Chapter 7 Electronic Structure of Atoms

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Chapter 7 Electronic Structure of Atoms

• Electromagnetic Radiation
• Flame Test/ Emission Spectra
• Quantized Energy Levels
• Bohr Model/ Rydberg Equation
• Principal Energy Levels, n
• First Ionization Energy
• 2nd , 3rd, 4th, etc Ionization Energy

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Chapter 7 Electronic Structureof Atoms

• Sublevels (s, p, d, f)
• Photoelectron Spectroscopy
• Electron Configuration
• Valence Electrons/ Core
• Good/ Bad Point of Atom Model
• Quantum Theory
• Dual Nature of the Electron
• Heisenberg Uncertainty Principle

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Chapter 7 ElectronicStructure of Atoms

• Quantum Numbers (n, l, ml, ms)
• Oribtal Diagrams
• Paramagnetism and Diamagnetism

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Wave Properties

• Frequency (n) is the number of waves that pass
  through a particular point in 1 second (Hz 1
  cycle/s)
• All waves travel through space at same rate c
• C 3.00 X108 m/s

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Wave Properties

• Higher frequency (?)
• Lower wavelength
• Higher Energy

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Higher Frequency and Energy
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Longer Wavelength, Lower Energy
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lithium
sodium
potassium
copper
16.11
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Flame Test

• Electron absorbs
• energy from the flame
• goes to a higher energy
• state.

2. Electron goes back down to lower energy state
and releases the energy it absorbed as light.
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Emission of Energy(2 Possibilities)
or
Continuous Energy Loss
Quantized Energy Loss
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Emission of Energy

• Continuous Energy Loss
• Any and all energy values possible on way down
• Implies electron can be anywhere about nucleus of
  atom
• Continuous emission spectra

• Quantized Energy Loss
• Only certain, restricted, quantized energy values
  possible on way down
• Implies an electron is restricted to quantized
  energy levels
• Line spectra

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Emission Spectrum
Continuous Emission Spectrum
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H2 Emission Spectrum
7.3
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Line Spectra vs. Continuous Emission Spectra

• The fact that the emission spectra H2 gas and
  other molecules is a line rather than continuous
  emission spectra tells us that electrons are in
  quantized energy levels rather than anywhere
  about nucleus of atom.

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Bohrs Model of the Atom (1913)

• e- can only have specific (quantized) energy
  values
• light is emitted as e- moves from one energy
  level to a lower energy level

n (principal quantum number) 1,2,3,
RH (Rydberg constant) 2.18 x 10-18J
7.3
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Electronic Structure

• Electrons are in Quantized Energy Levels
• The maximum number of electrons in principal
  energy level, n, is 2n2.
• Lower energy levels are completely filled before
  higher energy levels are filled.

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First Ionization Energy

• Experimental evidence to show the number
• of electrons in a principal energy level,n.

• Energy to remove the 1st (or most loosely bond)
  electron from an atom in the gaseous state.

-higher IE, harder to remove e-
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First Ionization Energy(MJ/mole)
Harder to remove electron
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Factors in First Ionization Energy

• Principal Energy Level, n, the Most Loosely Bond
  Electron is in.
• Higher n- Easier to Remove Lower IE
• (Explains Trend Going Down Group)
• Charge in the Nucleus of the Atom.
• Higher atomic Harder to Remove Higher IE
• (Explains Trend Going Across Period)

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H vs. He
The nuclear charge of Helium is twice that of
Hydrogen thus youd expect the ionization
energy for helium to be twice that of hydrogen
if the second electron is in the same principal
energy level.
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H vs. Li

• Since the ionization energy of lithium is less
  than half that of H even though the nuclear
  charge of Li is 3X more than H we put the third
  electron in a higher energy level.
• This is how we know principal energy level one
  contains a maximum of 2 electrons.

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Li vs. C
Carbon has twice the nuclear charge of lithium
and its ionization energy is about twice that of
lithium. Thus, we know the 4th, 5th, and 6th
electron are still in n2 since the ionization
energy has not decreased while the nuclear
charge increased.
n2
n1
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C IE 1.09
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Ne vs. C
The ionization energy still increases while the
nuclear charge increases between carbon and neon.
This means we are still filling n2.
n2
n2
n1
n1
Trend Across Period
10
6
Ne IE 2.08
C IE 1.09
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Ne vs. Na

• As the 11th electron is added the ionization
  energy decreases even though the nuclear charge
• increases. This means the 11th electron must go
  in a higher energy level.
• This is how we know n2 holds a maximum of 8
  electrons.

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Li vs. Na
Trend Going Down Group
Even though Na has a higher nuclear charge the
ionization energy is lower since the electron is
in a higher n.
Shielding effect Filled principal energy levels
shield full effect of positive charge of nucleus
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Explaining Trends in Ionization Energy

• Determine the number of electrons in each
  principal energy level.
• Look at the principal energy level the most
  loosely bond electron is in.
• Look at the nuclear charge (atomic )

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Experimental Data That Tells How Many Electrons
are in Each Principal Energy Level

• Trends in the Values of First Ionization Energies
  for Different Elements.
• Trends in the 1st, 2nd, 3rd, 4th, etc Ionization
  Energy for the Same Element.
• Photoelectron Spectroscopy

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Ionization Energies (IE)

• 1st IE Energy to remove the 1st (or most
  loosely bond) electron from an atom in the
  gaseous state.
• 2nd IE Energy to remove the 2nd most loosely
  bond electron from an atom in the gaseous state.
• 3rd IE Energy to remove the 3rd most loosely
  bond electron from an atom in the gaseous state.

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Ionization Energies for Selected Elements,
kJ/Mole
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Multiple Ionization Energies for Li
2nd IE is much larger than 1st IE since electron
is removed from lower n
No 4th Ionization Energy no more electrons
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Multiple Ionization Energies for Be
n2
n2
n1
n1
4
4
Be2 3rd IE 14850
Be 1st IE 899
Big jump in IE means removing electron from lower
n
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Multiple Ionization Energies for Be (continued)
n2
n2
n1
n1
4
4
Be2 4th IE 21005
Be2 5th IE NA
No 5th IE since only have 4 electrons in Be
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Multiple Ionization Energies for Mg
Big jump between 2nd and 3rd IE since removing
electron from lower n
n3
n2
n1
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Mg 1st IE 738
2nd IE is about 2X 1st IE since removing electron
from same n
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Multiple Ionization Energies for Mg (Continued)
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There is a big jump between the 2nd and 3rd IE
for both Be and Mg since both are removing an
electron from a lower n
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Mg always has a lower IE than Be since the
electron is always being removed from a higher
n.
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Experimental Evidence That Suggests That
Principal Energy Levels Have Sublevels Associated
With Them

• Emission Spectrum
• Photoelectron Spectrum

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Emission Spectrum for Barium
Emission Spectrum for Neon
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Ionization Energies are Measured by Photoelectron
Spectroscopy
Ionization Energy 143.4 114.8 28.6 MJ/mole
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Photoelectron Spectrum

• Peak Area is Proportional to Number of Electrons
• Breaks in Ionization Energy Scale Represent
  Different Energy Levels
• Number of Peaks Within Breaks Equals Number of
  Sublevels Within Principal Energy Level

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Taken from Chemistry A Guided Inquiry by
Richard S Moog John Farrell John Wiley
Sons, Inc1999
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Taken from Chemistry A Guided Inquiry by
Richard S Moog John Farrell John Wiley
Sons, Inc1999
Old Picture
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Taken from Chemistry A Guided Inquiry by
Richard S Moog John Farrell John Wiley
Sons, Inc1999
Revised Picture
2p
n2
2p
2s
1s
n1
1s
2s
10
Ne IE 2.08
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Taken from Chemistry A Guided Inquiry by
Richard S Moog John Farrell John Wiley
Sons, Inc1999
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3s
2p
n2
2s
1s
n1

Taken from Chemistry A Guided Inquiry by
Richard S Moog John Farrell John Wiley
Sons, Inc1999
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Taken from Chemistry A Guided Inquiry by
Richard S Moog John Farrell John Wiley
Sons, Inc1999
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4f
4d
4p
n4
4s
3d
n3
3p
3s
2p
n2
2s
1s
n1