Chemistry Protons, electrons and neutrons

Prévia do material em texto

1
Atoms and 
Electronic Structure
Atomic Structure and 
Subatomic Particles
• Protons and neutrons in 
nucleus
– Very small dense 
central core
• Electrons “move about” 
is the remaining space 
of the atom.
1
Protons Neutrons Electrons
• 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.
Relative
Charge
Charge 
(C)
Mass 
(amu)
Mass
(g)
Proton +1 +1.60218×10–19 1.00727 1.67261×10–24
Neutron 0 0 1.00866 1.67493×10–24
Electron -1 –1.60218×10–19 0.00055 0.00091×10–24
2
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
18
8A
1
H
1.008 2
2A
13
3A
14
4A
15
5A
16
6A
17
7A
2
He
4.003
3
Li
6.941
4
Be
9.012
5
B
10.81
6
C
12.01
7
N
14.01
8
O
16.00
9
F
19.00
10
Ne
20.18
11
Na
22.99
12
Mg
24.31 3
3B
4
4B
5
5B
6
6B
7
7B 8
9
8B
10 11
1B
12
2B
13
Al
26.98
14
Si
28.09
15
P
30.97
16
S
32.06
17
Cl
35.45
18
Ar
39.95
19
K
39.10
20
Ca
40.08
21
Sc
44.96
22
Ti
47.87
23
V
50.94
24
Cr
52.00
25
Mn
54.94
26
Fe
55.85
27
Co
58.93
28
Ni
58.69
29
Cu
63.55
30
Zn
65.39
31
Ga
69.72
32
Ge
75.59
33
As
74.92
34
Se
78.96
35
Br
79.90
36
Kr
83.80
37
Rb
85.47
38
Sr
87.62
39
Y
88.91
40
Zr
91.22
41
Nb
92.91
42
Mo
95.96
43
Tc
(98)
44
Ru
101.1
45
Rh
102.9
46
Pd
106.4
47
Ag
107.9
48
Cd
112.4
49
In
114.8
50
Sn
118.7
51
Sb
121.8
52
Te
127.6
53
I
126.9
54
Xe
131.3
55
Cs
132.9
56
Ba
137.3
57
La
138.9
72
Hf
178.5
73
Ta
180.9
74
W
183.8
75
Re
186.2
76
Os
190.2
77
Ir
192.2
78
Pt
195.1
79
Au
197.0
80
Hg
200.6
81
Tl
204.4
82
Pb
207.2
83
Bi
209.0
84
Po
(209)
85
At
(210)
86
Rn
(222)
87
Fr
(223)
88
Ra
(226)
89
Ac
(227)
104
Rf
(261)
105
Db
(262)
106
Sg
(266)
107
Bh
(264)
108
Hs
(269)
109
Mt
(268)
110
Ds
(271)
111
Rg
(272)
112
Cn
(285)
113 114
Fl
(289)
115 116
Lv
(292)
117 118
Lanthanide series
58
Ce
140.1
59
Pr
140.9
60
Nd
144.2
61
Pm
(145)
62
Sm
150.4
63
Eu
152.0
64
Gd
157.3
65
Tb
158.9
66
Dy
162.5
67
Ho
164.9
68
Er
167.3
69
Tm
168.9
70
Yb
173.0
71
Lu
175.0
Actinide series
90
Th
232.0
91
Pa
231.0
92
U
238.0
93
Np
(237)
94
Pu
(244)
95
Am
(243)
96
Cm
(247)
97
Bk
(247)
98
Cf
(251)
99
Es
(252)
100
Fm
(257)
101
Md
(258)
102
No
(259)
103
Lr
(262)
Chlorine Atomic No. 12
3
If the mass number is 37, how 
many neutrons does the atom 
have?
1. 37
2. 17
3. 20
4. 27
5. 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:
4
Isotope symbols
X
A
Z
Mass number
Atomic number
Determine the number of 
protons, neutrons and electrons. 
1. p=5, n=6, e=6
2. p=5, n=6, e=5
3. p=6, n=5, e=5
4. p=5, n=11, e=5
5. p=6, n=5, e=5
B11
5
Example
B
11
5
Is the “5” necessary ?
5
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
35Cl-
56Fe3+
15N
16O2-
27Al3+
6
Give the number of electrons and 
neutrons for 35Cl−
1. e = 16, n = 20
2. e = 18, n = 18
3. e = 17, n = 20
4. e = 18, n = 20
5. None of the above
The Periodic Law 
and the Periodic Table
1
1A
18
8A
1
H
1.008
2
2A
13
3A
14
4A
15
5A
16
6A
17
7A
2
He
4.003
3
Li
6.941
4
Be
9.012
5
B
10.81
6
C
12.01
7
N
14.01
8
O
16.00
9
F
19.00
10
Ne
20.18
11
Na
22.99
12
Mg
24.31
3
3B
4
4B
5
5B
6
6B
7
7B 8
9
8B
10 11
1B
12
2B
13
Al
26.98
14
Si
28.09
15
P
30.97
16
S
32.06
17
Cl
35.45
18
Ar
39.95
19
K
39.10
20
Ca
40.08
21
Sc
44.96
22
Ti
47.87
23
V
50.94
24
Cr
52.00
25
Mn
54.94
26
Fe
55.85
27
Co
58.93
28
Ni
58.69
29
Cu
63.55
30
Zn
65.39
31
Ga
69.72
32
Ge
75.59
33
As
74.92
34
Se
78.96
35
Br
79.90
36
Kr
83.80
37
Rb
85.47
38
Sr
87.62
39
Y
88.91
40
Zr
91.22
41
Nb
92.91
42
Mo
95.96
43
Tc
(98)
44
Ru
101.1
45
Rh
102.9
46
Pd
106.4
47
Ag
107.9
48
Cd
112.4
49
In
114.8
50
Sn
118.7
51
Sb
121.8
52
Te
127.6
53
I
126.9
54
Xe
131.3
55
Cs
132.9
56
Ba
137.3
57
La
138.9
72
Hf
178.5
73
Ta
180.9
74
W
183.8
75
Re
186.2
76
Os
190.2
77
Ir
192.2
78
Pt
195.1
79
Au
197.0
80
Hg
200.6
81
Tl
204.4
82
Pb
207.2
83
Bi
209.0
84
Po
(209)
85
At
(210)
86
Rn
(222)
87
Fr
(223)
88
Ra
(226)
89
Ac
(227)
104
Rf
(261)
105
Db
(262)
106
Sg
(266)
107
Bh
(264)
108
Hs
(269)
109
Mt
(268)
110
Ds
(271)
111
Rg
(272)
112
Cn
(285)
113 114
Fl
(289)
115 116
Lv
(292)
117 118
Lanthanide series
58
Ce
140.1
59
Pr
140.9
60
Nd
144.2
61
Pm
(145)
62
Sm
150.4
63
Eu
152.0
64
Gd
157.3
65
Tb
158.9
66
Dy
162.5
67
Ho
164.9
68
Er
167.3
69
Tm
168.9
70
Yb
173.0
71
Lu
175.0
Actinide series
90
Th
232.0
91
Pa
231.0
92
U
238.0
93
Np
(237)
94
Pu
(244)
95
Am
(243)
96
Cm
(247)
97
Bk
(247)
98
Cf
(251)
99
Es
(252)
100
Fm
(257)
101
Md
(258)
102
No
(259)
103
Lr
(262)
Elements with 
similar 
properties have 
a repeating 
pattern and are 
aligned in 
columns
• Groups
• Families 
7
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
2
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.
8
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 = × 
Properties of Waves
Put picture of ocean surf here
Visible light consists of electromagnetic waves.
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
Electric field component
Magnetic field component
c = 
For electromagnetic radiation:
Speed of light
wavelength
frequency
Speed of light in a vacuum: 3.00  10 8 m/s
9
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. 6.12  1023
2. 6.12  105
3. 6.12  1014 
4. 1.63  10-15
5. 1.63  10-6
10
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 aslit
• Slit must be a 
comparable size to 
the wavelength
11
Wave versus particle behavior
• Diffraction - the 
bending of waves as 
they pass through a 
slit
• Slit must be a 
comparable size to 
the wavelength
Interference Pattern
A
n 
in
he
re
nt
 p
ro
pe
rt
y 
of
 w
av
es
.
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...
3
12
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.
Planck’s Constant = 6.63 × 10-34 J.s
Frequency of light emitted
Energy of a single quanta of energy
E = h
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.
13
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.
14
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
Energy of the photon
Kinetic Energy of the electron
Binding energy of the electron
15
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.
1. 1.96  10181/s 1.29  10-15 J
2. 1.96  109 1/s 1.29  10-24 J
3. 4.62  107 1/s 3.03  10-26 J
4. 4.62  10-2 1/s 3.06  19-34 J
Frequency Energy
Forgot to convert from nm to m
16
Bohr’s Model of the 
Hydrogen atom
4
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.
17
1. e- can only have specific (quantized) energy values
2. light is emitted as e- moves from one energy level to 
another
Bohr’s Model of the Atom (1913)
n (principal 
quantum 
number) = 
1,2,3,…
18
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)
5
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
19

h 
mv
wavelength 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
20
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. 
6
Heisenberg’s Uncertainty 
Principle
4
h
vmx 
• 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.
21
• 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.
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)
7
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. 
22
Energy of an electron in hydrogen:






2
1
n
RE Hn
• 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 Name of 
Orbital/Subshell 
0 s 
1 p 
2 d 
3 f 
4 g 
 
 
 Note: Each principal quantum number has its own 
allowable values of (l) because l goes up to (n-1)
23
In the shell n=4, what are the 
names of the subshells it has?
1. s only
2. s and p
3. s, p and d
4. 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
24
Which of the following is not an 
allowable set of quantum numbers 
[n, l, ml]
1. [1,0,0]
2. [2,2,-2]
3. [3,2,0]
4. [4,1,-2]
5. Both 1 and 2
6. Both 2 and 4
Which set of quantum numbers 
will identify an electron in a 4p 
subshell?
1. [4, 3, 2]
2. [4, 1, 0]
3. [4, 1, -1]
4. [4, 2, 0]
5. Both 2 and 3
Atomic Spectroscopy Explained
• Atom absorbs energy, electron promoted to 
higher energy level. (Excited state.)
• Electron emits photon of light. (Returns to 
the ground state.
8
25
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.
E = –RH( )1 1n2f n2i–
This equation can only be used for 
the Hydrogen atom
Rydberg Constant =
2.18  10-18J
Final energy
level
Initial energy
level
This gives the change in energy of the electron.
26
Connection between energy of the 
electron and energy of the photon.electronphoton EE 




c
hEphoton
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. +2.91 x 10-20
2. -2.91 x 10-20
3. +1.55 x 10-19
4. -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. 323
2. 456
3. 646
4. 811
5. 1280
27
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
9
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
• l = 2, ml = -2, -1, 0, 1, 2
• Five numbers means five 
orientations
Nodal Plane –
electron
probability 
density is 0
28
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)
29
Orbital Diagram
shows what subshells (orbitals) are occupied by 
electrons.
• Ground state – lowest energy state of all 
electrons. 
10
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
1s
Name of subshell 
(and orbital)
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
30
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
• 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
11
31
• 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
32
Element Orbital Diagram Electron Config. Q.N.
C
N
O
F
Ne
Q
Using the periodic table
Which Q.N.’s are different for the 
last two electrons placed in 
oxygen?
1. n
2. l
3. ml
4. ms 
5. l and ms 
6. ml and ms 
Back to Fluorine
k2
Be able to duplicate this breakdown of the Periodic Table and 
you can do the configuration of any element.
1
1A
18
8A
1
H
1.00
8
2
2A
13
3A
14
4A
15
5A
16
6A
17
7A
2
He
4.00
3
3
Li
6.94
1
4
Be
9.01
2
5
B
10.8
1
6
C
12.0
1
7
N
14.0
1
8
O
16.0
0
9
F
19.0
0
10
Ne
20.1
8
11
Na
22.9
9
12
Mg
24.3
1
3
3B
4
4B
5
5B
6
6B
7
7B 8
9
8B
10 11
1B
12
2B
13
Al
26.9
8
14
Si
28.0
9
15
P
30.9
7
16
S
32.0
6
17
Cl
35.4
5
18
Ar
39.9
5
19
K
39.1
0
20
Ca
40.0
8
21
Sc
44.9
6
22
Ti
47.8
7
23
V
50.9
4
24
Cr
52.0
0
25
Mn
54.9
4
26
Fe
55.8
5
27
Co
58.9
3
28
Ni
58.6
9
29
Cu
63.5
5
30
Zn
65.3
9
31
Ga
69.7
2
32
Ge
75.5
9
33
As
74.9
2
34
Se
78.9
6
35
Br
79.9
0
36
Kr
83.8
0
37
Rb
85.4
7
38
Sr
87.6
2
39
Y
88.9
1
40
Zr
91.2
2
41
Nb
92.9
1
42
Mo
95.9
6
43
Tc
(98)
44
Ru
101.
1
45
Rh
102.
9
46
Pd
106.
4
47
Ag
107.
9
48
Cd
112.
4
49
In
114.
8
50
Sn
118.
7
51
Sb
121.
8
52
Te
127.
6
53
I
126.
9
54
Xe
131.
3
55
Cs
132.
9
56
Ba
137.
3
57
La
138.
9
72
Hf
178.
5
73
Ta
180.
9
74
W
183.
8
75
Re
186.
2
76
Os
190.
2
77
Ir
192.
2
78
Pt
195.
1
79
Au
197.
0
80
Hg
200.
6
81
Tl
204.
4
82
Pb
207.
2
83
Bi
209.
0
84
Po
(209)
85
At
(210)
86
Rn
(222)
87
Fr
(223)
88
Ra
(226)
89
Ac
(227)
104
Rf
(261)
105
Db
(262)
106
Sg
(266)
107
Bh
(264)
108
Hs
(269)
109
Mt
(268)
110
Ds
(271)
111
Rg
(272)
112
Cn
(285)
113 114
Fl
(289)
115 116
Lv
(292)
117 118
Lanthanide series
58
Ce
140.
1
59
Pr
140.
9
60
Nd
144.
2
61
Pm
(145)
62
Sm
150.
4
63
Eu
152.
0
64
Gd
157.
3
65
Tb
158.
9
66
Dy
162.
5
67
Ho
164.
9
68
Er
167.
3
69
Tm
168.
9
70
Yb
173.
0
71
Lu
175.
0
Actinide series
90
Th
232.
0
91
Pa
231.
0
92
U
238.
0
93
Np
(237)
94
Pu
(244)
95
Am
(243)
96
Cm
(247)
97
Bk
(247)
98
Cf
(251)
99
Es
(252)
100
Fm
(257)
101
Md
(258)
102
No
(259)
103
Lr
(262)
Slide 95
k2 diagram used for the question has points to the the two paramagnetic electrons of oxygen
kwoodru, 10/29/2007
33
What are the quantum numbers 
of the last two electrons of Be 
electron configuration?
1. [2, 0, 0, ½] [2, 0, 1, ½]
2. [2, 0, 0, ½] [2, 0, 0, ½]
3. [2, 0, 0, ½] [2, 0, 0, -½]
4. [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. 
34
Write the electron configuration of Cl.
1
1A
18
8A
1
H
1.00
8
2
2A
13
3A
14
4A
15
5A
16
6A
17
7A
2
He
4.00
3
3
Li
6.94
1
4
Be
9.01
2
5
B
10.8
1
6
C
12.0
1
7
N
14.0
1
8
O
16.0
0
9
F
19.0
0
10
Ne
20.1
8
11
Na
22.9
9
12
Mg
24.3
1
3
3B
4
4B
5
5B
6
6B
7
7B 8
9
8B
10 11
1B
12
2B
13
Al
26.9
8
14
Si
28.0
9
15
P
30.9
7
16
S
32.0
6
17
Cl
35.4
5
18
Ar
39.9
5
19
K
39.1
0
20
Ca
40.0
8
21
Sc
44.9
6
22
Ti
47.8
7
23
V
50.9
4
24
Cr
52.0
0
25
Mn
54.9
4
26
Fe
55.8
5
27
Co
58.9
3
28
Ni
58.6
9
29
Cu
63.5
5
30
Zn
65.3
9
31
Ga
69.7
2
32
Ge
75.5
9
33
As
74.9
2
34
Se
78.9
6
35
Br
79.9
0
36
Kr
83.8
0
37
Rb
85.4
7
38
Sr
87.6
2
39
Y
88.9
1
40
Zr
91.2
2
41
Nb
92.9
1
42
Mo
95.9
6
43
Tc
(98)
44
Ru
101.
1
45
Rh
102.
9
46
Pd
106.
4
47
Ag
107.
9
48
Cd
112.
4
49
In
114.
8
50
Sn
118.
7
51
Sb
121.
8
52
Te
127.
6
53
I
126.
9
54
Xe
131.
3
55
Cs
132.
9
56
Ba
137.
3
57
La
138.
9
72
Hf
178.
5
73
Ta
180.
9
74
W
183.
8
75
Re
186.
2
76
Os
190.
2
77
Ir
192.
2
78
Pt
195.
1
79
Au
197.
0
80
Hg
200.
6
81
Tl
204.
4
82
Pb
207.
2
83
Bi
209.
0
84
Po
(209)
85
At
(210)
86
Rn
(222)
87
Fr
(223)
88
Ra
(226)
89
Ac
(227)
104
Rf
(261)
105
Db
(262)
106
Sg
(266)
107
Bh
(264)
108
Hs
(269)
109
Mt
(268)
110
Ds
(271)
111
Rg
(272)
112
Cn
(285)
113 114
Fl
(289)
115 116
Lv
(292)
117 118
Lanthanide series
58
Ce
140.
1
59
Pr
140.
9
60
Nd
144.
2
61
Pm
(145)
62
Sm
150.
4
63
Eu
152.
0
64
Gd
157.
3
65
Tb
158.
9
66
Dy
162.
5
67
Ho
164.
9
68
Er
167.
3
69
Tm
168.
9
70
Yb
173.
0
71
Lu
175.
0
Actinide series
90
Th
232.
0
91
Pa
231.
0
92
U
238.
0
93
Np
(237)
94
Pu
(244)
95
Am
(243)
96
Cm
(247)
97
Bk
(247)
98
Cf
(251)
99
Es
(252)
100
Fm
(257)
101
Md
(258)
102
No
(259)
103
Lr
(262)
Sn:
V:
1
1A
18
8A
1
H
1.00
8
2
2A
13
3A
14
4A
15
5A
16
6A
17
7A
2
He
4.00
3
3
Li
6.94
1
4
Be
9.01
2
5
B
10.8
1
6
C
12.0
1
7
N
14.0
1
8
O
16.0
0
9
F
19.0
0
10
Ne
20.1
8
11
Na
22.9
9
12
Mg
24.3
1
3
3B
4
4B
5
5B
6
6B
7
7B 8
9
8B10 11
1B
12
2B
13
Al
26.9
8
14
Si
28.0
9
15
P
30.9
7
16
S
32.0
6
17
Cl
35.4
5
18
Ar
39.9
5
19
K
39.1
0
20
Ca
40.0
8
21
Sc
44.9
6
22
Ti
47.8
7
23
V
50.9
4
24
Cr
52.0
0
25
Mn
54.9
4
26
Fe
55.8
5
27
Co
58.9
3
28
Ni
58.6
9
29
Cu
63.5
5
30
Zn
65.3
9
31
Ga
69.7
2
32
Ge
75.5
9
33
As
74.9
2
34
Se
78.9
6
35
Br
79.9
0
36
Kr
83.8
0
37
Rb
85.4
7
38
Sr
87.6
2
39
Y
88.9
1
40
Zr
91.2
2
41
Nb
92.9
1
42
Mo
95.9
6
43
Tc
(98)
44
Ru
101.
1
45
Rh
102.
9
46
Pd
106.
4
47
Ag
107.
9
48
Cd
112.
4
49
In
114.
8
50
Sn
118.
7
51
Sb
121.
8
52
Te
127.
6
53
I
126.
9
54
Xe
131.
3
55
Cs
132.
9
56
Ba
137.
3
57
La
138.
9
72
Hf
178.
5
73
Ta
180.
9
74
W
183.
8
75
Re
186.
2
76
Os
190.
2
77
Ir
192.
2
78
Pt
195.
1
79
Au
197.
0
80
Hg
200.
6
81
Tl
204.
4
82
Pb
207.
2
83
Bi
209.
0
84
Po
(209)
85
At
(210)
86
Rn
(222)
87
Fr
(223)
88
Ra
(226)
89
Ac
(227)
104
Rf
(261)
105
Db
(262)
106
Sg
(266)
107
Bh
(264)
108
Hs
(269)
109
Mt
(268)
110
Ds
(271)
111
Rg
(272)
112
Cn
(285)
113 114
Fl
(289)
115 116
Lv
(292)
117 118
Lanthanide series
58
Ce
140.
1
59
Pr
140.
9
60
Nd
144.
2
61
Pm
(145)
62
Sm
150.
4
63
Eu
152.
0
64
Gd
157.
3
65
Tb
158.
9
66
Dy
162.
5
67
Ho
164.
9
68
Er
167.
3
69
Tm
168.
9
70
Yb
173.
0
71
Lu
175.
0
Actinide series
90
Th
232.
0
91
Pa
231.
0
92
U
238.
0
93
Np
(237)
94
Pu
(244)
95
Am
(243)
96
Cm
(247)
97
Bk
(247)
98
Cf
(251)
99
Es
(252)
100
Fm
(257)
101
Md
(258)
102
No
(259)
103
Lr
(262)
The electron configuration of 
Antimony is
1. [Kr]5s25d105p6
2. [Kr]5s24d105p3
3. [Xe]5s24d105p3
4. [Xe]5s25d105p2
35
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 NOT!
– [Ar] 4s13d5 ****this is correct
• silver
– [Kr]5s24d9 NOT!
– [Kr]5s14d10 ****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.