Chapter 4 Electron Configurations and Quantum Chemistry
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Transcript of Chapter 4 Electron Configurations and Quantum Chemistry
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©Bires, 2002 Slide 1©Bires, 2004
Chapter 4Electron Configurations and
Quantum Chemistry
Electron configurations determine how an atom behaves in bonding
with other atoms!Topics rearranged from your text. You should read pages 90-116. Atomic Emissions/Abortions deleted
Anyone who says that they can contemplate quantum mechanics without becoming dizzy has not understood the concept in the least. -Niels Bohr
electron
neutron
proton
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©Bires, 2002 Slide 2©Bires, 2004
The Bohr Model• Niels Bohr
– rebuilt the model of the atom placing the electrons in energy levels.
• Quantum chemistry– a discipline that states that energy can be given off
in small packets or quanta of specific size.• What would happen to an electron if the right
sized quanta of energy was added to it?
• What would happen when the electron came back down to its ground state?
EXCITED STATE
Ground state
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Atomic Absorption / Emission• Atomic Absorption
– Electron is given a specific amount of energy, it “quantum jumps” to a higher energy level (excited state).
• Atomic Emission– Electron returns to its ground state, it emits
energy equal to the amount of energy required to raise it to the higher level.
– The difference in energy of the levels produces photons of differing energy
– (blue = higher energy, red = lower energy)
Internet Animation of Atomic Absorption / Emission
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Atomic Absorption / Emission• White light is composed of all visible wavelengths
(colors) of light. (Electromagnetic Radiation)
• The continuous spectrum– having all wavelengths / colors of visible light
(white)
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Emissions Spectrum• When electrons in a gas are charged with energy
– those electrons make their quantum jumps…– Then return to their ground states…– resulting (unique) pattern of light given off is called
an emission spectrum.
• Each element has a different emission spectrum, a sort of quantum fingerprint, that we can use to identify elements in unknown samples.
Flame Tests of Alkali Metals.mov
What happens next?
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Some bright-line emission spectra:• Some emission spectra of elements
– from www.webelements.com:• Sodium, 11 electrons in 3 ground-shells:
• Mercury, 80 electrons in 6 ground-shells:
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Absorption Spectrum• Absorption spectrum
– Given off when white light is shined through a gas– electrons in the gas absorb some wavelengths of
the light
• Viewing light from distant stars through the gasses of our planets allows us to know what chemicals make up those planets.
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Bright-Line Spectra• Max Planck, German physicist
– calculated that energy required to make a quantum jump for specific energy levels and colors.
• Planck’s Constant, 6.6262x10-27
– multiplied by the frequency of the desired emission color equals energy required to produce that jump.
hE c
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Electron Configurations - overview• Bohr model
– electrons exist in specific energy levels.• Electron orbitals (shapes)
– Within each energy level, orbits the electrons can occupy.
• Within each orbital– electrons can be set “spin up” or “spin down”
• Electron configuration– The configuration of electrons in their levels,
orbitals, and spins gives us an atom’s.• Modern Quantum Model
– Electron exists in electron configurations
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Energy Levels (n)• The electrons exist in energy
levels or shells.• The first energy shell can
hold only 2 electrons.– Hydrogen and Helium in their
ground state have electrons that occupy this shell.
• The second shell can hold 8 electrons.
• The third can hold 18 electrons.
2 8
32
18
Shells
All shells after three can hold 32 electrons.
Old School: “KLM notation”
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Orbitals (Shapes)• Orbitals
– electrons travel in set paths.– These paths form shapes.
• Each “shape” can hold 2 electrons
• The smallest orbital is the “s” orbital. The “s” orbital:– Has only 1 shape (holds 2 e-)– Is spherical in shape– Is the lowest energy orbital
S-2
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p-Orbitals• The 2nd orbital shape is the “p” orbital shape.• There are 3 “p” shapes, each holding 2
electrons, for a total of 6 electrons in the “p” orbital.
• The “p” orbitals are:– Dumbbell-shaped
– Higher in energy than the “s”P-6
S-2
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d-Orbitals• The 3rd orbital shape is the “d” orbital shape.• There are 5 “d” orbital shapes, for a total of 10
electrons in the “d” orbital.• “d” orbitals are higher in energy than “p”
orbitals.
S-2
D-10P-6
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f-Orbitals• The last orbital
shape is the “ f ” orbital shape.– “ f ” orbitals have
irregular shapes due to quantum tunneling.
– There are 7 “ f ” shapes, for a total of 14 electrons.
Electrons in f orbitals are very high in energy
S-2 F-14D-10P-6
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Electron “Spin”• Electrons can be “spin up” or “spin down.”
– (by convention, an electron that is alone is “spin up”)
• Pauli’s Exclusion Principle– If two electrons share a shape, they must be spin-
paired (one up and one down).• Hund’s Rule
– As electrons fill orbitals, they first fill each shape available with one electron before spin pairing.
• For instance: take a “p” orbital…it has three orbital-shapes that can hold 2 e- each.
• It would fill like this:Electron Configurations.mov
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Writing Electron Configurations• The Aufbau principle
– electron will fill lower energy orbitals first.
• Energy of electrons:– low energy s < p < d < f high energy– low energy nearer < farther high energy– low energy level 1 < level 7 high energy
• Total energy of an electron:– Product of energy of its shell and the
energy of its orbital.– Guess: Which is lower in energy, an
electron found in 3d or one found in 4s?
s low energyd high energy
close low energy
far high energy
Total energy
=
Shell
x
orbital shape
The 4s electrons are lower in energy!
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©Bires, 2002 Slide 17©Bires, 2004
Writing Electron Configurations• Orbital filling diagram
– Shows how electrons fill into levels and orbitals
1s Electron 2s 2p Configurations 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6 6f14 Don’t Copy this
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©Bires, 2002 Slide 18©Bires, 2004
Building the Orbital Filling Diagram• Begin by listing the shells 1, 2,
3, 4, 5, 6, 7 vertically.• These are your “s” orbitals.• Next, add another column of
number, beginning with 2.• These are your “p” orbitals.• Do the same for “d” and “f”
orbitals, beginning with “3” for the “d” orbitals and “4” for the “f” orbitals.
• Next, add your orbital letters.• Finally, draw diagonal lines as
shown.
1
2
3
65
4
7
2
3
65
4
7
3
65
4
7
65
4
7s
s
s
s
s
s
s
p
p
p
p
p
p
d
d
d
d
d
f
f
f
f
s p d f
...7654654543433221 26101426102610262622 spdfspdspdspspss
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©Bires, 2002 Slide 19©Bires, 2004
Electron Configurations of Some Atoms• Consider Fluorine, with 9 electrons
• What about Copper, with 29 electrons?
Notice the position of the last electron…
522 221F pss
9262622 3433221Cu dspspss
2962622 4333221Cu sdpspss
Both used
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“Blocks” of the periodic table…• The periodic table tells us in which orbital the
last electron should be found.– The last electron in an atom is found in the…
s orbitals p orbitals
d orbitals
f orbitals
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©Bires, 2002 Slide 22©Bires, 2004
Noble Gas Shorthand• Notice the configurations of the noble gases:
• We can shorten the electron configuration of larger elements with NGS.
• Consider Mg:• We can substitute Neon’s e- config, and write Mg:• Similarly, Titanium’s (Ti) e- config:
• Can be shortened to:
21He s 622 221Ne pss 62622 33221Ar pspss
2622 3221Mg spss
23]Ne[Mg s2262622 3433221Ti dspspss
2234]Ar[Ti ds
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©Bires, 2002 Slide 23©Bires, 2004
Ion e- configurations• Ions (elements with more/less electrons) also have
electron configurations.• Consider Sulfur (S):• What if sulfur gained an electron?
• Consider Calcium (Ca):
• What if calcium lost two electrons?
4233]Ne[S ps
5233]Ne[-S ps
262622 433221Ca spspss
626222 33221Ca pspss
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©Bires, 2002 Slide 24©Bires, 2004
• Question:– Why do the atomic radii of atoms decrease as
electrons are added to the atom, as you move from left to right across a period?
• electrostatic attraction– attraction between the electrons (-) in the shells and
the protons(+) in the nucleus – pulls the electrons in
This is what we call a periodic trend
End of chapter 4