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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Electronic Structure Model

• Experimental Evidence
• Line Spectra
• Ionization Energies
• Photoelectron Spectrum
• Intensity/detail of Line Spectra

• What it means
• Electrons in quanitized n
• electrons in each n
• electrons in each n and each sublevel
• Indicates n have sublevels associated with them

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Electronic Structure
n of Sublevel e- in n (2n2) Sublevel Names e- in each sublevel
1 1 2 s s-2
2 2 8 s,p s-2, p-6
3 3 18 s,p,d s-2, p-6, d-10
4 4 32 s,p,d,f s-2, p-6, d-10, f-14
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Order of orbitals (filling) in multi-electron atom
1s lt 2s lt 2p lt 3s lt 3p lt 4s lt 3d lt 4p lt 5s lt 4d lt
5p lt 6s
7.7
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Mg 12 electrons
1s lt 2s lt 2p lt 3s lt 3p lt 4s
1s22s22p63s2
2 2 6 2 12 electrons
Abbreviated as Ne3s2
Ne 1s22s22p6
Cl 17 electrons
1s lt 2s lt 2p lt 3s lt 3p lt 4s
1s22s22p63s23p5
2 2 6 2 5 17 electrons
7.7
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Electron Configurations of Cations and Anions
Of Representative Elements
Na Ne3s1
Na Ne
Atoms lose electrons so that cation has a
noble-gas outer electron configuration.
Ca Ar4s2
Ca2 Ar
Al Ne3s23p1
Al3 Ne
H 1s1
H- 1s2 or He
Atoms gain electrons so that anion has a
noble-gas outer electron configuration.
F 1s22s22p5
F- 1s22s22p6 or Ne
O 1s22s22p4
O2- 1s22s22p6 or Ne
N 1s22s22p3
N3- 1s22s22p6 or Ne
8.2
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Na Ne
Al3 Ne
F- 1s22s22p6 or Ne
O2- 1s22s22p6 or Ne
N3- 1s22s22p6 or Ne
Na, Al3, F-, O2-, and N3- are all isoelectronic
with Ne
H- 1s2
same electron configuration as He
8.2
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Electron Configurations of Transition Metals

• Completely filled or half-completely filled
  d-orbitals have a special stability
• Some irregularities are seen in the electron
  configurations of transition and inner-transition
  metals.

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Electron Configurations of Cations of Transition
Metals
When a cation is formed from an atom of a
transition metal, electrons are always removed
first from the ns orbital and then from the (n
1)d orbitals.
Order of filling 3slt3plt4slt3d But when removing
electrons to form ions for transition
metals Order of removing electrons 4slt3dlt3plt3s
8.2
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Electronic Structure

• Good Points
• Electrons in Quantized Energy Levels
• Maximum electrons in each n is 2n2
• Sublevels (s,p,d,f) and electrons they hold

• Bad Points
• Electrons are placed in orbits about nucleus
• Only explains emission spectra of H2
• Does not address all interactions
• Treats electron as particle

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There are less interactions to take into account
in H than other elements

• Interactions
• Attraction between
• nucleus and negative
• electrons

• Interactions
• Attraction between nucleus
• and negative electrons
• Repulsion between electrons
• in same energy level.
• Shielding effect of filled
• principal energy levels.

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Quantum Theory Revised Electronic Structure
Model

• Dual Nature of the Electron
• Heisenberg Uncertainty Principle

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Dual Nature of Electron

• Previous Concept
• A Substance is Either Matter or Energy
• Matter Definite Mass and Position
• Made of Particles
• Energy Massless and Delocalized
• Position not Specificed
• Wave-like

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Dual Nature of Electron

• Electron is both particle-like and wave-like
  at the same time.
• Previous model only considered particle-like
  nature of the electron

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Heisenberg Uncertainty Principle

• Act of measuring the position and energy of
  electron changes the position of electron
• Better one variable is known (energy) the less
  well the other variable is known (position)

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Orbitals Replace Orbits

• Orbits- Both electron position and energy known
  with certainty
• Orbitals Regions of space where an electrons of
  a given energy will most likely be found

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Quantum TheoryOrbitals Replace Orbits
Orbitals
Orbits
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Schrodinger Wave Equation (Y) Describes
size/shape/orientation of orbitals

• Wave Equation is based on
• Dual Nature of Electron (Electron both
• particle and wave-like at the same time.)
• Heisenberg Uncertainty Principle
• (Orbitals describe a region in space
• an electron will most likely be.)

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Wave Equation (Y)

• Wave Equation describe the size, shape, and
  orientation of the orbital the electron (of a
  given energy) is in. There are four variables in
  the function
• -n Energy and size of orbital
• l Shape of orbital
• ml Orientation of orbital
• ms Electron Spin

(n, l, ml, ms)
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1. Each electron has a unique set of 4 Quantum
   Numbers
2. Each orbital described by the Quantum Numbers can
   hold a maximum of 2 electrons.

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Schrodinger Wave Equation 1st Quantum Number
Y fn(n, l, ml, ms)
principal quantum number n
n 1, 2, 3, 4, .
distance of e- from the nucleus
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Schrodinger Wave Equation 2nd Quantum Number
Y fn(n, l, ml, ms)
angular momentum quantum number l
for a given value of n, l 0, 1, 2, 3, n-1
l 0 s orbital l 1 p orbital l 2
d orbital l 3 f orbital
n 1, l 0 n 2, l 0 or 1 n 3, l 0, 1,
or 2
Shape of the volume of space that the e-
occupies
7.6
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Principal Energy Level, n Sublevel, l Quantum Electron Configuration
1 0 (1,0, , ) 1s
2 0 (2,0, , ) 2s
1 (2, 1, , ) 2p
3 0 (3,0, , ) 3s
1 (3, 1, , ) 3p
2 (3,2, , ) 3d
4 0 (4,0, , ) 4s
1 (4, 1, , ) 4p
2 (4, 2, , ) 4d
3 (4, 3, , ) 4f
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f-orbitals
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Orbital Shapes
Orbital Type Shape Name
s Spherical
p Dumbbell
d Complex
f More complex
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Schrodinger Wave Equation 3rd Quantum Number
Y fn(n, l, ml, ms)
magnetic quantum number ml
for a given value of l ml -l, ., 0, . l
if l 1 (p orbital), ml -1, 0, or 1 if l 2
(d orbital), ml -2, -1, 0, 1, or 2
orientation of the orbital in space
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Number of Degenerate Orbitals Needed for Each
Type of Orbital (Sublevel)
Type of Orbital Maximum of electrons in Orbital of Degenerate Orbitals
s 2 1
p 6 3
d 10 5
f 14 7
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ml -1
ml 0
ml 1
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Schrodinger Wave Equation 4th Quantum Number
Y fn(n, l, ml, ms)
spin quantum number ms
ms ½ or -½
ms -½
ms ½
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Valid Possibilities for Quantum Numbers
Chemistry The Science in Context by Thomas R
Gilbert, Rein V Kriss, and Geoffrey Davies,
Norton Publisher, 2004, p125
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If l 1, then ml -1, 0, or 1
2p
3 orbitals
If l 2, then ml -2, -1, 0, 1, or 2
3d
5 orbitals which can hold a total of 10 e-
7.6
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Three Manners to Convey How Electrons are Arranged

• Electron Configuration List Orbitals and Number
  of Electrons in Each
• (1s22s22p63s2)
• Quantum Numbers (2,0,0,1/2)
• Orbital Diagrams List Orbitals and show location
  of electrons and their spin

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Orbital Diagrams
Pauli exclusion principle - no two electrons in
an atom can have the same four quantum numbers.
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Orbital Diagrams
Carbon 6 electrons
Electron Configuration 1s22s22p2
Orbital Diagram
7.7
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Orbital Diagrams
Oxygen 8 electrons
Electron Configuration 1s22s22p4
Orbital Diagram
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Paramagnetic
Diamagnetic
unpaired electrons
all electrons paired
7.8