chapter 8 periodic properties of the elements. electron spin experiments by stern and gerlach showed...
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Chapter 8Chapter 8Periodic Periodic Properties Properties of the of the ElementsElements
Electron SpinElectron Spinexperiments 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 ExperimentElectron Spin Experiment
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Spin Quantum Number, Spin Quantum Number, mmssspin 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 PrinciplePauli Exclusion Principleno two electrons in an atom may have the same
set of 4 quantum numberstherefore no orbital may have more than 2
electrons, and they must have with opposite spinsknowing the number orbitals in a sublevel allows
us to determine the maximum number of electrons in the sublevels sublevel has 1 orbital, therefore it can hold
2 electronsp sublevel has 3 orbitals, therefore it can
hold 6 electronsd sublevel has 5 orbitals, therefore it can
hold 10 electronsf sublevel has 7 orbitals, therefore it can hold
14 electrons
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Allowed Quantum Allowed Quantum NumbersNumbers
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Quantum Number
Values Number of Values
Significance
Principal, n 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
Quantum Numbers of Quantum Numbers of Helium’s ElectronsHelium’s Electronshelium has two electrons both electrons are in the first energy levelboth electrons are in the s orbital of the first
energy levelsince they are in the same orbital, they must
have opposite spins
n l ml ms
first
electron1 0 0 +½
second
electron1 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 DiagramsOrbital 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
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orbital with1 electron
unoccupiedorbital
orbital with2 electrons
Sublevel Splitting in Sublevel Splitting in Multielectron AtomsMultielectron 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 ShieldingPenetrating and Shielding 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 Penetration & Shielding
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En
erg
y
1s
7s
2s
2p
3s
3p3d
6s6p
6d
4s
4p4d
4f
5s
5p
5d5f
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 Filling the Orbitals with ElectronsElectronsenergy shells fill from lowest energy to highsubshells 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 Electron Configurations of Multielectron AtomsMultielectron Atoms
n = 1
s orbital (l = 0)
1 electronH: 1s1
1s2
n = 1
s orbital (l = 0)
2 electronsHe:
n = 2
s orbital (l = 0)
1 electrons1s2 2s1Li:
Lowest energy to highest energy
Valence ElectronsValence Electrons 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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ExamplesExamplesFor 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 Electron configuration of transition metal and atoms in transition metal and atoms in higher energy statehigher 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