Electronic Structure and the Periodic Table

1
Chapter 6

• Electronic Structure and the Periodic Table

2
Electromagnetic Radiation

• Electromagnetic radiation energy that exhibits
  wavelike behavior as it travels through space
• Electromagnetic spectrum all electromagnetic
  radiation arranged in order of increasing
  wavelength

Figure from http//www.geog.ouc.bc.ca/physgeog/con
tents/6f.html
3
Wave Theory

• Wavelength (l) distance between corresponding
  points on adjacent waves
• Frequency (n) number of waves in a given amount
  of time
• Amplitude maximum deviation from equilibrium

Figure from http//marine.rutgers.edu/mrs/class/jo
sh/em_spec.html
4
c ln

• Speed of light in vacuum 3.00 x 108 m/s
• Wavelength (l) is inversely proportional to
  frequency (n).
• Wavelength is usually measured in nanometers
  (nm).
• 1 nm 1 x 10-9 m
• Frequency is usually measured in hertz (Hz).
• 1 Hz 1 /s

5
c ln

• What is the frequency of green light with a
  wavelength of 513 nm?
• c ln
• 3.00 x 108 m/s (5.13 x 10-7 m)n
• n 5.85 x 1014 Hz
• What is the wavelength, in nanometers, of a beam
  of light with a frequency of 3.67 x 1014 Hz?
• 3.00 x 108 m/s l(3.67 x 1014 Hz)
• l 817 nm, infrared

6
Ultraviolet Catastrophe

• Hot objects, which do not burn, glow red then
  white at very high temperatures.
• Wave theory predicted emission in UV region.
• Max Planck
• Proposed that hot objects lose energy only in
  small, specific amounts
• Energy is quantized.

7
E hn

• Quantum finite quantity of energy
• Photon particle of light
• How much energy does a photon with a frequency of
  5.7 x 1014 Hz have?
• E hn (6.626 x 10-34 Js)(5.7 x 1014 Hz)
• E 3.8 x 10-19 J
• How much energy does a photon with a wavelength
  of 400 nm have?
• c ln n 7.5 x 1014 Hz
• E 5.0 x 10-19 J

Frequency (Hz)
Quantum of Energy (Joules, J)
Plancks constant 6.626 x 10-34 Js
8
Energy States

n 4
excited states
n 3
E
?
n 2
hn
current
?
?
ground state
n 1
9
Spectral Lines of Hydrogen
IR
visible
UV
Figure from http//www.physics.nmt.edu/raymond/ph1
3xbook/node204.html
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Continuous vs. Line Spectra

• Continuous spectrum contains all wavelengths,
  produced from the separation of white light

Figure from http//csep10.phys.utk.edu/astr162/lec
t/light/spectrum.html
11
Continuous vs. Line Spectra

• Line spectrum contains only specific
  wavelengths, produced when electrons move from an
  excited state to ground state
• Examples of line spectra

12
Niels Bohr Model of the Atom
electrons
nucleus
orbits
Bohr pictured the atom as a nucleus surrounded by
specific orbits which contained electrons. The
electrons could only exist at certain orbits and
not between them.
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Energy of Electrons in a Hydrogen Atom

• En - RH (1/n2)
• RH Rydberg constant 2.179 x 10-18 J
• n quantum level

14
Wave Theory of Matter

• Louis de Broglie
• Suggested that electrons also have wave-particle
  duality
• Werner Heisenberg
• Heisenberg Uncertainty Principle it is
  impossible to simultaneously know the position
  and momentum of an electron
• Erwin Shroedinger
• Developed an equation to describe wave motions of
  electrons
• Schroedinger equation yield wave functions y
• Electron density probability of finding an
  electron in certain location, proportional to y2

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Quantum Numbers and Orbitals

• Orbital 3-D region around the nucleus where an
  electron is most likely located
• Quantum numbers specify properties of atomic
  orbitals and electrons
• Principal quantum number (n) main energy level
  (n 1, n 2, etc.)
• Orbital quantum number (l) shape of orbital (s,
  p, d, f)
• Magnetic quantum number (ml) orientation of
  orbital
• Spin quantum number (ms) spin of electron
  (1/2, -1/2)
• Pauli exclusion principle no two e-s in the
  same atom can have the same four quantum numbers

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Quantum Number Values
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Orbital Shapes

• s orbitals
• spherical
• only one possible orientation

Figure adapted from http//www.science.nus.edu/we
bchm/1101/1101ch1.htm
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Orbital Shapes

• p orbitals
• Dumbell shaped
• 3 possible orientations

Figure adapted from http//www.chem.ufl.edu/chm20
40/Notes/Chapter_9/quantum.html
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Orbital Shapes

• d orbitals
• 5 possible orientations

Figure adapted from http//www.chem.ufl.edu/chm20
40/Notes/Chapter_9/quantum.html
20
Orbital Shapes

• f orbitals
• 7 possible orientations

Figure adapted from http//www.chem.ufl.edu/chm20
40/Notes/Chapter_9/Coolstuff/forb.html
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Electron Configuration

• Arrangement of electrons in atoms
• Electrons exist in orbitals which exist in
  subshells (sublevels) which exist at energy
  levels
• Aufbau principle electrons occupy the lowest
  energy level available (Electrons are lazy!)

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Electron Configuration

• Principal quantum number indicates the number of
  sublevels at that energy level
• s orbitals have lowest energy, followed by p, d,
  f
• Recall that there is one s orbital, three p, five
  d, seven f
• Maximum of 2 electrons can occupy each orbital,
  must have opposite spins

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Electron Configuration
Orbital notation
H
He
Li
1
s
1
s
1
s
2
s
Electron-configuration notation
e-s
H 1s1 He 1s2
Li 1s22s1
Energy level
subshell
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Electron Configuration
Orbital notation
(write all available orbitals at a sublevel even
if they are not filled)
B
1
s
2
s
2
p
Electron-configuration notation B 1s22s22p1
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Electron Configuration

• Hunds rule if there is more than one orbital
  at a sublevel, put one electron in each (all with
  the same spin) before doubling up

Orbital notation
O
1
s
2
s
2
p
Electron-configuration notation O 1s22s22p4
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Orbital Filling Order
Once you get to 3p, the 4s orbitals have lower
energy than the 3d. Follow this diagram for
filling orbitals
1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d
6f 7s 7p 7d 7f
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Electron Configuration
Orbital notation
Sr
1s
2s
2p
3s
3p
4s
3d
4p
5s
Electron-configuration notation Sr
1s22s22p63s23p64s23d104p65s2
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Exceptions
Expected orbital configuration of chromium
Cr
1s
2s
2p
3s
3p
4s
3d
but
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Exceptions
It is more stable for the chromium to have a
half-filled 3d sublevel, so one of the 4s
electrons is moved to 3d.
Cr
1s
2s
2p
3s
3p
4s
3d
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Exceptions
Expected orbital configuration of copper
Cu
1s
2s
2p
3s
3p
4s
but
3d
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Exceptions
It is more stable for the copper to have a filled
3d sublevel, so one of the 4s electrons is moved
to 3d.
Cu
1s
2s
2p
3s
3p
4s
3d
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Noble Gas Configuration

• Write atomic symbol of the previous noble gas in
  brackets.
• Write the remainder of the electron configuration
  which comes after that of the noble gas.

The electron configuration for He is 1s2, so you
dont have to include this part.
B He2s22p1
The electron configuration for Kr is
1s22s22p63s23p64s23d104p6, so you dont have to
include this part.
Sr Kr5s2
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Blocks of the Periodic Table
s block
p block
d block
f block
34
Valence Configuration

• Valence electrons electrons in the outer shell
  (highest energy level) of an atom, number valence
  e-s group
• Valence configuration include only electrons at
  the highest energy level and those in sublevels
  of a lower energy level that are not full

Sr 1s22s22p63s23p64s23d104p65s2
Highest energy level is 5
Valence config 5s2
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Valence Configuration
Br 1s22s22p63s23p64s23d104p5
Highest energy level is 4
Valence configuration 4s24p5
Ni 1s22s22p63s23p64s23d8
Highest energy level is 4
3d sublevel is not full
Valence configuration 4s23d8
36
Configurations of Ions

• Write out the electron configuration of the
  neutral atom first, then add or remove electrons
  from the highest energy level.
• Al 1s22s22p63s23p1
• Al3 1s22s22p6
• Se 1s22s22p63s23p64s23d104p4
• Se2- 1s22s22p63s23p64s23d104p6
• Fe 1s22s22p63s23p64s23d6
• Fe3 1s22s22p63s23p63d5

Highest energy level is 3. Remove electrons from
here first. Remove p electrons before s
electrons because they are higher in energy.
Highest energy level is 4. Add electrons here.
Highest energy level is 4. Remove electrons from
here first. If more need to be removed, take
them from the highest sublevel of energy level 3
next.
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Why do atoms form ions?

• All elements want 8 valence electrons.
• This will result in a full s and a full p
  sublevel at the highest energy level.
• It will make them stable.
• They will be isoelectronic (same electron
  configuration) with a noble gas.
• Al3 is isoelectronic with Ne.
• Se2- is isoelectronic with Kr.
• (Transition metals are more complicated because
  their d sublevels are partially filled. The
  charges of these are not predictable.)

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Atomic Radii

• Atomic radius ½ the distance between nuclei of
  two identical bonded atoms
• Decreases left to right across period
• Valence electrons at same energy level
• More protons in nucleus pull electrons closer

• Increases going down a group
• Valence electrons at higher energy level
• Valence electrons more shielded from positive
  nucleus

39
Ionic Radii

• Atoms form ions in order to become isoelectronic
  (same e- config) with a noble gas.
• Elements like to have 8 electrons in their
  valence shells.
• Positive ions
• Much smaller than respective atom
• All electrons from valence shell have been
  removed
• Negative ions
• Somewhat larger than respective atom
• Electrons added to valence shell cause more
  repulsion

40
Ionization Energy

• First ionization energy energy required to
  remove highest energy electron from atom
• Second ionization energy energy required to
  remove second highest energy electron from atom
  (after first has been removed)
• IE increases left to right across a period
• Valence electrons at same energy level
• Greater nuclear charge
• IE decreases down group
• Valence electrons at higher energy level
• Valence electrons more shielded from nuclear
  charge

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Electron Affinity

• Electron affinity energy change that occurs
  when an electron is added to a neutral atom
• Increases left to right across a period
• valence electrons gets closer to eight
• Adding one more electron becomes more favorable
• Valence electrons at same energy level
• Greater nuclear charge
• Decreases down a group
• Valence energy lebel farther from nucleus
• Nuclear charge has less effect on added electron

42
Electronegativity

• Electronegativity ability of an atom in a
  molecule to draw bonded electrons toward itself
• Increases left to right across a period
• Valence electrons at same energy level
• Greater nuclear charge
• Decreases down a group
• Valence (bonding) electrons at higher energy
  level
• Bonding electrons more shielded from nuclear
  charge