Transcript CHAPTER 5

Unit 2: Electronic Structures
of Atoms
Part I
Introduction
• Rutherford’s model of the atom was consistent with
the evidence but it had some serious limitations
– How does chemical bonding occur?
– Why do elements form compounds with characteristic
formulae?
• Better understanding gained from theory of
arrangement of electrons
– Based on light given off and absorbed by atoms
3
Electromagnetic Radiation
• Electromagnetic radiation is a form of
energy that consists of electric and
magnetic fields at right angles
4
Electromagnetic Radiation
• All types of Electromagnetic Radiation can be
described in terms of waves
• Any wave is characterized by:
– wavelength: has the symbol  .
– Frequency: has the symbol .
– Amplitude
5
Example of Electromagnetic Radiation
• Visible light is a type of electromagnetic radiation that
can be broken up into a spectrum
• A rainbow is an example of visible light that has been
broken up naturally when rain or mist refracts sunlight
6
Electromagnetic Radiation – Described in the
terminology of waves
(a)
(b)
•
•
the distance b/w any 2 crests (or troughs) is a wavelength λ
The frequency υ, of the wave is the number of crests or troughs that pass
a given point per second.
a) Two waves traveling at the same speed. The upper wave has a longer
wavelength but a shorter frequency. The lower wave has a shorter
wavelength but a higher frequency
7
b) Variation in amplitude
Fig. 5-11, p. 181
Electromagnetic Radiation
• The wavelength , is measured in units of distance such
as m, cm, nm, Å (angstrom).
– 1 Å = 1 x 10-10 m = 1 x 10-8 cm
• The frequency , is measured in units of 1/time e.g. s-1
• Wavelength and frequency are related to each other by:

  = speed of propagation of wave or
 = c
• In a vacuum, the speed of electromagnetic radiation, c, is
the same for all wavelengths, 3.00 x 108 m/s
– Therefore:   = c = 3.00 x 108 m/s
8
Electromagnetic Radiation
• Light from a source of white light is passed
through a slit and then a prism. It spreads into a
continuous spectrum of all wavelengths of visible
light
• Sir Isaac Newton was the first to recorded the
separation of light into tis component colours
9
Fig. 5-13, p. 182
Electromagnetic Radiation
• Visible light is only a small portion of the electromagnetic spectrum
10
Electromagnetic Radiation
• Molecules interact with electromagnetic
radiation.
– Molecules can absorb and emit light.
• Once a molecule has absorbed light
(energy), the molecule can:
1. Rotate
2. Translate
3. Vibrate
4. Electronic transition
11
Electromagnetic Radiation
• Example 1: What is the frequency of green
light of wavelength 5200 Å?
c =    =
c

 1 x 10-10 m 
 = 5.200  10-7 m
(5200 Å) 
 1Å

3.00  108 m/s
=
5.200  10-7 m
 = 5.77  1014 s -1
• Review example 4-4 & Work exercise 54
12
Electromagnetic Radiation
• Under certain conditions it is also possible to describe light
( electromagnetic radiation) as composed of particles or
photons
• According to Max Planck, each photon of light had a
particular amount (a quantum) of energy
– The amt of energy possessed by a photon depends on the
frequency of light
• The energy of a photon is given by Planck’s equation:
E = h  or E =
hc

h = Planck’ s constant = 6.626 x 10
-34
J s
13
Electromagnetic Radiation
• Example 2: What is the energy of a photon
of green light with wavelength 5200 Å?
 What is the energy of 1.00 mol of these
photons?
From Example 1, we know that  = 5.77 x 1014 s -1
E = h
E = (6.626  10-34 J  s)(5.77  1014 s -1 )
E = 3.83  10-19 J per photon
For 1.00 mol of photons :
14
(6.022  10 23 photons)(3 .83  10-19 J per photon) = 231 kJ/mol
Electromagnetic Radiation
Energy of Light
• The frequency of UV light is 2.73 x 1016s-1 and that
of yellow light is 5.26 x 1014s-1. Calculate the
energy, in Joules, of an individual photon of each.
• UV
 E=hv
• Yellow  E =hv
= (6.626 x 10-34 J.s) x (2.73 x 1016s-1 ) = 1.81 x 10-17J
= (6.626 x 10-34 J.s) x (5.26 x 1014s-1) = 3.49 x 10-19 J
As you can see the energy of UV light is higher
and this is one reason why UV light damages
your skin more rapidly than visible light !
Review examples 4-5 and 4-6 & Work exercises 56 and 58
15
The Photoelectric Effect
• The photoelectric effect is one phenomenon that has not
been satisfactorily explained by the wave theory of light
– Apparatus: Light strikes the surface of some metals causing an
electron to be ejected. The electrons travel to the anode and
form a current flowing through the circuit
16
Video: The Photoelectric Effect
Questions:
1. How does the brightness (intensity) of light affect the
number of electrons emitted per second (current)?
2. How does the colour of light affect the number of
electrons emitted per second (current)?
17
The Photoelectric Effect
Observations:
1. Electrons are ejected only if light of sufficient
energy (short wavelength) is used, no matter how
long or how brightly the light shines
•
This wavelength limit is different for different metals
18
The Photoelectric Effect
2.
The # of electrons emitted per second (current) increases
as the brightness (intensity) of the light increases, once the
photon energy is high enough to start the photoelectric
effect
• The amount of current does not depend on the wavelength
(colour) of light used, after the minimum photon energy needed to
start the effect is reached.
• Problem: According to classical theory, even “low” energy light
should cause current to flow if the metal is irradiated long enough.
BUT this was not the case!
The intensity of light is the brightness of the light.
• In wave terms, it is related to the amplitude
• In photon terms, it is the number of photons hitting the target
19
The Photoelectric Effect
• Albert Einstein explained the photoelectric effect
– Explanation involved light having particle-like
behavior.
• Light behaved as though it were composed of photons, each
with a particular amount (a quantum) of energy
– According to Einstein, each photon can transfer its
energy to a single electron during a collision.
• Einstein won the 1921 Nobel Prize in Physics for this work.
20
Atomic Spectra and the Bohr Atom
21
Atomic Spectra and the Bohr Atom
• Every element has a unique spectrum.
• Thus we can use spectra to identify
elements.
– This can be done in the lab, stars,
fireworks, etc.
22
Atomic Spectra and the Bohr Atom
• Atomic and molecular spectra are
important indicators of the underlying
structure of the species.
• In the early 20th century several eminent
scientists began to understand this
underlying structure.
– Included in this list are:
– Niels Bohr
– Erwin Schrodinger
– Werner Heisenberg
23
Atomic Spectra and the Bohr Atom
• When an electric current
 1
1
1 
is passed through H2 gas
= R  2  2 
at very low pressures,

 n1 n 2 
several lines in the
R is the Rydberg constant
spectrum of H2 are
produced
R = 1.097  107 m -1
• In the 19th century Balmer
n1  n 2
and Rydberg showed that
the wavelengths of the
n’ s refer to the numbers
lines in H2 spectrum can
of the energy levels in the
be related by a
mathematical equation:- emission spectrum of hydrogen
24
Atomic Spectra and the Bohr Atom
• In 1913 Neils Bohr provided an explanation for
Rydberg and Balmer’s observations.
• Here are the postulates of Bohr’s theory.
1.
Atom has a number of definite and
discrete energy levels (orbits) in
which an electron may exist
without emitting or absorbing
electromagnetic radiation.
As the orbital radius increases so does
the energy
1<2<3<4<5......
25
Atomic Spectra and the Bohr Atom
2. An electron may move from one discrete
energy level (orbit) to another, but, in so
doing, monochromatic radiation is emitted or
absorbed in accordance with the following
equation.
E 2 - E1 = E = h =
hc

E 2  E1
26
Light is emitted when
an electron FALLS
from a higher energy
orbit to a lower energy
one
Light is absorbed
when an electron
JUMPS to a higher
energy orbit
Arrows showing some possible electronic transitions corresponding to lines in the emission
spectrum for Hydrogen. Transitions in the opposite direction account for the lines in the
absorption spectrum. The biggest energy change occurs when an electron jumps b/w n=1
and n=2; a considerably smaller energy change occurs when an electron jumps b/w n=3 and
n=4
Fig. 5-17, p. 187
Atomic Spectra and the Bohr Atom
3. An electron moves in a circular orbit about the
nucleus and it motion is governed by the
ordinary laws of mechanics and electrostatics,
with the restriction that the angular momentum of
the electron is quantized (can only have certain
discrete values).
angular momentum = mvr = nh/2
h = Planck’s constant
n = 1,2,3,4,...(energy levels)
v = velocity of electron
m = mass of electron
r = radius of orbit
28
Atomic Spectra and the Bohr Atom
Application Rydberg Equation
• Example: What is the wavelength of light
emitted when the hydrogen atom’s energy
changes from n = 4 to n = 2?
n 2 = 4 and n1 = 2
 1
1 

= R  2  2 

 n1 n 2 
1
1
7
-1  1
= 1.097  10 m  2  2 

2 4 
1
1
7
-1  1
= 1.097  10 m   

 4 16 
1
29
Atomic Spectra and the Bohr Atom
1

= 1.097  107 m -1 0.250  0.0625
1

= 1.097  107 m -1 0.1875
1

= 2.057  106 m -1
 = 4.862 10-7 m
-7
 = 1 nm x 4.862 10 m
1 x 10-9 m
λ = 4.862 x 102m =486.2 nm
30
Atomic Spectra and the Bohr Atom
Notice that the wavelength calculated from
the Rydberg equation matches the wavelength
of the green colored line in the H spectrum.
31
(a) For each H atom several transitions are possible
and correspond to light of a specific wavelength. E.g.
n=3  n=2; n= 4  n=1;etc. Because higher levels
become closer in energy, the difference in energy b/w
successive transitions become smaller emission
lines get closer  continuous spectrum
(b) Emission spectrum of H. The lines produced by
electrons falling to n=1 level is known as the Lyman
series; it is in the UV region.
Transitions to n=2 level, known as the Balmer series is in
the VIS region.
32
Atomic Spectra and the Bohr Atom
• Each line in the emission spectrum represents the
difference in energies between two allowed energy
levels for electrons
• When an electron goes from level n2 to n1, the
difference in energy is given off as a single
photon.
• The energy of this photon can be calculated from
Bohr’s equation:
E of photon = E2 – E1
Where E2 > E1
33
Various metals emit distinctive colours of visible light when heated high enough
(flame test). This is the basis for all fireworks, which use salts of different metals like
strontium (red), barium (green), and copper (blue) to produce beautiful colours. p. 190
Atomic Spectra and the Bohr Atom
• Bohr’s theory correctly explains the H emission
spectrum (& other species with containing one
electron e.g. Li2+, He+)
• The theory fails for all other elements
– Failed because it modified classical mechanics to
solve a problem that could not be solved using
classical mechanics
• Led to the development of new physicsQuantum mechanics
35
The Wave Nature of the Electron
• Based on Einstein’s idea that light exhibited both
properties of waves & particles suggested to Louis de
Broglie that very small particles e.g. electrons, also
have wave-like properties.
– The electron wavelengths are described by the de Broglie
relationship.
h
=
mv
h = Planck’ s constant
m = mass of particle
v = velocity of particle
36
The Wave Nature of the Electron
• De Broglie’s assertion was verified by
other scientists within two years.
• Consequently, we now know that electrons
(in fact - all particles) have both a particle
and a wave like character.
– This wave-particle duality is a fundamental
property of submicroscopic particles.
37
The Wave Nature of the Electron
• It was shown that subatomic particles behave very
differently from macroscopic objects
• From the work of de Broglie & others we now know
that they do not obey the laws of classical
mechanics (Newton’s Laws) like larger objects do
• A new kind of mechanics – Quantum Mechanics
which is based on the wave properties of matter,
describes the behaviour of very small particles
– Quantization of energy is a consequence of these
properties
38
Video: The Quantum Mechanical Picture of the Atom
Look at the video and answer the following questions.
1. What are the names of three (3) scientists who contributed to
quantum theory?
2. (a) What are the four (4) quantum numbers?
(b) What is n? What does it tell us?
(c) What is l? What does it tell us?
(d) What is ml? What does it tell us?
39
(e) What is ms? What does it tell us?
Video: The Quantum Mechanical Picture of the Atom
Answers
1. What are the names of three (3) scientists who contributed
to quantum theory?
40
Heisenberg, De Broglie & Schrodinger.
Video: The Quantum Mechanical Picture of the Atom
2. (a) What are the four (4) quantum numbers?
n, l, ml, ms
(b) What is n? What does it tell us?
The principal quantum number. Describes the main energy level or shell
that an electron occupies.
(c) What is l? What does it tell us?
The angular momentum quantum number. Designates the sublevel or
specific shape of atomic orbital that an electron may occupy.
(d) What is ml? What does it tell us?
Magnetic quantum number. Designates a specific orbital within a
subshell.
(e) What is ms? What does it tell us?
Spin quantum number. Refers to the spin of an electron and the
orientation of the magnetic field produced by this spin.
41
The Quantum Mechanical Picture of the Atom
Heisenberg Uncertainty Principle
– It is impossible to determine simultaneously both
the position and momentum of an electron (or
any other small particle).
– Consequently, we must speak of the electrons’
position about the atom in terms of probability
functions.
– These probability functions are represented as
orbitals in quantum mechanics.
42
The Quantum Mechanical Picture of the Atom
Basic Postulates of Quantum Theory
1. Atoms and molecules can exist only in
certain energy states. In each energy
state, the atom or molecule has a definite
energy. When an atom or molecule
changes its energy state, it must emit or
absorb just enough energy to bring it to
the new energy state (the quantum
condition).
43
The Quantum Mechanical Picture of the Atom
2. The frequency of the light emitted or
absorbed by atoms or molecules is
related to the energy change by a
simple equation.
E = h =
hc

44
The Quantum Mechanical Picture of the Atom
3. The allowed energy states of atoms and
molecules can be described by sets of
numbers called quantum numbers.
•
Quantum numbers are the solutions of the
Schrodinger, Heisenberg & Dirac equations.
•
Four quantum numbers are necessary to
describe energy states of electrons in atoms.
45
Quantum Numbers
• The principal quantum number, n describes
the main energy level that an electron
occupies.
– It may be any positive integer
– n = 1, 2, 3, 4, ...... “shells”
46
Quantum Numbers
• Within each shell (defined by the value of n),
different sublevels/ subshells are possible
• The angular momentum quantum number with
the symbol , designates each sublevel
 = 0, 1, 2, 3, 4, 5, .......(n-1)
 Thus the maximum value of  is (n-1)
A letter notation is given to each value of 
  = s, p, d, f, g, h, .......(n-1)
•  tells us the shape of the orbitals.
• These orbitals are the volume around the atom that
the electrons occupy 90-95% of the time.
– This is one of the places where Heisenberg’s Uncertainty principle47
comes into play.
Quantum Numbers
• The symbol for the magnetic quantum number
is m.
• m designates a specific orbital in a subshell
• Within each subshell m may take on any integral
values from – through +  ( zero is included):
– m = (–), ….0,…..(+)
EXAMPLE
• If  = 0 (or an s orbital), then m = 0.
– Notice that there is only 1 value of m.
This implies that there is one s orbital per n value. n  1
48
Quantum Numbers
• If  = 1 (or a p orbital), then m = -1,0,+1.
– There are 3 values of m.
Thus there are three p orbitals per n value. n  2
• If  = 2 (or a d orbital), then m = -2,-1,0,+1,+2.
– There are 5 values of m.
Thus there are five d orbitals per n value. n  3
49
Quantum Numbers
• If  = 3 (or an f orbital), then m = -3,-2,1,0,+1,+2, +3.
– There are 7 values of m.
Thus there are seven f orbitals per n value, n  4
• Theoretically, this series continues on to g,h,i,
etc. orbitals.
– Atoms that have been discovered or made up to this
point in time only have electrons in s, p, d, or f orbitals
in their ground state configurations.
50
Quantum Numbers
• The last quantum number is the spin quantum
number, ms.
• The spin quantum number only has two possible
values.
ms = +½ or -½
• This quantum number tells us the spin and
orientation of the magnetic field of the electrons.
• Wolfgang Pauli in 1925 discovered the Exclusion
Principle.
– No two electrons in an atom can have the same set of
4 quantum numbers.
51
Quantum Numbers
NOTE: Each orbital has only hold a maximum of 2 electrons with opposite spins
52
Table 5-4, p. 195
Part II
Quantum Numbers - Review
n – principal quantum number – describes the main energy level or shell
that an electron occupies (n = 1, 2, 3, 4, …)
l - angular momentum quantum number – designates a sublevel or
specific shape of atomic orbital that an electron may occupy
(l = 0, 1, 2, 3, …, (n-1)
s p d f
ml – magnetic quantum number - designates a specific orbital within a
subshell (ml = (-l),…, 0,….(+l)
ms – spin quantum number - refers to the spin of an electron and the
orientation of the magnetic field produced by this spin (ms = +½ or -1/2)
54
Quantum Numbers - Review
NOTE: Each orbital has only hold a maximum of 2 electrons with opposite spins
55
Table 5-4, p. 195
Atomic Orbitals - Review
Recall: Atomic orbitals are regions of space
where the probability of finding an electron about
an atom is highest.
56
S Orbitals
s orbital properties:
- There is one s orbital per n level.
=0
m = 0 → there is only 1 value of m
- s orbitals are spherically symmetric.
57
p Orbitals
• p orbital properties:
– The first p orbitals appear in the n = 2 shell.
• p orbitals are peanut or dumbbell shaped volumes.
– They are directed along the axes of a Cartesian
coordinate system.
• There are 3 p orbitals per n level.
– They have an  = 1.
– m = -1,0,+1  3 values of m
– The three orbitals are named px, py, pz.
58
p Orbitals
• Each electron in a p orbital has
an equal probability of being in
either lobe
• In the 2 lobes, the wave ψ,
represents that the electron has
opposite phases, corresponding
to the crests and troughs of
waves
– These phases correspond to
mathematical wave functions
with +ve and –ve signs.
– BUT these signs DO NOT
represent charges
• The nucleus defines the origin of
the axes
• Nodal plane  least probability of
finding an electron
59
Fig. 5-23, p. 197
d Orbitals
• d orbital properties:
– The first d orbitals appear in the n = 3 shell.
• The five d orbitals have two different shapes:
– 4 are clover leaf shaped.
– 1 is peanut shaped with a doughnut around it.
– The orbitals lie directly on the Cartesian axes or are
rotated 45o from the axes.
• There are 5 d orbitals per n level.
– The five orbitals are named – d xy , d yz , d xz , d x 2 - y2 , d z 2
– They have an  = 2.
– m = -2,-1,0,+1,+2
 5 values of m
60
d Orbital Shapes
• d orbital shapes
61
Fig. 5-25, p. 197
f Orbitals
• f orbital properties:
– The first f orbitals appear in the n = 4 shell.
• The f orbitals have the most complex shapes.
• There are seven f orbitals per n level.
– The f orbitals have complicated names.
– They have an  = 3
– m = -3,-2,-1,0,+1,+2, +3  7 values of m
– The f orbitals have important effects in the
lanthanide and actinide elements.
62
f Orbital Shapes
• f orbital shapes
63
Generalizations - Atomic Orbitals
• Some generalizations about
atomic orbital size:
– Larger values of n
correspond to larger orbital
size (1s versus 3s)
– In any atom, all orbitals with
same principal quantum
number n are similar in size
(compare 2s with 2p)
– Each orbital with a given n
value becomes smaller as
nuclear charge increases.64
How can two electrons occupy the same orbital?
• Electrons are negatively
charged & behave as
though they were spinning
on their axes  so they
act as tiny magnets
• The motions of e-s
produce magnetic fields
which can interact with
one another.
When 2 e-s have opposite spins, the attraction
due to their opposite magnetic fields helps to
overcome the repulsion of their like charges.
This permits 2 e-s to occupy the orbital.
• Two e-s in the same
orbital have opposite ms
values & are said to be
spin-paired or paired.
65
Fig. 5-28, p. 199
Spin Quantum Number Effects
• Spin quantum number effects:
– Every orbital can hold up to two electrons.
• Consequence of the Pauli Exclusion Principle.
– The two electrons are designated as having
one spin up  and one spin down 
• Spin describes the direction of the
electron’s magnetic fields.
66
Paramagnetism and Diamagnetism
• Unpaired electrons have their spins
aligned   or  
– This increases the magnetic field of the
atom.
• Atoms with unpaired electrons are called
paramagnetic .
– Paramagnetic atoms are attracted to a
magnet.
67
Paramagnetism and Diamagnetism
• Paired electrons have their spins
unaligned .
– Paired electrons have no net magnetic field.
• Atoms with paired electrons are called
diamagnetic.
– Diamagnetic atoms are repelled by a
magnet.
68
Calculating Number of Orbitals & Electrons
• Because two electrons in the same orbital must be
paired, it is possible to calculate the number of
orbitals and the number of electrons in each n
shell.
• The number of orbitals per n level is given by n2.
• The maximum number of electrons per n level is
2n2.
– The value is 2n2 because of the two paired
electrons.
69
Calculating Number of Orbitals & Electrons
Energy Level
n
1
2
3
4
# of Orbitals
Max. # of e-
n2
1(1s)
4 (2s, 2px, 2py, 2pz)
You do it!
9 (3s, three 3p’s, five 3d’s)
16 (4s, three 4p’s, five 4d’s, 7 4f’s)
2n2
2
8
18
32
70
Electron Configurations
• This describe how electrons are assigned in orbitals.
• The arrangement which gives the atom its lowest
energy or unexcited state is called the ground state.
• To determine these configurations, we use the
Aufbau Principle as a guide:
Each atom is “built up” by adding the necessary number
of electrons into orbitals in the way that gives the lowest
total energy for the atom.
71
Electron Configurations
Two general rules help us to predict electron
configurations:
1. Electrons are assigned to orbitals in order of increasing
value of (n + l )
e.g. the 2s subshell has (n + l = 2 + 0 = 2) but the 2p subshell has
(n + l = 2 + 1 = 3), therefore the 2s subshell is filled before the 2p
e.g. the 4s subshell has (n + l = 4 + 0 = 4) but the 3d subshell has
(n + l = 3 + 2 = 5), therefore the 4s subshell is filled before the 3d
2. For subshells with the same value of (n + l ), electrons are
assigned first to the subshell with the lower n.
e.g. 2p subshell has ( n + l = 2 + 1 = 3), but 3s has (n + l = 3 + 0
= 3) therefore the 2p subshell is filled before the 3s because it
72
has a lower value of n.
Electron Configurations
• The Aufbau Principle describes the electron
filling order in atoms.
73
Electron Configurations
You can use this mnemonic to remember the correct filling
order for electrons in atoms.
In general , the (n + 1)s
orbital fills before the nd
orbital.
This is sometimes
referred to as the (n + 1)
rule
74
Electron Configurations
A few things to note:
– The electron configuration of the lowest total energy don’t
always match those predicted by the Aufbau principle
• There are some exceptions e.g. Transitions elements
– The electronic structures are governed by the Pauli
Exclusion Principle:
• No 2 electrons in an atom can have the same four
quantum numbers
– Degenerate orbitals are orbitals of the same energy e.g.
px, py and pz
– Electrons occupy all the orbitals of a given subshell singly
before pairing begins. These unpaired electrons have
parallel spins  Hund’s Rule
75
The Periodic Table & Electron Configurations
• Now we will use the Aufbau Principle to
determine the electronic configurations of the
elements on the periodic chart.
• 1st row elements:
1s
1
H

2 He 
Configurat ion
1
1s
1s 2
76
The Periodic Table & Electron Configurations
• 2nd row elements:
77
The Periodic Table & Electron Configurations
• You try it! 3rd row elements:
3s
11 Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
Ne
Ne
Ne
Ne
Ne
Ne
Ne
Ne
3p
Configurat ion












 


  

  
Ne 3s1
Ne 3s2
Ne 3s2 3p1
Ne 3s2 3p2
Ne 3s2 3p3
Ne 3s2 3p4
Ne 3s2 3p5
Ne 3s2 3p6
78
The Periodic Table & Electron Configurations
• 4th row elements:4s orbital is filled before 3d
3d
4s
19 K Ar 

20
Ca Ar 

21
Sc You do it!
4p
Configurat ion
Ar  4s1
Ar  4s2
79
The Periodic Table & Electron Configurations
• 4th row elements:
3d
19
20
21
4s
K Ar 

Ca Ar 

Sc Ar  

4p
Configurat ion
Ar  4s
Ar  4s2
Ar  4s2 3d1
1
Ti You do it!
22
80
The Periodic Table Electron Configurations
3d
4s
19 K Ar 

20
Ca Ar 

Sc Ar  

22
Ti Ar   

23
V Ar    

Cr Ar      

21
24
4p
Configurat ion
Ar  4s1
Ar  4s2
Ar  4s2 3d1
Ar  4s2 3d 2
Ar  4s2 3d 3
Ar  4s1 3d5
There is an extra measure of stability associated
with half - filled and completely filled orbitals.
81
The Periodic Table &Electron Configurations
3d
25 Mn Ar      
26
4s

4p
Configurat ion
Ar  4s2 3d5
Fe You do it!
82
The Periodic Table &Electron Configurations
3d
25 Mn Ar      
26
27
28
29
4s
4p

Fe Ar      

Co Ar      

Ni Ar      

Cu Ar       
Configurat ion
Ar  4s2 3d5
Ar  4s2 3d 6
Ar  4s2 3d 7
Ar  4s2 3d8
Ar  4s1 3d10
Another exception like Cr and
for essentiall y the same reason.
83
The Periodic Table & Electron Configurations
3d
4s
4p
31 Ga Ar        
32
Ge Ar         
33
As Ar          
34
Se You do it!
Configurat ion
Ar  4s2 3d10 4p1
Ar  4s2 3d10 4p2
Ar  4s2 3d10 4p3
Why isn’t the configuration of Ge, [Ar] 4s1 3d10 4 p3 ?
Ans: It does not occur because of the large energy gap
between ns and np orbitals
84
Electron Configurations & Quantum Numbers
• Now we can write a complete set of
quantum numbers for all of the electrons
in any atom e.g. 11Na:
– When completed there must be one set of 4
quantum numbers for each of the 11
electrons in 11Na (remember Ne has 10
electrons)
3s
11 Na Ne 
3p
Configurat ion
Ne 3s1
85
The Periodic Table & Electron Configurations
n
1st e-
n
1


0
- 1
12stndee
1
00
2 nd e - 1
0
3rd e -
2
4 th e -
2
m
m
ms
m
 1/2s
s electrons
1
0

1/2
0
 1/2  1 s electrons
0
0
 1/2 
0
0
0
0
 1/2 
2 s electrons
 1/2 
86
The Periodic Table & Electron Configurations
n

m
1
0
0
2 nd e - 1
0
0
3rd e -
2
0
0
4 th e -
2
0
0
5 th e -
2
1
-1
6 th e -
2
1
0
7 th e -
2
1
1
8 th e -
2
1
1
9 th e -
2
1
0
10 th e -
2
1
1
 1/2 

 1/2 
 1/2 

2 p electrons
 1/2 
 1/2 

 1/2 

11th e -
3
0
0
 1/2 3 s electron
1st e -
ms
 1/2 

1 s electrons
 1/2 

 1/2 

2 s electrons
 1/2 


87
* n is the principal quantum number. The d1s2, d2s2… designations represent known
configurations. They refer to (n-1)d and ns orbitals. Exceptions shown in grey.