Chapter 8
Chapter 8
1.1 Orbital Energies
1.2 Orbital Energies
 First three quantum numbers uniquely describe
orbitals
1s
2s
3s
4s
2p
3p
4p
3d
4d
 What general principle explains orbital energies?
 Which orbital has higher energy, 1s or 2s?
Why?
4f
 level: contains orbitals with the same n
 Which orbital has higher energy, 2s or 2p?
Why?
y
 sublevel: contains orbitals with the same n and l
 What determines the relative energies of these
orbitals? Which are lower in energy? Which are
higher in energy?
Chapter 8
 Which orbital has higher energy, 2px, 2py or 2pz?
Why?
Chapter 8
1.3 Orbital Energies
1.4 Orbital Energies
 Which orbitals have higher energy, 3s, 3p or 3d?
Why? (what do you need to know?)
 Will the five 3d orbitals have the same energy?
Why or why not?
1
Chapter 8
Chapter 8
1.5 Energy of Orbitals: Summary
 For the same type of orbital (same ______), energy
increases as n increases
((1s < 2s < 3s < 4s…))
 For the same n, energy increases s < p < d < f
(3s < 3p < 3d)
 All orbitals of the same sublevel have the same
energy (they are _________________))
(3px = 3py = 3pz)
2.1 Magnetic Properties: Electron
 A physical phenomenon:
spinning,
p
g charged
g p
particles
produce magnetic fields
 Spinning electrons
produce tiny magnetic
fields
 Electrons can spin in one
of two directions
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p
Chapter 8
2.2 Magnetic Properties: Electron
 Diamagnetic
 Paramagnetic
Chapter 8
2.3 The 4th Quantum Number
 Electron spin, ms: ms = ½ or -½
 Pauli exclusion principle:
2
Chapter 8
Chapter 8
3.1 Electron Configuration Rules
 Electrons fill the lowest energy orbital first (Aufbau
principle)
3.2 Electron Configurations
 Two notations for the arrangement of electrons in
atoms
 Two electrons (max) per orbital
spdf notation
 Maximize parallel spins when filling a sublevel
 Fill orbitals of equal energy with one electron each
before pairing up
 Why don’t we double-up first?
Chapter 8
noble gas notation
Chapter 8
3.3 Electron Configurations
 Hydrogen
3.4 More Examples
 Provide the electron configurations (in spdf and
noble gas notation)
phosphorus
 Lithium
vanadium
 Oxygen
iodine
3
Chapter 8
Chapter 8
3.5 More Examples: Ions
4.1 Periodic Table Organization
(a) S2–
So does S2– = Ar?
(b) Br –
(c) Al3+
Chapter 8
Chapter 8
4.2 Periodic Table Organization
s-block
p-block
d block
d-block
atoms where an s sublevel is being filled
atoms where a p sublevel is being filled
atoms where a d sublevel is being filled
4.3 Periodic Properties
 You will need to know the following:
1 Definitions and chemical equations where
1.
appropriate
 Valence electrons
2. Periodic trends moving up and down and left to right
across the periodic table
 Core electrons
3. Explanations of the trends
 Same group = same number and type of valence
electrons
4. How the atomic properties affect chemical properties
4
Chapter 8
Chapter 8
4.4 Effective Nuclear Charge
 Valence electrons don’t “feel” the full charge of the
nucleus
 Valence electrons are shielded from nuclear charge (Z)
 Take the case of Li
1s22s1
4.5 Effective Nuclear Charge
 Zeff: the positive charge “felt” by a valence electron
 An crude approximation: Zeffff = Z – core electrons
 where Z = atomic number
 Zeff increases across the periodic table





Chapter 8
Lithium: Zeff = 3 – 2 = 1
Carbon: Zeff = 6 – 2 = 4
Fluorine: Zeff = 9 – 2 = 7
Sodium: Zeff =
Silicon: Zeff =
Chapter 8
4.6 Atomic Size
4.7 Atomic Size
 The distance from the nucleus to the edge of the
outermost electron
 Periodic trend:
 Explanation:
5
Chapter 8
Chapter 8
5.1 Ionization Energy (IE)
 The energy required to remove an electron from a
gaseous atom
A(g) + energy  A+(g) + e-
5.2 Sign Conventions
 Energy absorbed (in) = a positive value + 165 kJ
 Energy required (input, raw material)
 Energy released (out) = a negative value - 165 kJ
 Energy produced (output, product)
 Energy input required
The sign tells us which way energy is going
The magnitude tells us how much energy is involved
Chapter 8
Chapter 8
5.3 Ionization Energies
5.4 Ionization Energies: Summary
6
Chapter 8
Chapter 8
5.5 Successive Ionizations
IE1
IE2
IE3
Na
495
4560
Mg
735
1445
IE4
IE5
IE6
IE7
5.6 Successive Ionizations
 For Mg, 2nd IE > 1st IE
 For Al, 3rd IE > 2nd IE > 1st IE
 Why?
7730
Al
580
1815
2740
11600
Si
780
1575
3220
4350
16100
P
1060
1890
2905
4950
6270
21200
S
1005
2260
3375
4565
6950
8490
27000
 For Mg, 3rd IE >>> 2nd IE
 For Al, 4th IE >>> 3rd IE
Example:
Na(g) + IE1  Na+(g) + e-
 Why?
Na+(g) + IE2  Na2+(g) + e-
Chapter 8
Chapter 8
5.7 Electron Affinity
 The energy released when an electron is added to a
gaseous atom
A(g) +
e-

A-(g)
+ energy
 A free electron is not a stable beast. It would rather
be associated with an atom.
5.8 Electron Affinity Predictions
A(g) + e-  A-(g) + energy
 Across a period (left to right): Should it get easier or
harder to add an electron?
 Down a group: Should it get easier or harder to add an
electron?
 If it’s “easy” to add an electron, is the EA a large
negative number or a small negative number?
 Deviations from the general trends
7
Chapter 8
Chapter 8
5.9 Electron Affinity Trends
Chapter 8
5.10 Electron Affinity Summary
Chapter 8
6.1 Ionization: Change in Size
 Why does the size decrease?
6.2 Ionization: Change in Size
 Why does the size increase?
8