Transcript Chapter 8-part 1

Chapter 8
Periodic
Properties of
the Elements
Electron Spin
experiments by Stern and Gerlach showed a beam of silver
atoms is split in two by a magnetic field
 the experiment reveals that the electrons spin on their axis
 as they spin, they generate a magnetic field
◦ spinning charged particles generate a magnetic field
 if there is an even number of electrons, about half the atoms
will have a net magnetic field pointing “North” and the other
half will have a net magnetic field pointing “South”

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Electron Spin Experiment
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Spin Quantum Number, ms

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spin quantum number describes how the electron spins on
its axis
◦ clockwise or counterclockwise
◦ spin up or spin down
spins must cancel in an orbital
◦ paired
ms can have values of ±½
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Pauli Exclusion Principle



no two electrons in an atom may have the same set of 4
quantum numbers
therefore no orbital may have more than 2 electrons, and
they must have with opposite spins
knowing the number orbitals in a sublevel allows us to
determine the maximum number of electrons in the sublevel
 s sublevel has 1 orbital, therefore it can hold 2 electrons
 p sublevel has 3 orbitals, therefore it can hold 6
electrons
 d sublevel has 5 orbitals, therefore it can hold 10
electrons
 f sublevel has 7 orbitals, therefore it can hold 14
electrons
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Allowed Quantum Numbers
Quantum
Number
Principal, n
Number of Significance
Values
1, 2, 3, ...
distance from
nucleus
Azimuthal, l 0, 1, 2, ..., n-1
n
shape of
orbital
Magnetic, ml -l,...,0,...+l
2l + 1 orientation of
orbital
Spin, ms
-½, +½
2
direction of
electron spin
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Values
Quantum Numbers of
Helium’s Electrons
helium has two electrons
 both electrons are in the first energy level
 both electrons are in the s orbital of the first energy level
 since they are in the same orbital, they must have opposite
spins

first
electron
second
electron
n
l
ml
ms
1
0
0
+½
1
0
0
-½
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Electron Configurations
the ground state of the electron is the lowest energy
orbital it can occupy
 the distribution of electrons into the various orbitals in
an atom in its ground state is called its electron
configuration
 the number designates the principal energy level
 the letter designates the sublevel and type of orbital
 the superscript designates the number of electrons in that
sublevel
 He = 1s2

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Orbital Diagrams

we often represent an orbital as a square and the
electrons in that orbital as arrows
◦ the direction of the arrow represents the spin of
the electron
unoccupied
orbital
orbital with
1 electron
orbital with
2 electrons
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Sublevel Splitting in
Multielectron Atoms
the sublevels in each principal energy level of Hydrogen all
have the same energy – we call orbitals with the same energy
degenerate
◦ or other single electron systems
 for multielectron atoms, the energies of the sublevels are split
◦ caused by electron-electron repulsion
 the lower the value of the l quantum number, the less energy
the sublevel has
◦ s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3)
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Penetrating and Shielding

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the radial distribution function shows that
the 2s orbital penetrates more deeply into
the 1s orbital than does the 2p
the weaker penetration of the 2p sublevel
means that electrons in the 2p sublevel
experience more repulsive force, they are
more shielded from the attractive force of
the nucleus
the deeper penetration of the 2s electrons
means electrons in the 2s sublevel
experience a greater attractive force to the
nucleus and are not shielded as effectively
the result is that the electrons in the 2s
sublevel are lower in energy than the
electrons in the 2p
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Penetration & Shielding
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7s
6s
Energy
5s
4s
6d
6p
5p
4d
3d
3p
2p
1s
4f
4p
3s
2s
5d
5f
Notice the following:
1. because of penetration, sublevels within an energy
level are not degenerate
2. penetration of the 4th and higher energy levels is so
strong that their s sublevel is lower in energy than
the d sublevel of the previous energy level
3. the energy difference between levels becomes
smaller for higher energy levels
Filling the Orbitals with
Electrons
energy shells fill from lowest energy to high
 subshells fill from lowest energy to high
◦ s→p→d→f
◦ Aufbau Principle
 orbitals that are in the same subshell have the same energy
 no more than 2 electrons per orbital
◦ Pauli Exclusion Principle
 when filling orbitals that have the same energy, place one
electron in each before completing pairs
◦ Hund’s Rule
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Electron Configurations of
Multielectron Atoms
H:
He:
1s1
1s2
1 electron
s orbital (l = 0)
n=1
2 electrons
s orbital (l = 0)
n=1
Lowest energy to highest energy
Li:
1s2 2s1
1 electrons
s orbital (l = 0)
n=2
Valence Electrons
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the electrons in all the
subshells with the highest
principal energy shell are
called the valence
electrons
electrons in lower energy
shells are called core
electrons
chemists have observed
that one of the most
important factors in the
way an atom behaves,
both chemically and
physically, is the number
of valence electrons
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Examples

For the following atom, write:
◦ the Ground State Electron Configuration
◦ Use short hand notation to write orbital Diagram
◦ Determine the core electrons and valence electrons
 Carbon
 Magnesium
 Sulfur
 Potassium
Electron configuration of transition
metal and atoms in higher energy
state

For the following atom, write:
◦ the Ground State Electron Configuration
◦ Use short hand notation to write orbital Diagram
◦ Determine the core electrons and valence electrons
 Cr
 Br
 Pd
 Bi