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