chapter 13 electrons in atoms walla walla high school mr. carlsen
TRANSCRIPT
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Chapter 13Electrons in Atoms
Walla Walla High School
Mr. Carlsen
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Section 13.1Models of the Atom
OBJECTIVES:Summarize the development of
atomic theory.Explain the significance of quantized
energies of electrons as they relate to the quantum mechanical model of the atom.
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Greek Idea Democritus and
Leucippus Matter is made up
of solid indivisible particles
John Dalton - one type of atom for each element
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J. J. Thomson’s Model
Discovered electrons Atoms were made of
positive stuff Negative electron
floating around “Plum-Pudding”
model
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Ernest Rutherford’s Model Discovered dense
positive piece at the center of the atom- nucleus
Electrons would surround it
Mostly empty space
“Nuclear model”
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Niels Bohr’s Model He had a question: Why don’t the
electrons fall into the nucleus? Move like planets around the sun. In circular orbits at different levels. Amounts of energy separate one
level from another. “Planetary model”
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Bohr’s planetary model Energy level of an electron
analogous to the rungs of a ladder electron cannot exist between energy
levels, just like you can’t stand between rungs on ladder
Quantum of energy required to move to the next highest level
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The Quantum Mechanical Model
Energy is quantized. It comes in chunks. A quanta is the amount of energy needed to
move from one energy level to another. Since the energy of an atom is never “in
between” there must be a quantum leap in energy.
Erwin Schrödinger derived an equation that described the energy and position of the electrons in an atom
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Things that are very small behave differently from things big enough to see.
The quantum mechanical model is a mathematical solution
It is not like anything you can see.
The Quantum Mechanical Model
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Has energy levels for electrons.
Orbits are not circular. It can only tell us the
probability of finding
an electron a certain
distance from the nucleus.
The Quantum Mechanical Model
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The atom is found inside a blurry “electron cloud”
A area where there is a chance of finding an electron.
Draw a line at 90 % Think of fan blades
The Quantum Mechanical Model
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Atomic Orbitals Principal Quantum Number (n) = the
energy level of the electron. Within each energy level, the complex
math of Schrödinger's equation describes several shapes.
These are called atomic orbitals - regions where there is a high probability of finding an electron.
Sublevels - like theater seats arranged in sections
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Summary
s
p
d
f
# of shapes
Max electrons
Starts at energy level
1 2 1
3 6 2
5 10 3
7 14 4
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By Energy Level First Energy Level only s orbital only 2 electrons 1s2
Second Energy Level
s and p orbitals are available
2 in s, 6 in p 2s22p6
8 total electrons
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By Energy Level Third energy level s, p, and d
orbitals 2 in s, 6 in p, and
10 in d 3s23p63d10
18 total electrons
Fourth energy level
s,p,d, and f orbitals
2 in s, 6 in p, 10 in d, and 14 in f
4s24p64d104f14
32 total electrons
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By Energy Level Any more than
the fourth and not all the orbitals will fill up.
You simply run out of electrons
The orbitals do not fill up in a neat order.
The energy levels overlap
Lowest energy fill first.
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Section 13.2Electron Arrangement in Atoms
OBJECTIVES: Apply the aufbau principle, the Pauli
exclusion principle, and Hund’s rule in writing the electron configurations of elements.
Explain why the electron configurations for some elements differ from those assigned using the aufbau principle.
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Aufbau diagram - page 367
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Electron Configurations The way electrons are arranged in
atoms. Aufbau principle- electrons enter the
lowest energy first. This causes difficulties because of the
overlap of orbitals of different energies. Pauli Exclusion Principle- at most 2
electrons per orbital - different spins
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Electron Configuration Hund’s Rule- When electrons
occupy orbitals of equal energy they don’t pair up until they have to.
Let’s determine the electron configuration for Phosphorus
Need to account for 15 electrons
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The first two electrons go into the 1s orbital
Notice the opposite spins
only 13 more to go...
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
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The next electrons go into the 2s orbital
only 11 more...Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
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The next electrons go into the 2p orbital
only 5 more...Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
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• The next electrons go into the 3s orbital
• only 3 more...Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
• The last three electrons go into the 3p orbitals.
• They each go into separate shapes
• 3 unpaired electrons
• = 1s22s22p63s23p3
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The easy way to remember
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2
• 2 electrons
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Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2 2s2
• 4 electrons
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Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2 2s2 2p6 3s2
• 12 electrons
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Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2 2s2 2p6 3s2
3p6 4s2
• 20 electrons
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Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2 2s2 2p6 3s2
3p6 4s2 3d10 4p6
5s2
• 38 electrons
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Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2 2s2 2p6 3s2
3p6 4s2 3d10 4p6
5s2 4d10 5p6 6s2
• 56 electrons
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Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2 2s2 2p6 3s2
3p6 4s2 3d10 4p6
5s2 4d10 5p6 6s2
4f14 5d10 6p6 7s2
• 88 electrons
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Fill from the bottom up following the arrows
1s2s 2p3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
• 1s2 2s2 2p6 3s2
3p6 4s2 3d10 4p6
5s2 4d10 5p6 6s2
4f14 5d10 6p6 7s2
5f14 6d10 7p6 • 108 electrons
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Exceptional Electron Configurations
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Orbitals fill in order Lowest energy to higher energy. Adding electrons can change the
energy of the orbital. Half filled orbitals have a lower
energy. Makes them more stable. Changes the filling order
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Write these electron configurations
Titanium - 22 electrons 1s22s22p63s23p64s23d2
Vanadium - 23 electrons 1s22s22p63s23p64s23d3
Chromium - 24 electrons 1s22s22p63s23p64s23d4 expected But this is wrong!!
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Chromium is actually: 1s22s22p63s23p64s13d5
Why? This gives us two half filled orbitals. Slightly lower in energy. The same principal applies to
copper.
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Copper’s electron configuration
Copper has 29 electrons so we expect: 1s22s22p63s23p64s23d9
But the actual configuration is: 1s22s22p63s23p64s13d10
This gives one filled orbital and one half filled orbital.
Remember these exceptions: d4, d9
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Section 13.3Physics and the Quantum
Mechanical Model OBJECTIVES:
Calculate the wavelength, frequency, or energy of light, given two of these values.
Explain the origin of the atomic emission spectrum of an element.
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Light The study of light led to the
development of the quantum mechanical model.
Light is a kind of electromagnetic radiation.
Electromagnetic radiation includes many kinds of waves
All move at 3.00 x 108 m/s = c
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Parts of a wave
Wavelength
AmplitudeOrigin
Crest
Trough
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Parts of Wave - p.372 Origin - the base line of the energy. Crest - high point on a wave Trough - Low point on a wave Amplitude - distance from origin to crest Wavelength - distance from crest to
crest Wavelength is abbreviated by the Greek
letter lambda =
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Frequency The number of waves that pass a
given point per second. Units: cycles/sec or hertz (hz or sec-1) Abbreviated by Greek letter nu =
c =
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Frequency and wavelength Are inversely related As one goes up the other goes down. Different frequencies of light are
different colors of light. There is a wide variety of frequencies The whole range is called a spectrum,
Fig. 13.10, page 373
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Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
Low energy
High energy
Low Frequency
High Frequency
Long Wavelength
Short WavelengthVisible Light
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Prism White light is
made up of all the colors of the visible spectrum.
Passing it through a prism separates it.
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If the light is not white By heating a gas
with electricity we can get it to give off colors.
Passing this light through a prism does something different.
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Atomic Spectrum Each element
gives off its own characteristic colors.
Can be used to identify the atom.
How we know what stars are made of.
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• These are called discontinuous spectra, or line spectra
• unique to each element.
• These are emission spectra
• The light is emitted given off
• Sample 13-2 p.375
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Light is a Particle Energy is quantized. Light is energy Light must be quantized These smallest pieces of light are
called photons. Photoelectric effect? Energy & frequency: directly related.
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Energy and frequency E = h x E is the energy of the photon is the frequency h is Planck’s constant h = 6.6262 x 10 -34 Joules x sec. joule is the metric unit of Energy
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The Math in Chapter 13
2 equations so far:
c = E = h Know these!
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Examples What is the wavelength of blue light
with a frequency of 8.3 x 1015 hz? What is the frequency of red light
with a wavelength of 4.2 x 10-5 m? What is the energy of a photon of
each of the above?
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Explanation of atomic spectra When we write electron
configurations, we are writing the lowest energy.
The energy level, and where the electron starts from, is called it’s ground state- the lowest energy level.
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Changing the energy Let’s look at a hydrogen atom
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Changing the energy Heat or electricity or light can move the
electron up energy levels (“excited”)
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Changing the energy As the electron falls back to ground
state, it gives the energy back as light
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May fall down in steps Each with a different energy
Changing the energy
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{{{
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Further they fall, more energy, higher frequency.
This is simplified the orbitals also have different energies
inside energy levels All the electrons can move around.
Ultraviolet Visible Infrared
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What is light? Light is a particle - it comes in chunks. Light is a wave- we can measure its
wavelength and it behaves as a wave If we combine E=mc2 , c=, E = 1/2 mv2
and E = h We can get: = h/mv called de Broglie’s equation Calculates the wavelength of a particle.
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Sample problem What is the approximate mass of a
particle having a wavelength of 10-7 meters, and a speed of 1 m/s? Use = h/mv
= 6.6 x 10-27
(Note: 1 J = N x m; 1 N = 1 kg x m/s2
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Matter is a Wave Does not apply to large objects Things bigger than an atom A baseball has a wavelength of
about 10-32 m when moving 30 m/s An electron at the same speed has
a wavelength of 10-3 cm Big enough to measure.
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The physics of the very small Quantum mechanics explains how
the very small behaves. Classic physics is what you get
when you add up the effects of millions of packages.
Quantum mechanics is based on probability
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Heisenberg Uncertainty Principle
-It is impossible to know exactly the location and velocity of a particle.
The better we know one, the less we know the other.
Measuring changes the properties. Instead, analyze interactions with
other particles
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More obvious with the very small
To measure where a electron is, we use light.
But the light moves the electron And hitting the electron changes the
frequency of the light.
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Moving Electron
Photon
Before
ElectronChanges velocity
Photon changes wavelength
After
Fig. 13.19, p. 382