### Electron Structure

```Electron Structure
By
Most Important Ideas of
Quantum Mechanics

Orbital is not an orbit

Electron NOT moving around the nucleus

DO NOT KNOW how it is moving
What does it all mean?



Wave function has no easily visualized
physical meaning
Square of the function indicates the
probability of finding an electron near a
particular point in space
Probabilty distribution
Relative Orbital Size




Difficult to define precisely.
Orbital is a wave function.
Picture an orbital as a three-dimensional
electron density map.
Hydrogen 1s orbital:

Radius of the sphere that encloses 90% of
the total electron probability.
7.5
Probability
Distribution
for the 1s
Wave
Function
Quantum Numbers




Three “variables” contained within the
wave function – n, l, ml
First, the value of n determines the
possible values of l
The values of l, then determine the
possible values of ml
The magnetic spin, ms, is not determined
by the wave function
Quantum Numbers
(wave function variables)



Principal quantum number (n) – size and
energy of the orbital
Angular momentum quantum number (l)
– shape of atomic orbitals
Magnetic quantum number (ml) –
orientation of the orbital in space relative
to the other orbitals in the atom
7.6
Table 7.1 The Angular Momentum
Quantum Numbers (l) and
Corresponding Letters Used to
Designate Atomic Orbitals
Table 7.2 Quantum Numbers for the First
Four Levels of Orbitals in the Hydrogen Atom
Figure 7.13 Two
Representations of the
Hydrogen 1s, 2s, and
3s Orbitals (a) The
Electron Probability
Distribution (b) The
Surface Contains 90%
of the Total Electron
Probability (the Size of
the Oribital, by
Definition)
Figure 7.18 Orbital Energy
Levels for the Hydrogen Atom
1s Orbital
7.7
2px Orbital
7.7
The Boundary Surface
Representations of All Three 2p
Orbitals
7.7
3d x2  y 2
Orbital
7.7
3dxy Orbital
7.7
3d z2 Orbital
7.7
The Boundary Surfaces of All of
the 3d Orbitals
Representation of the 4f Orbitals in
Terms of Their Boundary Surfaces
7.7
Electron Spin



Electron spin quantum number (ms) – can
be +½ or -½.
Pauli exclusion principle - in a given atom
no two electrons can have the same set
of four quantum numbers.
An orbital can hold only two electrons,
and they must have opposite spins.
7.8
Figure 7.19 A Picture of the
Spinning Electron
Orbital Energies
7.9
A Comparison of the Radial Probability
Distributions of the 2s and 2p Orbitals
7.9
Distribution of the 3s Orbital
7.9
Probability Distributions of the
3s, 3p, and 3d Orbitals
7.9
Polyelectronic Atoms

Electron correlation problem:


Since the electron pathways are unknown,
the electron repulsions cannot be calculated
exactly
When electrons are placed in a particular
quantum level, they “prefer” the orbitals
in the order s, p, d, and then f.
7.9
Figure 7.22 The Orders of the Energies
of the Orbitals in the First Three Levels
of Polyelectronic Atoms
Aufbau Principle

As protons are added one by one to the
nucleus to build up the elements,
electrons are similarly added to hydrogen
like orbitals.
7.11
Hund’s Rule

The lowest energy configuration for an
atom is the one having the maximum
number of unpaired electrons allowed by
the Pauli principle in a particular set of
degenerate orbitals.
7.11
The Orbitals Being Filled for
Elements in Various Parts of the
Periodic Table
7.11
Orbital Diagrams and Electron Configuration Worksheet
Using the following diagram, complete the orbital diagram and electron
configuration of the elements assigned by the instructor.
Atom
Orbital diagram
E configuration
__ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __
1s 2s
2p
3s
3p
4s
3d
4p
__ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __
1s 2s
2p
3s
3p
4s
3d
4p
__ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __
1s 2s
2p
3s
3p
4s
3d
4p
__ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __
1s 2s
2p
3s
3p
4s
3d
4p
__ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __
1s 2s
2p
3s
3p
4s
3d
4p
__ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __
1s 2s
2p
3s
3p
4s
3d
4p
Valence Electrons


The electrons in the outermost principal
quantum level of an atom.
1s22s22p6 (valence electrons = 8)
The elements in the same group on the
periodic table have the same valence
electron configuration.
7.11
Exercise

Determine the expected electron configurations
for each of the following.
a) S
b) Ba
c) Eu
7.11
History of the Periodic Table


Many early scientists organized the
elements in a table.
Mendeleev


organized the elements according to physical
and chemical properties.
Correctly predicted the existence of previously
undiscovered elements and their properties
Figure 7.24 Mendeleev's Early
Periodic Table, Published in 1872
Periodic Trends



Ionization Energy
Electron Affinity
7.12
Ionization Energy



Energy required to remove an electron
from a gaseous atom or ion.
In general, as we go across a period from
left to right, the first ionization energy
increases.
In general, as we go down a group from
top to bottom, the first ionization energy
decreases.
7.12
Ionization Energy

Why does the IE decrease down a group?

Why does the IE increase across a period?


The effective nuclear charge increases across a
period and decreases down a column.
What is effective nuclear charge?
The Values of First Ionization Energy
for the Elements in the First Six Periods
7.12
Concept Check


Explain why the graph of ionization energy
versus atomic number (across a row) is not
linear.
Where are the exceptions?
7.12
Concept Check

Which atom would require more energy to
remove an electron? Why?
Li
Cs
7.12
Concept Check

Which has the larger second ionization
energy? Why?
Lithium or Beryllium
7.12
Successive Ionization Energies (KJ
per Mole) for the Elements in
Period 3
7.12
Electron Affinity



Energy change associated with the
addition of an electron to a gaseous
atom.
In general as we go across a period from
left to right, the electron affinities
become more negative.
In general electron affinity becomes more
positive in going down a group.
7.12
Figure 7.32 The Electron Affinity Values for
Atoms Among the First 20 Elements that
Form Stable, Isolated X- Ions


In general as we go across a period from
left to right, the atomic radius decreases.
In general atomic radius increases in
going down a group.
7.12
Figure 7.33 The Radious of an Atom
(r) is Defined as Half the Distance
Between the Nuclei in a Molecule
Consisting of Identical Atoms
Concept Check

Which should be the larger atom? Why?
Na
Cl
7.12
Concept Check
Which is larger?
 The hydrogen 1s orbital
 The lithium 1s orbital
Which is lower in energy?
 The hydrogen 1s orbital
 The lithium 1s orbital
7.12
7.12
7.12
Atomic
Selected
Atoms
Exercise

Arrange the elements oxygen, fluorine, and
sulfur according to increasing:
 Ionization energy
 Atomic size
7.12
Final Thoughts


It is the number of valence electrons that
chemists use to explain an atom’s
chemistry.
BUT do not forget the involvement of the
nucleus.

Electrostatic interaction of positive and
negative charges is the fundamental force
that explains chemical interactions.

i.e., the nucleus of one atom MUST attract the
electrons of another atom to create a chemical
bond . . .
7.13
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