Atoms and
Electronic Structure
Atomic Structure and
Subatomic Particles
1
• Protons and neutrons in
nucleus
– Very small dense
central core
• Electrons “move about”
is the remaining space
of the atom.
Protons Neutrons Electrons
Relative
Charge
Proton
+1
Neutron
0
Electron
-1
Charge
(C)
+1.60218×10–19
0
–1.60218×10–19
Mass
(amu)
1.00727
Mass
(g)
1.67261×10–24
1.00866
1.67493×10–24
0.00055
0.00091×10–24
• Note:
– Proton and neutron about the same mass
– Electron  2000 times smaller
– Electron and proton, same charge, opposite in sign.
atomic mass unit (amu) – 1/12 the mass of a single carbon
atom containing 6 protons and 6 neutrons.
1
Elements: Defined by their
number of protons
• Atomic number (Z) - The number of protons in
the nucleus
• This number identifies the element. (See the
numbers on the periodic table.)
• Mass number (A) - The total number of neutrons
and protons in the nucleus of the atom.
• So, if given the Z and A, how will you determine
the number of neutrons?
Some questions….
•
•
•
•
What is the atomic number of Chlorine?
17
How many protons does chlorine have?
17
Periodic Table
1
1A
1
H
18
8A
2
He
3
Li
2
2A
4
Be
13
3A
5
B
14
4A
6
C
15
5A
7
N
16
6A
8
O
17
7A
9
F
6.941
9.012
10.81
12.01
14.01
16.00
19.00
11
Na
12
Mg
7
7B
25
Mn
10
12
2B
30
Zn
26.98
28.09
30.97
32.06
35.45
39.95
28
Ni
11
1B
29
Cu
18
Ar
8
26
Fe
9
8B
27
Co
17
Cl
20
Ca
6
6B
24
Cr
16
S
24.31
5
5B
23
V
15
P
19
K
4
4B
22
Ti
14
Si
22.99
3
3B
21
Sc
13
Al
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.10
40.08
44.96
47.87
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
75.59
74.92
78.96
79.90
83.80
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.47
87.62
88.91
91.22
92.91
95.96
(98)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
180.9
183.8
186.2
190.2
192.2
195.1
197.0
207.2
209.0
(209)
(210)
(222)
114
Fl
115
116
Lv
117
118
1.008
200.6
204.4
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
Cn
113
(223)
(226)
(227)
(261)
(262)
(266)
(264)
(269)
(268)
(271)
(272)
(285)
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
Lanthanide series
Actinide series
140.1
(289)
10
Ne
20.18
(292)
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
(237)
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(262)
Chlorine
4.003
Atomic No. 12
2
1.
2.
3.
4.
5.
If the mass number is 37, how
many neutrons does the atom
have?
37
17
20
27
30
• How many electrons does an atom of chlorine
have?
• 17
• What element has an atomic number of 12?
• Magnesium
Periodic Table
Isotopes
• Isotopes are atoms that have the same atomic
number but different mass number.
• Most elements have two or more isotopes.
• Symbols can be used to distinguish the different
isotopes:
3
Isotope symbols
Mass number
Atomic number
11
5
B
1.
2.
3.
4.
5.
A
Z
X
Determine the number of
protons, neutrons and electrons.
p=5, n=6, e=6
p=5, n=6, e=5
p=6, n=5, e=5
p=5, n=11, e=5
p=6, n=5, e=5
Example
11
5
B
Is the “5” necessary ?
4
More about isotopes:
• Hydrogen is the only element in which the
different isotopes has their own names.
• 1H is hydrogen
• 2H is deuterium
• 3H is tritium
Ions: Losing and Gaining
Electrons
• Ions - a charged species formed from a neutral
atom or molecule when electrons are gained or
lost.
• Cation - positive charged ion.
– Formed by electrons being lost.
• Anion - negative charged ion.
– Formed by electrons being gained.
Symbol
electrons
protons
neutrons
24Mg
23Na+
35Cl
35Cl56Fe3+
15N
16O227Al3+
5
Give the number of electrons and
neutrons for 35Cl−
1.
2.
3.
4.
5.
e = 16, n = 20
e = 18, n = 18
e = 17, n = 20
e = 18, n = 20
None of the above
The Periodic Law
and the Periodic Table
Elements with
similar
properties have
a repeating
pattern and are
aligned in
columns
• Groups
• Families
1
1A
1
H
3
Li
2
2A
4
Be
6.941
9.012
1.008
11
Na
12
Mg
22.99
24.31
19
K
10
8
26
Fe
9
8B
27
Co
54.94
55.85
43
Tc
44
Ru
(98)
101.1
20
Ca
3
3B
21
Sc
4
4B
22
Ti
5
5B
23
V
6
6B
24
Cr
7
7B
25
Mn
39.10
40.08
44.96
47.87
50.94
52.00
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
85.47
87.62
88.91
91.22
92.91
95.96
13
3A
5
B
14
4A
6
C
15
5A
7
N
16
6A
8
O
17
7A
9
F
10.81
12.01
14.01
16.00
19.00
18
8A
2
He
4.003
10
Ne
20.18
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
26.98
28.09
30.97
32.06
35.45
39.95
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
65.39
69.72
75.59
74.92
78.96
79.90
83.80
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
112.4
114.8
118.7
121.8
127.6
126.9
131.3
28
Ni
11
1B
29
Cu
12
2B
30
Zn
58.93
58.69
63.55
45
Rh
46
Pd
47
Ag
102.9
106.4
107.9
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
132.9
137.3
138.9
178.5
180.9
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
Cn
113
114
Fl
115
116
Lv
117
118
(223)
(226)
(227)
(266)
(264)
Lanthanide series
Actinide series
(261)
(262)
(269)
(268)
(271)
(272)
(285)
(289)
(292)
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
70
Yb
71
Lu
140.1
140.9
144.2
(145)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.0
231.0
238.0
(237)
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(262)
6
2
Understanding Light
• Classical Physics viewed energy as
continuous….ie. Any amount of energy could be
released.
• This was found to be false by Max Planck when
concerning the radiation emitted by a heated solid.
• Planck discovered that atoms and molecules emit
energy only in discrete quantities or quanta. - thus
started quantum theory
Properties of Waves
• Waves - a vibrating disturbance by which energy
is transmitted.
• Waves are characterized by…
Wavelength () is the distance between
identical points on successive waves.
Amplitude is the vertical distance from the
midline of a wave to the peak or trough.
7
Properties of Waves
Put picture of ocean surf here
Frequency () is the number of waves that pass through a
particular point in 1 second (Hz = 1 cycle/s).
The speed (u) of the wave = × 
Visible light consists of electromagnetic waves.
Electric field component
Magnetic field component
Electromagnetic
radiation is the
emission and
transmission of
energy in the form of
electromagnetic
waves.
Energy – the
capacity to do work.
Units – Joule
1 J = 1 kg m2/s2
For electromagnetic radiation:
c = 
Speed of light
wavelength
frequency
Speed of light in a vacuum: 3.00  10 8 m/s
8
A photon has a frequency of 3.5 × 105 Hz.
Convert this frequency into wavelength (nm).
Does this frequency fall in the visible region?
What is the frequency (in Hz) of
light with a wavelength of 490 nm?
1.
2.
3.
4.
5.
6.12  1023
6.12  105
6.12  1014
1.63  10-15
1.63  10-6
9
Interactions of Waves
• Interference – the way waves interact with
each other.
• Constructive Interference – waves align
and increase the amplitude
• Destructive Interference – waves which are
out of phase
Wave versus particle behavior
• Diffraction - the
bending of waves as
they pass through a
slit
• Slit must be a
comparable size to
the wavelength
10
Wave versus particle behavior
An inherent property
of waves.
• Diffraction - the
bending of waves as
they pass through a
slit
• Slit must be a
comparable size to
the wavelength
Interference Pattern
3
Planck’s Quantum Theory
• When solids are heated they emit electromagnetic
radiation.
• It was determined that the amount of radiation
energy emitted was related to its wavelength.
• Classical physics could not account for this fact.
• Planck solved the problem...
11
Planck’s Quantum Theory
• Planck’s assumption: atoms and molecules could
emit (or absorb) energy only in discrete quantities.
• These “bundles” of energy were called quantum the smallest quantity of energy that can be emitted.
E = h
Frequency of light emitted
Planck’s Constant = 6.63 × 10-34 J.s
Energy of a single quanta of energy
The Particle Nature of Light
• Planck did not know the “why” of his discovery.
• Einstein used Planck’s Quantum Theory to help
explain something called the photoelectric effect
and then explained the “why” of Planck’s theory.
12
The Photoelectic Effect
• Light strikes the metal and ejects electrons.
• Expected…any frequency of light would eject the
electron if the light was intense enough.
What “They” Found…
• There was a certain
frequency where
below this frequency
no electrons were
ejected, no matter
how intense the light
was.
What “They” Found…
• Increasing the intensity
increased the number
of electrons that came
off.
13
What “They” Found…
• Increasing the
frequency cause the
electrons which were
ejected to have more
and more Kinetic
energy (meaning they
were moving at greater
speeds)
Einstein’s Explanation of the
Photoelectric Effect
• Light is made of a stream of particles (called
photons).
• Each photon has energy-- E = h
• Each photon, if it has enough energy, can knock
off one electron. (It must overcome the binding
energy ( BE ) of the electron.)
• The more intense the light, the more photons
strike the plate.
h = KE + BE
Binding energy of the electron
Kinetic Energy of the electron
Energy of the photon
14
Dual Nature of Light
1. Waves
2. Particles
• Depending on the experiment, light behaves one
way or the other.
• We will see later that matter has this nature also.
Calculations
• So now you have these two equations:
c=
E=h
• With these two equations if you know one of the
following, you can calculate the other two:
– Energy of photon,
– wavelength of light
– frequency of light
When copper is bombarded with high-energy electrons, X rays
are emitted. Calculate the frequency and energy (in joules)
associated with the photons if the wavelength of the X rays is
0.154 nm.
Frequency
1.
2.
3.
4.
1.96  10181/s
1.96  109 1/s
4.62  107 1/s
4.62  10-2 1/s
Energy
1.29  10-15 J
1.29  10-24 J
3.03  10-26 J
3.06  19-34 J
Forgot to convert from nm to m
15
4
Bohr’s Model of the
Hydrogen atom
Emission Spectra
• The continuous or line spectra of radiation emitted
by substances.
• Obtained by…
– energizing a sample until it produces light
– the light is passed through a prism
– the “rainbow” produced is the spectrum
• The spectrum is not necessarily in the visible region
of electromagnetic radiation.
16
Bohr’s Model of the Atom (1913)
1. e- can only have specific (quantized) energy values
2. light is emitted as e- moves from one energy level to
another
n (principal
quantum
number) =
1,2,3,…
17
5
The Dual Nature of Electrons
• Electron only occupies certain fixed
distances….Why?
• Louis de Broglie provided a solution.
• Electrons are not only particles but are waves
(Dual Nature)
The circumference of the orbit
is equal to an integral number
of wavelengths.
Here the wave does not close on
itself evenly and is a non-allowed
orbit
Expected behavior of particles
Actual electron behavior
18

wavelength
h
mv
mass
velocity
This equation is typically used to calculate the
wavelength
of a particle when the mass and velocity are
known.
Watch your units!
What is the de Broglie wavelength (in nm)
associated with a 2.5 g Ping-Pong ball traveling at
15.6 m/s?

h
mv
19
6
The Uncertainty Principle
• We know electrons
have a wave nature.
• We know electrons
have a particle
nature.
• If we try to observe
both aspects
simultaneously, we
ALWAYS fail.
Heisenberg’s Uncertainty
Principle
x  mv 
h
4
• Uncertainty in position = x
• Uncertainty in velocity = v
• The more you know about position, the less
you know about velocity.
Quantum Mechanics and the
Atom
• Electrons do not move as orbits about the nucleus.
• Due to Heisenberg’s Uncertainty Principle we can
only define regions in space where we have a high
probability of finding an electron.
20
• Schrödinger equations - mathematical equations
used to define the region in space which has a high
probability of finding the electrons. (electron
density)
• These equations take into account the particle and
wave nature of the electron
• These equations launched quantum mechanics.
• These regions in space with high electron density are
called orbitals.
7
Solutions to the Schrödinger
Equation for the Hydrogen Atom
• Complex mathematical functions but they give us
quantum numbers which define the orbitals.
• The four quantum numbers:
– The principal quantum number (n)
– The angular momentum quantum number (l)
– The magnetic quantum number (ml)
– The spin quantum number (ms)
Principal Quantum number, n
• n = 1, 2, 3, 4……
• defines the relative average distance from the
nucleus and the energy of the electron.
• the larger the n value the farther away the electron
is.
• The farther out the electron is, the larger, higher in
energy and more unstable the orbital.
• All electrons with the same n value are in the same
(principal) shell.
21
Energy of an electron in hydrogen:
 1 
En   RH  2 
n 
• RH = 2.1810–18 J
• Rydberg constant
for Hydrogen
The Angular Momentum Q.N., (l)
• l = 0, 1, ….(n-1)
• Tells the shape of the orbital. (we shall see the
shape in a minute.)
• All electrons with the same value of n and l are said
to be in the same subshell.
• Usually we call the subshells by the following
“names”.
Orbital or subshell “names”
l
0
1
2
3
4
Name of
Orbital/Subshell
s
p
d
f
g
Note: Each principal quantum number has its own
allowable values of (l) because l goes up to (n-1)
22
In the shell n=4, what are the
names of the subshells it has?
1.
2.
3.
4.
s only
s and p
s, p and d
s, p, d and f
Magnetic Quantum Number, ml
• ml = -l, ….0…..+l
• Gives the orientation in space of the orbital.
• And, gives the number of orientations (I’ll show
you how here in a minute)
• All electrons with the same n, l, ml are said to be
in the same orbital.
• Let’s stop and derive a table of quantum numbers
[n, l, ml].
Connections between Q.N.’s
n
l
ml
23
Which of the following is not an
allowable set of quantum numbers
[n, l, ml]
1.
2.
3.
4.
5.
6.
[1,0,0]
[2,2,-2]
[3,2,0]
[4,1,-2]
Both 1 and 2
Both 2 and 4
Which set of quantum numbers
will identify an electron in a 4p
subshell?
1.
2.
3.
4.
5.
[4, 3, 2]
[4, 1, 0]
[4, 1, -1]
[4, 2, 0]
Both 2 and 3
8
Atomic Spectroscopy Explained
• Atom absorbs energy, electron promoted to
higher energy level. (Excited state.)
• Electron emits photon of light. (Returns to
the ground state.
24
For
Hydrogen
• The energy of the
transition (E) must
equal the energy of
the photon emitted
(h).
• Notice how the levels
get closer together as
they go farther away
from the nucleus.
This gives the change in energy of the electron.
E = –RH
( n1 – n1 )
2
f
Rydberg Constant =
2.18  10-18J
Final energy
level
2
i
Initial energy
level
This equation can only be used for
the Hydrogen atom
25
Connection between energy of the
electron and energy of the photon.
E photon  Eelectron
E photon  h
c  
Calculate the E of the electron of a hydrogen
atom as the electron drops from the n = 5 state to
the n = 3 state.
1.
2.
3.
4.
+2.91 x 10-20
-2.91 x 10-20
+1.55 x 10-19
-1.55 x 10-19
You forgot to square the n’s
You put n’s in the wrong order
Calculate the wavelength (in nm) of a photon
emitted by a hydrogen atom when its electron
drops from the n = 5 state to the n = 3 state.
1.
2.
3.
4.
5.
323
456
646
811
1280
26
9
Atomic Orbitals
• Orbitals are defined by
the Schrödinger
equations.
• Regions in space where
there is a high
probability of finding
an electron.
• s orbital (when l = 0)
is a sphere
p orbital l = 1
• The ml values when l = 1 are: -1, 0, 1
• Three values means three orientations
in space.
d orbitals l = 2
Nodal Plane –
electron
probability
density is 0
• l = 2, ml = -2, -1, 0, 1, 2
• Five numbers means five
orientations
27
f orbitals l = 3
• l = 3, ml = -3, -2, -1, 0, 1, 2, 3
• Seven numbers means seven orientations.
The Phase of Orbitals
• Phase - the sign of the amplitude of a wave
• Two dimensional waves.
• Three dimensional waves
For a many electron atom:
E(s orbital) < E(p orbital) < E(d orbital) < E(f orbital)
28
10
Orbital Diagram
shows what subshells (orbitals) are occupied by
electrons.
• Ground state – lowest energy state of all
electrons.
Electron Spin and the Pauli
Exclusion Principle
• Electrons spin, either one way or the other.
• All electrons have the same amount of spin.
Spin Quantum number (ms)
ms = +½ or –½
• Example of an orbital diagram for hydrogen
• Arrow shows the spin
• Up arrow = +½; down arrow = –½
Name of subshell
(and orbital)
1s
Pauli Exclusion Principle
• No two electrons in an atom can have the same
four quantum numbers.
• Result – no more than 2 electrons can fit into any
orbital – they will spin in opposite directions.
• Helium – has two electrons in the atom
• Electron configuration: 1s2
• Orbital diagram:
1s
29
Quantum Numbers and
Orbital Diagrams
• Each electron has a set of four quantum numbers
associated with it.
• The first three, give the electron’s location
• The forth gives the spin
[1, 0, 0, +½]
[1, 0, 0, –½]
1s
For a many electron atom:
E(s orbital) < E(p orbital) < E(d orbital) < E(f orbital)
Electron Configuration for
Multielectron atoms
11
• We will learn to write the configuration for
“ground state” atoms.
– Electrons are in their lowest energy state
possible.
• Aufbau principle – building up from lowest to
highest energy
30
• It will be necessary for you to know the
order of orbitals from lowest in energy to
highest energy.
• The following is one way to learn the order.
Element Orbital Diagram Electron Config. Q.N.
H
He
Li
Be
B
Q
Hund’s rule next
Hund’s Rule
• The most stable arrangement of electrons in
subshells is the one with the greatest number of
parallel spins.
• Result: Half fill orbitals in subshells prior to filling
31
Element Orbital Diagram Electron Config. Q.N.
C
N
O
Q
F
Ne
Using the periodic table
k2
Which Q.N.’s are different for the
last two electrons placed in
oxygen?
1.
2.
3.
4.
5.
6.
n
l
ml
ms
l and ms
ml and ms
Back to Fluorine
1
1A
1
H
18
8A
2
He
3
Li
2
2A
4
Be
13
3A
5
B
14
4A
6
C
15
5A
7
N
16
6A
8
O
17
7A
9
F
6.94
1
9.01
2
10.8
1
12.0
1
14.0
1
16.0
0
19.0
0
11
Na
12
Mg
7
7B
25
Mn
10
12
2B
30
Zn
26.9
8
28.0
9
30.9
7
32.0
6
35.4
5
39.9
5
28
Ni
11
1B
29
Cu
18
Ar
8
26
Fe
9
8B
27
Co
17
Cl
20
Ca
6
6B
24
Cr
16
S
24.3
1
5
5B
23
V
15
P
19
K
4
4B
22
Ti
14
Si
22.9
9
3
3B
21
Sc
13
Al
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.1
0
40.0
8
44.9
6
47.8
7
50.9
4
52.0
0
54.9
4
55.8
5
58.9
3
58.6
9
63.5
5
65.3
9
69.7
2
75.5
9
74.9
2
78.9
6
79.9
0
83.8
0
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.4
7
87.6
2
88.9
1
91.2
2
92.9
1
95.9
6
(98)
101.
1
102.
9
106.
4
107.
9
112.
4
114.
8
118.
7
121.
8
127.
6
126.
9
131.
3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
207.
2
209.
0
(209)
(210)
(222)
114
Fl
115
116
Lv
117
118
1.00
8
200.
6
204.
4
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
Cn
113
(223)
132.
9
(226)
137.
3
(227)
138.
9
(261)
(262)
(266)
(264)
(269)
(268)
(271)
(272)
(285)
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
140.
9
144.
2
(145)
Lanthanide series
Actinide series
178.
5
140.
1
180.
9
183.
8
186.
2
190.
2
150.
4
192.
2
152.
0
195.
1
157.
3
197.
0
158.
9
162.
5
(289)
67
Ho
164.
9
68
Er
167.
3
4.00
3
10
Ne
20.1
8
(292)
69
Tm
168.
9
70
Yb
173.
0
71
Lu
175.
0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.
0
231.
0
238.
0
(237)
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(262)
Be able to duplicate this breakdown of the Periodic Table and
you can do the configuration of any element.
32
Slide 95
k2
diagram used for the question has points to the the two paramagnetic electrons of oxygen
kwoodru, 10/29/2007
What are the quantum numbers
of the last two electrons of Be
electron configuration?
1.
2.
3.
4.
[2, 0, 0, ½] [2, 0, 1, ½]
[2, 0, 0, ½] [2, 0, 0, ½]
[2, 0, 0, ½] [2, 0, 0, -½]
[2, 0, 0, ½] [3, 0, 0, ½]
Now to Boron
Procedure for writing the
Electron Configuration
• Find the nearest noble gas which comes before the
element.
• Place the noble gas symbol in square brackets.
This is called the noble gas core.
– Example: [He]
• Now use the breakdown of the periodic table that
you learned to add electron in until you have
reached the element of interest.
33
Write the electron configuration of Cl.
1
1A
1
H
3
Li
2
2A
4
Be
6.94
1
9.01
2
1.00
8
11
Na
12
Mg
22.9
9
24.3
1
19
K
20
Ca
3
3B
21
Sc
4
4B
22
Ti
5
5B
23
V
6
6B
24
Cr
7
7B
25
Mn
39.1
0
40.0
8
44.9
6
47.8
7
50.9
4
52.0
0
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
85.4
7
87.6
2
88.9
1
91.2
2
72
Hf
92.9
1
73
Ta
95.9
6
74
W
10
8
26
Fe
9
8B
27
Co
54.9
4
55.8
5
43
Tc
44
Ru
(98)
75
Re
101.
1
76
Os
12.0
1
14.0
1
16.0
0
19.0
0
10
Ne
20.1
8
15
P
16
S
17
Cl
18
Ar
32.0
6
35.4
5
39.9
5
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
65.3
9
69.7
2
75.5
9
74.9
2
78.9
6
79.9
0
83.8
0
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
58.6
9
63.5
5
46
Pd
47
Ag
79
Au
10.8
1
18
8A
2
He
4.00
3
30.9
7
45
Rh
107.
9
17
7A
9
F
14
Si
58.9
3
78
Pt
16
6A
8
O
28.0
9
28
Ni
106.
4
15
5A
7
N
13
Al
12
2B
30
Zn
77
Ir
14
4A
6
C
26.9
8
11
1B
29
Cu
102.
9
13
3A
5
B
112.
4
80
Hg
114.
8
81
Tl
118.
7
82
Pb
121.
8
127.
6
126.
9
131.
3
55
Cs
56
Ba
57
La
83
Bi
84
Po
85
At
86
Rn
132.
9
137.
3
138.
9
178.
5
180.
9
183.
8
186.
2
190.
2
192.
2
195.
1
197.
0
200.
6
204.
4
207.
2
209.
0
(209)
(210)
(222)
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
Cn
113
114
Fl
115
116
Lv
117
118
(223)
(226)
(227)
(261)
(262)
(266)
(268)
(271)
(272)
(285)
(264)
(269)
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
Er
69
Tm
140.
1
140.
9
144.
2
(145)
150.
4
152.
0
157.
3
158.
9
162.
5
164.
9
167.
3
168.
9
173.
0
175.
0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.
0
231.
0
238.
0
(237)
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(262)
3
Li
2
2A
4
Be
13
3A
5
B
14
4A
6
C
15
5A
7
N
16
6A
8
O
17
7A
9
F
6.94
1
9.01
2
10.8
1
12.0
1
14.0
1
16.0
0
19.0
0
11
Na
12
Mg
7
7B
25
Mn
10
12
2B
30
Zn
26.9
8
28.0
9
30.9
7
32.0
6
35.4
5
39.9
5
28
Ni
11
1B
29
Cu
18
Ar
8
26
Fe
9
8B
27
Co
17
Cl
20
Ca
6
6B
24
Cr
16
S
24.3
1
5
5B
23
V
15
P
19
K
4
4B
22
Ti
14
Si
22.9
9
3
3B
21
Sc
13
Al
31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
39.1
0
40.0
8
44.9
6
47.8
7
50.9
4
52.0
0
54.9
4
55.8
5
58.9
3
58.6
9
63.5
5
65.3
9
69.7
2
75.5
9
74.9
2
78.9
6
79.9
0
83.8
0
37
Rb
38
Sr
39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd
49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe
85.4
7
87.6
2
88.9
1
91.2
2
92.9
1
95.9
6
(98)
101.
1
102.
9
106.
4
107.
9
112.
4
114.
8
118.
7
121.
8
127.
6
126.
9
131.
3
55
Cs
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn
207.
2
209.
0
(209)
(210)
(222)
114
Fl
115
116
Lv
117
118
Lanthanide series
Actinide series
(289)
(292)
70
Yb
71
Lu
Sn:
1
1A
1
H
V:
1.00
8
200.
6
204.
4
87
Fr
88
Ra
89
Ac
104
Rf
105
Db
106
Sg
107
Bh
108
Hs
109
Mt
110
Ds
111
Rg
112
Cn
113
(223)
132.
9
(226)
137.
3
(227)
138.
9
(261)
(262)
(266)
(264)
(269)
(268)
(271)
(272)
(285)
Lanthanide series
Actinide series
178.
5
58
Ce
140.
1
180.
9
183.
8
186.
2
59
Pr
60
Nd
61
Pm
140.
9
144.
2
(145)
190.
2
62
Sm
150.
4
192.
2
63
Eu
152.
0
195.
1
64
Gd
157.
3
197.
0
65
Tb
158.
9
66
Dy
162.
5
(289)
67
Ho
164.
9
68
Er
167.
3
18
8A
2
He
4.00
3
10
Ne
20.1
8
(292)
69
Tm
168.
9
70
Yb
173.
0
71
Lu
175.
0
90
Th
91
Pa
92
U
93
Np
94
Pu
95
Am
96
Cm
97
Bk
98
Cf
99
Es
100
Fm
101
Md
102
No
103
Lr
232.
0
231.
0
238.
0
(237)
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(262)
The electron configuration of
Antimony is
1.
2.
3.
4.
[Kr]5s25d105p6
[Kr]5s24d105p3
[Xe]5s24d105p3
[Xe]5s25d105p2
34
A few other points to know.
• Transition metals - have incompletely filled d
subshells or readily give rise to cations that have
incompletely filled d subshells.
• Exception to learn: If one electron away from the
d subshell being half-full or full, the s electron will
be promoted to fill or half fill it.
• This is due to the stability achieved with half filled
or filled subshells.
Examples of the Exceptions
• chromium
– [Ar] 4s23d4
– [Ar] 4s13d5
NOT!
****this is correct
• silver
– [Kr]5s24d9
– [Kr]5s14d10
NOT!
****this is correct
• Lanthanides (rare earths) - incompletely filled
4f subshells or readily give rise to cations that
have incompletely filled 4f subshells.
• Actinide series - most of these not found in nature
but have been synthesized.
35