1
Many electron atoms
and the Periodic Table
2
Objectives







Explain the scientific basis for the Periodic Table
Apply the Aufbau principle, the Pauli Exclusion
principle and Hund’s rule to electrons in an atom
Explain the concept of energy levels in an atom and
the order of filling these levels
Write the electronic configuration of the first 20
elements
Draw and explain a block diagram of the Periodic Table
Explain the meaning and position of the transition
elements
Explain the periodic variations of atomic size, ionisation
energy, electron affinity and electronegativity
3
The Periodic Law
The physical and chemical properties of the
elements are a function of the electronic
configuration of the their atoms which vary with
increasing atomic number in a periodic manner
 hence the Periodic Table

 elements
are grouped according to their electronic
structure
 provides information on chemical properties
4
Many Electron Atoms

Describe in terms of hydrogen orbitals
 Same
quantum numbers and shapes
 energies different

For Hydrogen:
 Depend

on n
3s, 3p, 3d all the same energy
For many e- atom:
 Different
subshells at different energies
2s < 2p depend on n and l
5
Orbital Energies
Many e- Atoms
Hydrogen
0
0
n=4
3d
4s
3p
n=3
n=3
3s
3p
3s
3d
n=2
n=2
2s 2p
n=1
1s
Need to consider
2p
2s
n=1
1s
Effect of increased nuclear charge
Repulsions between electrons
6
Effective Nuclear Charge

Net positive charge from nucleus
attracting an electron




electron shielded by inner eeffective nuclear charge
Zeff = Z - S
S = No. of e- between atom and
nucleus
hence outer shell e- experience less +
charge
effect of “screening” depends
on e- distribution

need to consider orbital shape
Assume 5
Z
e- of interest
3 e- found in sphere
Eff. Nuclear charge = 5 - 3
=2
7
Orbital shape and Energy
s-orbital e- can be close to nucleus
 p-orbital further away than an “s” e d-orbital further from nucleus than “p” e therefore,

 “s”
e- has least screening by other e so a larger effective nuclear charge and is more
tightly bound
(lower energy than “p” ie. more stable)
8
Order of Energy Levels for many e- Atom
In general,
1s < 2s < 2p < 3s < 3p…...
9
An easy way to
remember…...
1s
Increasing Energy
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
10
Energy
11
Electron Spin
need to consider another property of electrons
to determine how electrons populate orbitals
 envisage electron as spinning on own axis

 quantized
 only
2 spin states
 distinguished
by the spin -magnetic quantum
number ms
1
+ 2
1
2
REMIND!
Electron Spin

Stern-Gerlach experiment - when a beam of
ground state H atoms (1s) is passed through a
magnetic field, the beam splits into two beams
12
13
Pauli and Hund
Pauli Exclusion Principle
 no two electrons in an atom can have the same
4 quantum numbers
 there are only 2 values of ms
 hence, an orbital can only hold 2 electrons and
they must have opposite spins
Hund’s Rule
 If orbitals have the same energy, add electrons
singly with spins parallel first
14
The Aufbau Principle
Building up

fill available orbitals with available electrons
starting with lowest energy orbitals (most stable)

this gives ground state
Note

don’t forget Pauli and Hund!!
Building up atoms
15
16
Aufbau Principle
4d
Increasing
energy
4p
3d
3p
4s
3s
2p
2s
1s
shell 1
shell 2
shell 3
shell 4

The aufbau
principle shows
how orbitals are
filled: in the
order to the left.
Two extra rules
are needed.
17
Aufbau Principle-2


Hund’s rule states that when filling a set of orbitals at
the same energy (sunshell), one electron is placed in
each orbital before pairing occurs.
Pauli’s principle tells us that when placing a second
electron in an orbital, its spin must be opposite to the
electron already in the orbital. (Spin is usually
represented by an arrow in an orbital box.)
Aufbau Principle-3a

“Box” or “orbital diagrams for electron configurations.
H (1e)
1S
Start at 1s
He (2e)
1S
The 1s orbital is filled.
Second electron is paired. (Pauli)
18
19
Aufbau Principle-3b
Li (3e)
1S 2S
1s2, 2s1
C (6e)
1S 2S
2p 1s2, 2s2, 2p2
Hund’s rule applies to p subshell
Aufbau Principle-3c
O (8e)
1S 2S

2p 1s2, 2s2, 2p4
In 2p subshell, Hund’s rule! Next electron follows Pauli
principle
Ti (22e) 1s2, 2s2, 2p6, 3s2,
6, 4s2, 3d2
3p
(22e)
4S
3d
20
21
Electrons and the Periodic Table

Electron fit logically into the periodic table. The s block
elements (see next slide) start filling at level 1, the p block at
level 2, and the d block at level 3.
22
Excited State Atoms

Occur when energy has been supplied to raise
e- energy
Ne
1s22s22p6
1s22s22p5….5s1
1s22s22p5….3p1
Ground state
High energy excited state
Low energy excited state
23
Transition metals
Consider the elements in the 4th period
(after Ar 1s22s22p63s23p6)
 after 3p natural sequence would be 3d
 but 4s has (slightly) lower energy than3d
 according to Aufbau must fill 4s before 3d
 ground

state for K is [Ar]4s1 and Ca is [Ar]4s2
as charge increases the energy of 3d decreases
 in
Sc 3d < 4s
 Sc+ [Ar]3d14s1 Sc2+[Ar]3d1
24
Transition metals

energy drop in 3d continues through to Zn

consequences
 for
elements Cu  Zn oxidation state 1+ or 2+
 beyond

Zn 3d electrons have no chemical role
elements from Sc to Zn called d-block
elements
 filling
up the d orbital
25
Transition metals

anomalies
[Ar]3d54s1
expect [Ar]3d44s2
 Cu
[Ar]3d104s1
expect [Ar]3d94s2
 similar occur in fifth period
 also Ru
[Kr]4d75s1
expect [Kr]4d65s2
 Cr

sixth period filling is erratic
 energies

of 4f, 5d & 6s comparable
seventh period all are radioactive
26
The Periodic Table
Organisation of the elements
 electronic
configurations related to position of
element
 elements grouped according to type of orbital the
outer shell electrons are in
BLOCK:
GROUP:
Named for last subshell occupied
the columns
elements have same outer orbital e- configuration
 similar chemical properties
 all
PERIODS: rows
 all
elements same shell
27
The Periodic Table

Subshell
orbitals with same energy

Shell
orbitals with similar energy

eg. 2p
eg. 2s, 2p
Valence Electrons
occupy outermost shell

Core Electrons
occupy filled inner shells

Cl 1s22s22p6 3s23p5
Ne core valence
Closed Shell Atoms
full outer shell - very stable - noble gases
28
29
Periodic Properties
Predicted by considering e- configurations

Sizes of atoms and ions

Ionisation energies

Electron affinities

Electronegativities

Polarising powers and polarisabilities
30
Sizes Of Atoms and Ions
Atoms do not have sharply defined boundaries

Hence, need to define atomic size

Atomic size depends on chemical environment
 ie.
Bonding etc
31
32
This shielding means that each valence electron in effect only “feels”
a +1 charge form the nucleus; this occurs for an highly excited
valence electron. Otherwise the shielding makes the “seen” charge is
higher than +1
33
Defining Atomic and Ionic
Size
estimating size
atomic radius = half the
distance between nearest atoms
in element (in condensed phases)
2r
for ions, base ionic radii on interatomic
distance in ionic crystals. (depends on charge...)
Cu
atom
Cu
covalent bonding
Cu+
Cu+2
1.28
1.17
0.96
0.69
Ao
34
Sizes of Atoms and Ions
Decrease
Increase
Why?
Consider:
1. Principle Quantum number
(shell)
2. Effective nuclear charge
35
Across a Period…….
Increase nuclear charge
but no. of core electrons stay the same
so effective nuclear charge increases while
shell remains the same
hence electrons drawn closer to nucleus
hence decrease atomic size
eg.
Na
1s22s22p63s1
1.91Ao
Mg
1s22s22p63s2
1.60 Ao
36
Down A Group….

More distant electron shell occupied while
effective nuclear charge the same

hence atomic size increases
eg.
Li
1s22s1
1.57 A0
Na
1s22s22p63s1
1.91 Ao
Atomic radii
37
38
Radius of Ions
Cation < Atom
eg.
Na+
0.96 Ao
<
Na
1.91 Ao
1s22s22p6
1s22s22p63s1

lost an ecore electrons exposed
more tightly bound
Decreases across a period eg. Na+ > Mg2+
39
Radius of Ions
Atom < Anion
eg.
Cl
<
0.99 Ao
[Ne]3s23p5
Cl
1.81 Ao
[Ne]3s23p6

gained an eelectron cloud greater
decreases nuclear pull by each electron
Decreases across a period e.g. S2- > Cl-
40
Ionisation Energy
The energy required to remove an electron from a
ground state atom
X(g)  X+(g) + eE = IE1
Measure of stability of outer shell electron configuration
 Depends on




size of the atom
effective nuclear charge
screening effect of inner electrons
type of electron
41
Ionisation Energy
Increase
Decrease
Why?
Consider
1. Effective nuclear charge
2. Distance of e- from nucleus
42
Across a Period….
Increase in effective nuclear charge
Decrease in radius
hence increase attraction between e- and nucleus
hence increase IE
Exceptions: “p” less stable than “s” (B < Be)
orbitals “singly occupied” more stable
than “doubly occupied” (O < N)
43
Down a Group…...
Increase radius while
effective nuclear charge the same
hence Decrease attraction between e- and nucleus
hence decrease IE
Ionization energy
44
45
Electron Affinity
The energy released when an e- added to atom to
form anion
eg.
F(g) + e- 
EA = 328 kJ/mol


F-(g)
a small EA means e- must be forced to stick
measure of ability of atom to accept e-
46
Electron Affinities
Increase
Decrease
Same as IE
Why?
Consider
1. Size
2. Effective Nuclear charge
Low for
Noble gases
47
Electron Affinities
1.
F + e- 
F-
EA = 328.0 kJ mol-1
1s2 2s2 2p6
- stable closed shell = Ne
2.
Ne + e-  Ne-
EA = negative
1s2 2s2 2p6 3s1
- new shell, further from nucleus
- almost totally screened from nuclear charge
- so unstable
48
Electron Affinities
1s22s2
Low EA
Filled s subshell
Next e- higher energy level
so need energy to add e-
Be
N
1s22s22p3
Low EA
Half filled “p”
Adding another e- will cause e- repulsion
hence unfavourable
49
Electronegativity
The ability of an atom to draw e- to itself in a
chemical bond
useful for
predicting extent of charge
transfer between atoms
eg.
“Covalent” 
“Ionic”
H—H
C—H
N—H
NaCl
50
Electronegativity
Related to EA and IE
Cs and F
Cs
F
IE1
low
high
EA
small
large
(Cs gives up e- easily, while F accepts e- easily.)
Electron donor
Electropositive
Electron acceptor
Electronegative
51
Electronegativity
Increase ( size, nuclear charge )
Decrease
Size, same effective nuclear charge
52
Electronegativity and Bond
Type
Numerical scale of electronegativities developed
Paulings electronegativity scale ()
caesium =0.79
fluorine =3.98
For two bonded atoms
(), is a measure of the bond polarity
with the more electronegative atom having more
of the electron density
53
Electronegativity and Bond
Type
Examples
Na
(0.93)
Cl
(3.16)
H
(2.20)
Cl
(3.16)
Cl
(3.16)
Cl
(3.16)
() = 2.23, ionic, Na+Cl>2.0
() = 0.96, polar covalent
H+Cl() = 0,
<0.4
covalent, Cl-Cl