
3 A photon is emitted e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e-
e-e- e-e-
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- Bohr Model 1 2 3 4 5 6 Energy of photon depends on the
difference in energy levels Energy of photon depends on the
difference in energy levels Bohrs calculated energies matched the
bright-line spectrum for the H atom Bohrs calculated energies
matched the bright-line spectrum for the H atom nucleus
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- 70 Bohr Model Limitations Unfortunately, this model only works
for Hydrogen! The success of Bohrs model of the hydrogen atom is
explaining observed spectral lines led many scientist to conclude
that a similar model could be applied to all atoms. It was soon
recognized, however, that Bohrs approach did not explain the
spectra of atoms with more than one electron. Nor did Bohrs theory
explain the chemical behavior of atoms.
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- With more sophisticated equipment, spectral lines were found to
consist of closely spaced lines called Doublets Hydrogen (pretty
simple!) Helium (not so basic!) doublets So there had to be more to
Bohrs energy levels (orbits) than he realized. MORE SCIENTIFIC
ADVANCEMENTS!
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- De Broglies Hypothesis - De Broglies Hypothesis - Duality of
Matter Since waves have particle characteristics (Dual Nature of
Light) Since waves have particle characteristics (Dual Nature of
Light) Moving particles have wave characteristics Moving particles
have wave characteristics Louis de Broglie ~1924 In 1924, Louis
DeBroglie suggested that every moving particle has a wave nature
just like light!
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- The Wave-like Electron Louis deBroglie The electron propagates
through space as an energy wave. To understand the atom, one must
understand the behavior of electromagnetic waves.
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- Wave-Particle Duality JJ Thomson won the Nobel prize for
describing the electron as a particle. His son, George Thomson won
the Nobel prize for describing the wave-like nature of the
electron. The electron is a particle! The electron is an energy
wave!
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- According to Isaac Newton, we can determine both the position
& momentum of a large body. (like an airplane) However, we
CANNOT accurately predict where an electron will be at some future
time! Heisenberg Uncertainty Principle (1926) says that it is
impossible to know both the location and the momentum of an
electron simultaneously. Werner Heisenberg 1901-1976
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- 78 Heisenberg Uncertainty Principle You can find out where the
electron is, but not where it is going. OR You can find out where
the electron is going, but not where it is! One cannot
simultaneously determine both the position and momentum of an
electron. Werner Heisenberg
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- Microscope Electron In order to observe an electron, one would
need to hit it with photons having a very short wavelength. Short
wavelength photons would have a high frequency and a great deal of
energy. If one were to hit an electron, it would cause the motion
and the speed of the electron to change. Heisenberg Uncertainty
Principle
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- heck
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- In the Bohr Model of the atom, the electron is at a fixed
distance from the nucleus. He assumed we knew both the position
& the momentum of electrons. The Uncertainty Principle
disproves this!
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- 83 A New Model! (The Last One!) So De Broglie and Heisenbergs
contributions lead us to a new atomic model. It will recognize the
wave nature of the electron and describe it in terms appropriate to
waves. The resulting model will precisely describe the ENERGY of
the electron, while describing its location as a probability.
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- The Quantum Mechanical Model of the Atom Erwin Schrodinger
1887-1961
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- Schrdingers Quantum Mechanical Model Used to determine the
PROBABILITY of finding an electron at any given distance from the
nucleus Describes the electron as a 3-dimensional wave surrounding
the nucleus. (fan blades) (fan blades) Schrodinger applied
DeBroglies idea of electrons behaving as waves to the problem of
electrons in atoms.
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- Today we say that the Today we say that the electrons are
located in a region of space outside the nucleus called the.
electron cloud. Quantum Mechanical Model The Quantum Mechanical
Model of the atom describes the electronic structure of the atom as
the probability of finding electrons within certain regions of
space (orbitals). The Quantum Mechanical Model of the atom
describes the electronic structure of the atom as the probability
of finding electrons within certain regions of space (orbitals).
1926
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- Quantum Mechanics Orbital - densist, darkest region of the
electron cloud) Orbital - densist, darkest region of the electron
cloud) Region in space where there is a high (90%) probability of
finding an electron Region in space where there is a high (90%)
probability of finding an electron Electron Probability vs.
Distance Electron Probability (%) Distance from the Nucleus (pm)
100150200250500 0 10 20 30 40 90% probability of finding the
electron Electron cloud
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- LOHS AP Chemistry Fall 2007Dr. Schrempp89 Erwin Schrodinger
(1887-1961) Won Nobel Prize in 1933 for his equation. Came up with
a paradoxical thought experiment to show problems in observing
isolated systems (Schrodingers Cat) Experiment: A cat is placed in
a sealed box containing a device that has a 50% chance of killing
the cat.
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- 94 Unfortunately, Schrodingers cat could not cope with a life
of uncertainty
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- Modern View The atom is mostly empty space The atom is mostly
empty space Two regions Two regions Nucleus Nucleus protons and
neutrons protons and neutrons Electron cloud Electron cloud region
where you are likely to find an electron region where you are
likely to find an electron I don't like it, and I'm sorry I ever
had anything to do with it. - Erwin Schrodinger talking about
Quantum Physics
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- Feeling overwhelmed? Just a little more! "Teacher, may I be
excused? My brain is full." Chemistry
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- Quantum Numbers Four Quantum Numbers: Specify the address of
each electron in an atom 4 4 4 4 4 4 4 4 4 4 The QM model makes it
possible to describe the location of an electron Using four quantum
numbers.
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- Quantum Numbers Principal Quantum Number ( n ) Angular Momentum
Quantum # ( l ) Magnetic Quantum Number ( m ) Spin Quantum Number (
s )
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- Quantum Numbers 1. Principal Quantum Number ( n ) Energy level
Energy level Size of the orbital cloud Size of the orbital cloud n
= 1, 2, 3, 4 n = 1, 2, 3, 4 n 2 = # of orbitals in the energy level
n 2 = # of orbitals in the energy level 2n 2 = # of electrons per
energy level 2n 2 = # of electrons per energy level 1s1s 2s2s
3s3s
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- Electron Energy Level (Shell) Principle Quantum number
Generally symbolized by n, it denotes the probable distance of the
electron from the nucleus. n is also known as the Principle Quantum
number
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- Quantum Numbers s p d f 2. Angular Momentum Quantum # ( l )
Corresponds to: Energy sublevel Shape of the orbital l = s, p, d, f
(in order of increasing energy) s cloud is spherical, p cloud is
dumb-bell shaped
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- Sublevel names
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- Quantum Numbers 3. Magnetic Quantum Number ( m ) Orientation
(direction in space) of orbital Specifies the exact orbital within
each sublevel that the electron occupies.
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- Quantum Numbers pxpxpxpx pzpzpzpz pypypypy x y z x y z x y z A
p sublevel has 3 possible orbitals oriented along the x, y, & z
axes.
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- pxpx
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- pypy
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- pzpz
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- Copyright 2007 Pearson Benjamin Cummings. All rights
reserved.
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- d-orbitals A d sublevel has 5 possible orbital clouds
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- f Orbitals An f sublevel has 7 possible orbitals
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- Quantum Numbers 4. 4. Spin Quantum Number ( s ) Electron spin +
or - Electron spin + or - An orbital can hold 2 electrons that spin
in opposite directions clockwise or counter- clockwise. An orbital
can hold 2 electrons that spin in opposite directions clockwise or
counter- clockwise. +-
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- 3s3s 3p3p 3d3d 2s2s 2p2p A Cross Section of an Atom 1s1s
n0p+n0p+ The first energy level has only one sublevel (1s). The
second energy level has two sublevels (2s and 2p). The third energy
level has three sublevels (3s, 3p, & 3d). *Although the diagram
suggests that electrons travel in circular orbits, this is a
simplification and is not actually the case. Rings of Saturn # of
sublevels per energy level = n
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- Quantum Numbers n = 3 n = 2n = 1 Principallevel Sublevel
Orbital ssp sp d pxpx pypy pzpz d xy d xz d yz dz2dz2 d x 2 - y 2
pxpx pypy pzpz
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- Maximum Electron Capacities of Subshells and Principal Shells n
1 2 3 4...n l 0 0 1 0 1 2 0 1 2 3 Sublevel designation designation
s s p s p d s p d f Orbitals in sublevel sublevel 1 1 3 1 3 5 1 3 5
7 Sublevel capacity capacity 2 2 6 2 6 10 2 6 10 14 Principal
energy Level capacity Level capacity 2 8 18 32...2n 2
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- Quantum Numbers 1. Principal # n 2. Ang. Mom. # l 3. Magnetic #
m 4. Spin # s energy level sublevel (s,p,d,f) orbitalelectron Each
electron has a unique address: Each electron has a unique address:
2p x +1/2 Energy level sublevel orbital Electron spin
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- ATOMIC STRUCTURE There are 3 ways to represent the electron
arrangement of an atom: 1.Electronic Configuration 3.Lewis Dot
Diagrams 2. Orbital Filling Diagrams
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- 3 ways to represent electron arrangements in atoms: Orbital
Notation (orbital filling diagrams): An orbital is represented by a
line or box. An orbital is represented by a line or box. The lines
are labeled with the principal quantum number and the sublevel
letter. The lines are labeled with the principal quantum number and
the sublevel letter. Arrows represent the electrons. Arrows
represent the electrons. An orbital containing one electron is
written as , an orbital with two electrons is written as . An
orbital containing one electron is written as , an orbital with two
electrons is written as .
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- Electron Configurations 2p42p42p42p4 Energy Level Sublevel
Number of electrons in the sublevel A list of all the electrons in
an atom (or ion) 1s 2 2s 2 2p 4
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- Filling Rules for Electron Orbitals Pauli Exclusion Principle
No two electrons in an atom can have the same 4 quantum numbers. No
two electrons in an atom can have the same 4 quantum numbers. Each
orbital can only hold TWO electrons with opposite spins. Each
orbital can only hold TWO electrons with opposite spins. Wolfgang
Pauli
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- General Rules Aufbau Principle Electrons fill the lowest energy
orbitals first. Electrons fill the lowest energy orbitals first.
Lazy Tenant Rule Lazy Tenant Rule 2s2s 3s3s 4s4s 5s5s 6s6s 7s7s
1s1s 2p2p 3p3p 4p4p 5p5p 6p6p 3d3d 4d4d 5d5d 6d6d 4f4f 5f5f 1s1s
2s2s 2p2p 3s3s 3p3p 4s4s 4p4p 3d3d 4d4d 5s5s 5p5p 6s6s 7s7s 6p6p
6d6d 4f4f 5f5f 5d5d Energy * Aufbau is German for building up
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- Examples: (Remember that you must place one electron into each
orbital of a sublevel before a second electron in placed into an
orbital.) Hydrogen 1s 1sHelium Lithium 1s 2s 1s 2s Carbon 1s 2s 2p
x 2p y 2p z 1s 1 1s 2 1s 2 2s 1 Be 1s 2s 1s 2 2s 2 Boron 1s 2s 2p x
2p y 2p z 1s 2 2s 2 2p 1 1s 2 2s 2 2p 2
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- 1s 2s 2p x 2p y 2p z Carbon 1s 2 2s 2 2p 2 Hunds Rule Within a
sublevel, place one electron per orbital before pairing them. Empty
Bus Seat Rule
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- Hunds Rule Also called The Monopoly Rule Also called The
Monopoly Rule In Monopoly, you have to build houses EVENLY. You can
not put 2 houses on a property until all the properties has at
least 1 house. In Monopoly, you have to build houses EVENLY. You
can not put 2 houses on a property until all the properties has at
least 1 house.
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- O 8e - Orbital Diagram Orbital Diagram Electron Configuration
Electron Configuration 1s 2 2s 2 2p 4 Orbital Notation 1s 2s 2p O
15.9994 8
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- Orbital Filling Element 1s 2s 2p x 2p y 2p z 3s Configuration
Orbital Filling Element 1s 2s 2p x 2p y 2p z 3s Configuration
Electron H He Li C N O F Ne Na 1s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2
2p 6 1s 2 2s 2 2p 5 1s 2 2s 2 2p 4 1s 2 2s 2 2p 3 1s 2 2s 2 2p 2 1s
2 2s 1 1s 2 Orbital Notations Electron H He Li C N O F Ne Na 1s 1
1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 1s 2 2s 2 2p 5 1s 2 2s 2 2p 4 1s
2 2s 2 2p 3 1s 2 2s 2 2p 2 1s 2 2s 1 1s 2
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- 4f4f 4d4d 4p4p 4s4s n = 4 3d3d 3p3p 3s3s n = 3 2p2p 2s2s n = 2
1s1s n = 1 Energy Filling order becomes irregular after Ar Because
of overlapping of electron clouds (orbitals) in larger atoms Write
the electron config. for K Sc ?
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- Slide 134
- Diagonal Rule Must be able to write it for the test! Without
it, you may not get correct answers ! Must be able to write it for
the test! Without it, you may not get correct answers ! The
diagonal rule is a memory device that helps you remember the order
of the filling of the orbitals from lowest energy to highest energy
The diagonal rule is a memory device that helps you remember the
order of the filling of the orbitals from lowest energy to highest
energy Order in which orbitals are filled with electrons 1s 2 2s 2
2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p
6 7s 2
- Slide 135
- 1s22s23s24s25s26s27s21s22s23s24s25s26s27s2 2p63p64p65p66p67p6
2p63p64p65p66p67p6 3d 10 4d 10 5d 10 6d 10 7d 10 4f 14 5f 14 6f 14
7f 14 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2
4f 14 5d 10 6p 6 7s 2
- Slide 136
- Electron Configurations Ws. #3 Symbol# e - Orbital Diagram and
Longhand Electron Configuration Mg 12 1s 2s 2p 3s P 15 1s 2s 2p 3s
3p V 23 1s 2s 2p 3s 3p 3d 4s Ge 32 1s 2s 2p 3s 3p 3d 4s ___ 4p Kr
36 1s 2s 2p 3s 3p 3d 4s 4p O 8 1s 2s 2p
- Slide 137
- Part B Rules of Electron Configurations Which of the following
rules is being violated in each electron configuration below?
Explain your answer for each. Hunds Rule, Pauli Exclusion
Principle, Aufbau Principle __ __ Hund s Rule should be 1 electron
in each of 1s 2s 2p the 1st 2 orbitals (not doubled up) ___ _ _
Aufbau Principle need 1s 2s 2p 3s 3p to fill 3s before 3p _ Pauli
Exclusion Principle 1s 2s 2p 3s 3p 1 electron in 3s needs to point
down (opposite spins) 1s 2s 2p 3s 3p 3d Aufbau must fill 4s before
3d
- Slide 138
- Write out the complete electron configuration for the
following: 1) An atom of nitrogen 2) An atom of silver Fill in the
orbital boxes for an atom of nickel (Ni) 2s2s2p2p 3s3s 3p3p 4s4s
3d3d1s1s 1s 2 2s 2 2p 3 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s
2 4d 9
- Slide 139
- Shorthand Notation A way of abbreviating long electron
configurations A way of abbreviating long electron configurations
Since we are only concerned about the outermost electrons, we can
skip to places we know are completely full (noble gases), and then
finish the configuration Since we are only concerned about the
outermost electrons, we can skip to places we know are completely
full (noble gases), and then finish the configuration
- Slide 140
- Shorthand Notation Step 1: Its the Showcase Showdown! Step 1:
Its the Showcase Showdown! Find the closest noble gas to the atom
(or ion), WITHOUT GOING OVER the number of electrons in the atom
(or ion). Write the noble gas in brackets [ ]. Step 2: Find where
to resume by finding the next energy level. Step 2: Find where to
resume by finding the next energy level. Step 3: Resume the
configuration until its finished. Step 3: Resume the configuration
until its finished.
- Slide 141
- neon's electron configuration (1s 2 2s 2 2p 6 ) Shorthand
Configuration for Na [Ne] 3s 1 third energy level one electron in
the s orbital orbital shape Na = [1s 2 2s 2 2p 6 ] 3s 1 electron
configuration A B C D
- Slide 142
- Part B Shorthand Electron Configuration Use the Noble gas that
comes before an element on the periodic table to represent all
inner electrons. Put the symbol of the Noble gas in parentheses Sy
mb ol # e - Shorthand Electron Configuration Ca 20[Ar] 4s 2 Pb
82[Xe] 6s 2 4f 14 5d 10 6p 2 F 9[He] 2s 2 2p 5 U 92[Rn] 7s 2 5f
4
- Slide 143
- Shorthand Notation Chlorine Chlorine Longhand is 1s 2 2s 2 2p 6
3s 2 3p 5 Longhand is 1s 2 2s 2 2p 6 3s 2 3p 5 You can abbreviate
the first 10 electrons with a noble gas, Neon. [Ne] replaces 1s 2
2s 2 2p 6 The next energy level after Neon is 3 is 3 So you start
at level 3 on the diagonal rule (all levels start with s) and
finish the configuration by adding 7 more electrons to bring the
total to 17 [Ne] 3s 2 3p 5
- Slide 144
- Boron is 1s 2 2s 2 p 1 Boron is 1s 2 2s 2 2p 1 The noble gas
preceding Boron is He, so the short way is [He]2s 2 p. The noble
gas preceding Boron is He, so the short way is [He]2s 2 2p 1.
Sulfur is 1s 2 2s 2 Sulfur is 1s 2 2s 2 2p 6 3s 2 3p 4 Short way:
[Ne]3s3p Short way: [Ne]3s 2 3p 4 Example: Titanium [Ar]4s 2 3d 2
[Ar]4s 2 3d 2
- Slide 145
- Practice Shorthand Notation Write the shorthand notation for
each of the following atoms: Write the shorthand notation for each
of the following atoms:KCaIBi
- Slide 146
- Shorthand Configuration [Ar] 4s 2 Electron configurationElement
symbol [Ar] 4s 2 3d 3 [Rn] 7s 2 5f 14 6d 4 [He] 2s 2 2p 5 [Kr] 5s 2
4d 9 [Kr] 5s 2 4d 10 5p 5 [Kr] 5s 2 4d 10 5p 6 Ca V Sg F Ag I Xe Fe
[Ar] 4s 2 3d 6
- Slide 147
- Valence Electrons Electrons are divided between core and
Electrons are divided between core and valence electrons electrons
in the highest energy level of an atom. These are the electrons
that take part in reactions (so most important!) Never more than 8
valence electrons in an atom (2 s & 6 p)
- Slide 148
- Lewis (Electron) Dot Structures Shorter than configs. &
only show valence electrons Element symbol surrounded by # dots =
to number of valence electrons. No more than 2 electrons per side
on symbol. Start at top of symbol, add dots clockwise, one per side
before doubling them up.
- Slide 149
- B 1s 2 2s 2 2p 1, Core = [He], valence = 2s 2 2p 1 Br [Ar] 4s 2
3d 10 4p 5 Core = [Ar] 3d 10 valence = 4s 2 4p 5 Br is in 7A, so 7
valence electrons No. of valence electrons = Group number (for A
groups) B B is in Group 3A, so 3 valence electrons (dots) Br
- Slide 150
- It is very important to define stable here. STABLE means:
STABLE means: (in order of stability) 1. all equal energy orbitals
are FULL 2. 2. all orbitals are half-full 3. 3. all orbitals are
totally empty.
- Slide 151
- 1. The highest energy electron is the LAST one you write in the
electron configuration. 1s 2 2s 2 2p 6 3s 2 3p 5 -- the 3p 5
electron is the last written. *Remember Aufbaus Principle,
electrons fill from the lowest to the highest energy. 2. 2. The
outermost electron is the one with the LARGEST principle quantum
number. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2. The 4 p 2 is the
farthest from the nucleus. OR (2) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d
10. Here, it is the 4s 2 electron, because it has the largest
principle q.n. Some More Stuff!!
- Slide 152
- Irregular Electron configurations sometimes the electron
configuration is NOT what we would predict it to be. Sometimes
electrons are moved because (1) it will result in greater stability
for that atom or (2) for some unknown reason??
- Slide 153
- Examples Predict the electron configuration for Cr #24: [Ar]4s
2 3d 4 However, the real E. C. is [Ar]4s 1 3d 5. The 4s 2 electron
has been moved to 3d to achieve greater stability.
- Slide 154
- 153 Exceptions, Cont. Cr: [Ar] 4s 1 3d 5 NOT Cr: [Ar] 4s 2 3d 4
Because lower energy results from half- filling 6 orbitals with
spins aligned instead of causing repulsion in one of the 3d
orbitals Cu: [Ar] 4s 1 3d 10 NOT [Ar] 4s 2 3d 9 Because easier to
add the electron to a sublevel with four electrons that already
have the same spin than causing repulsion in a different
orbital
- Slide 155
- Electron Orbitals: Electron orbitals Equivalent electron shells
(Bohr) Neon Ne-10: 1s, 2s and 2p
- Slide 156
- Relative Sizes 1s and 2s 1s 2s
- Slide 157
- s, p, and d-orbitals A s orbitals: Hold 2 electrons (outer
orbitals of Groups 1 and 2) B p orbitals: Each of 3 pairs of lobes
holds 2 electrons = 6 electrons (outer orbitals of Groups 13 to 18)
C d orbitals: Each of 5 sets of lobes holds 2 electrons = 10
electrons (found in elements with atomic no. of 21 and higher)
- Slide 158
- Principal Energy Levels 1 and 2
- Slide 159
- Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d
5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration
CLICK ON ELEMENT TO FILL IN CHARTS N HH He Li C N Al Ar F Fe
LaHeLiCNAlArFFeLa
- Slide 160
- Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d
5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration
CLICK ON ELEMENT TO FILL IN CHARTS N H = 1s 1 Hydrogen H He Li C N
Al Ar F Fe LaHeLiCNAlArFFeLa
- Slide 161
- Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d
5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration
CLICK ON ELEMENT TO FILL IN CHARTS N He = 1s 2 Helium HH He Li C N
Al Ar F Fe LaLiCNAlArFFeLa
- Slide 162
- Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d
5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration
CLICK ON ELEMENT TO FILL IN CHARTS N Li = 1s 2 2s 1 Lithium HH He
Li C N Al Ar F Fe LaHeCNAlArFFeLa
- Slide 163
- Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d
5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration
CLICK ON ELEMENT TO FILL IN CHARTS N C = 1s 2 2s 2 2p 2 Carbon HH
He Li C N Al Ar F Fe LaHeLiNAlArFFeLa
- Slide 164
- Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d
5s 5p 4d 6s 6p 5d 4f NUCLEUS Electron Configuration CLICK ON
ELEMENT TO FILL IN CHARTS N N = 1s 2 2s 2 2p 3 Bohr Model Nitrogen
Hunds Rule maximum number of unpaired orbitals. HH He Li C N Al Ar
F Fe LaHeLiCAlArFFeLa
- Slide 165
- Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d
5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration
CLICK ON ELEMENT TO FILL IN CHARTS N F = 1s 2 2s 2 2p 5 Fluorine HH
He Li C N Al Ar F Fe LaHeLiCNAlArFeLa
- Slide 166
- Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d
5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration
CLICK ON ELEMENT TO FILL IN CHARTS N Al = 1s 2 2s 2 2p 6 3s 2 3p 1
Aluminum HH He Li C N Al Ar F Fe LaHeLiCNArFFeLa
- Slide 167
- Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d
5s 5p 4d 6s 6p 5d 4f NUCLEUS Electron Configuration CLICK ON
ELEMENT TO FILL IN CHARTS N Ar = 1s 2 2s 2 2p 6 3s 2 3p 6 Bohr
Model Argon HH He Li C N Al Ar F Fe LaHeLiCNAlFFeLa
- Slide 168
- Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d
5s 5p 4d 6s 6p 5d 4f NUCLEUS CLICK ON ELEMENT TO FILL IN CHARTS Fe
= 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 N HH He Li C N Al Ar F Fe
LaHeLiCNAlArFLa Bohr Model Iron Electron Configuration
- Slide 169
- Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d
5s 5p 4d 6s 6p 5d 4f NUCLEUS CLICK ON ELEMENT TO FILL IN CHARTS La
= 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4s 2 3d 10 4p 6 5s 2 4d 10 5p
6 6s 2 5d 1 N HH He Li C N Al Ar F Fe LaHeLiCNAlArFFe Bohr Model
Lanthanum Electron Configuration
- Slide 170
- Slide 171
- Copyright 2006 Pearson Benjamin Cummings. All rights reserved.
Electron capacities
- Slide 172
- So, where are the electrons of an atom located? Various Models
of the Atom Daltons Model Thompsons Plum Pudding Model Rutherfords
Model Bohrs Planetary Model electrons rotate around the nucleus
Quantum Mechanics Model modern description of the electron in
atoms, derived from a mathematical equation (Schrodingers wave
equation)
- Slide 173
- Development of Atomic Models Rutherford model In the early
twentieth century, Rutherford showed that most of an atom's mass is
concentrated in a small, positively charged region called the
nucleus. Bohr model After Rutherford's discovery, Bohr proposed
that electrons travel in definite orbits around the nucleus.
Thomson model In the nineteenth century, Thomson described the atom
as a ball of positive charge containing a number of electrons.
Quantum mechanical model Modern atomic theory described the
electronic structure of the atom as the probability of finding
electrons within certain regions of space.
- Slide 174
- Slide 175
- Diagonal Rule s s 3p 3d s 2p s 4p 4d 4f s 5p 5d 5f s 6p 6d 6f s
7p 7d 7f 1234567 Steps: 1.Write the energy levels top to bottom.
2.Write the orbitals in s, p, d, f order. Write the same number of
orbitals as the energy level. 3.Draw diagonal lines from the top
right to the bottom left. follow the arrows! 4.To get the correct
order, follow the arrows! By this point, we are past the current
periodic table so we can stop.
- Slide 176
- Atoms like to either empty or fill their outermost level. Since
the outer level contains two s electrons and six p electrons (d
& f are always in lower levels), the optimum number of
electrons is eight. This is called the octet rule.