The document discusses the history and development of atomic theory from ancient Greek philosophers to modern quantum mechanics. It summarizes key contributors and their models including Dalton's atomic theory, Thomson's plum pudding model, Rutherford's nuclear model, Bohr's planetary model, Schrödinger's wave mechanical model, and the modern quantum mechanical model incorporating quantum numbers. Diagrams illustrate many of the historic atomic models.

1. Once Upon a Time…
Leucippus’ Story

4. Different
Theories of
the Atom

5. Greek Philosophers
• Aristotle: relied on logic-continuous
matter.
• Leucippus & Democritus: they introduce
the word “atomos” individual particles of
matter which could not be subdivided.
Substance were mixtures of different types
of atoms.
• Lucretius: Roman poet who wrote about
atoms.

6. Atoms
• The Greeks based their models on logic
and speculation, not experiment.
• Several observations led up to the
formulation of atomic theory by Dalton.
Dalton’s model of the atom was based on
observation and experiment, not
speculation.
• Based on Laws of Nature as discovered
by several people.

7. Dalton’s atomic theory
• All matter is composed of extremely
small particles called atoms.
• All atoms of a given element are alike,
but atoms of one element differ from
the atoms of any other element.
• Compounds are formed when atoms
of different elements combine in fixed
proportion.

8. Dalton’s Atomic Model:
Solid
bowling
ball

9. J. J Thomson discovered the electron
using the cathode ray tube.

12. Thomson
proposed a
model of a
spherical atom
composed of
diffuse,
positively
charged matter,
in which
electrons were
embedded like
“plum pudding”

13. Goldstein discovered positive particles,
which we now know are protons.

14. Rutherford proposed the nuclear
model of the atom.
Atoms consisted of a central nucleus
which had a positive charge and
which had a very small volume, but it
also contained most of the mass of
the atom. Surrounding the nucleus
were electrons, which had very little
mass, but which occupied most of the
volume of the atom.

15. Rutherford set out to test Thomson’s
hypothesis.

16. Microscopic
View of the
Alpha
Particle
Scattering
Experiment

17. Rutherford's Atomic Model
Discover
the
nucleus
(proton+
neutron)
and
electron
flying.

18. The Bohr Atom
• Bohr was able to accurately predict
the energy levels of the one-electron
atom, hydrogen.
• He deduced that multi-electron atoms
would have electrons placed in the
energy levels described by his theory.
• A certain maximum number of
electrons could be in each level.

19. Bohr’s Atomic model
(Energy levels)

20. Niels Borh Atomic Model
“Planetary System”
Electrons
flying in
energy’s
levels

21. The charge cloud
representations illustrate the
regions of high electron
probability around a nucleus
as calculated using
Schrödinger's equations.
These complex
mathematical equations
combine the wave
properties and particle
nature of an electron with
quantum restrictions.
Schrödinger Electron Cloud
Model

22. Sommerfeld
Talk about
Elliptical
orbits and
sublevels of
energy

23. Atomic Models
4 atomic models build the:
“Quantum Mechanical Model of
the Atom”

24. QUANTUM
NUMBER
SYMBOL REPRESENTS VALUES PROVIDES
INFORMATION ABOUT
First or Principal
Quantum Number n
Energy Level It can take whole numbers
from 1 to n (1 to 7 most
common)
The electron cloud size
Second Quantum
Number l
Energy Sublevel It can range in values from
0 to n-1 (i.e.: when n=3,
values of l are 0, 1, and 2)
The shape of the electron
cloud
Third Quantum
Number m
Orbital (regions
in which 0, 1 or
up to 2
electrons are
likely to be
found)
It can have integral values
form -l to +l
m=n2
, i.e: when n=2, m= 4,
that means that in 2nd
level
there is one 2s orbital and
three 2p orbitals for a total
of four.
s= 1 orbital
p= 3 orbitals
d= 5 orbitals
f= 7 orbitals
The orientation in the space
of the orbital
Fourth Quantum
Number s
Spin of the
electron
It can be either +1/2 or
-1/2
The rotation direction of the
electron, either clockwise or
counter-clockwise.
p. 44, 45

25. Niels Borh Atomic Model
“Planetary System”
Electrons
flying in
energy’s
levels
“n”

26. “s” orbitals,
“p” orbitals,
“d” orbitals,
“f” orbitals.
Sommerfeld talk about Elliptical orbits
and sublevels of energy “l” The shape.

27. s, p, d, and f Orbitals

28. Charge Cloud Representations of
“s” Orbitals

29. Shapes of
“p” Orbitals

30. Schrödinger
Magnetic Zones in the Atom “m”
To know
the
orientation
of the
e-
3D

31. px, py, pz orbitals

32. s, p, d, and f Orbitals

33. Dirac-Jordan: Electron’s Spin “s”
counterclockwise clockwise

34. Electrons in Energy Levels
The maximum number of electrons in
any energy level is 2n2
.
Level 2n2
max # e-
1 2(1)2
2
2 2(2)2
8
3 2(3)2
18
4 2(4)2
32

35. Z: Atomic number = # of protons
He2
A: Atomic mass =
# of protons + # of neutrons
He : 4 uma
Isotopes =
Same # of protons, lack of neutrons
He: 2p+
, 1no

36. Z: Atomic Number
A: Atomic Mass
Z: Atomic Number
A: Atomic Mass
Z: Atomic Number
A: Atomic Mass

37. +
+
o
o
e-
e-
Helium atom
2 p+
, Z = 2
2 n0
+ 2 p+, A = 4
2 e-

38. ANION
+
+
o
o
e-
e-
2 p+
= +2
3 e-
= -3
net charge = -1
e
-
Helium atom

39. CATION
e-
+
+
o
o
e-
2 p+
= +2
1 e-
= -1
net charge = +1
Helium atom

40. ISOTOPES
+
+o
e-
e-
2 p+
, Z = 2
1 n0
+ 2 p+, A = 3
2 e-
Helium atom

41. Elemento Símbolo Número de
Protones
(p+
)
Número de
Neutrones
(no
)
Número de
Electrones
(e-
)
Número
Atómico
(Z)
Número
de Masa
(A)
Oxígeno O 8
8 8
8 16
Silicio Si 14 14
14 14 28
Aluminio Al
13 14 13 13
27
Hierro Fe
26 30 26 26 56
Calcio Ca 20 20
Sodio Na 11 23
Cobre Cu 29 35 29
Magnesio Mg 12 24
Oro Au 79 197
Plata Ag 61 47

42. Elemento Símbolo Número de
Protones
(p+
)
Número de
Neutrones
(no
)
Número de
Electrones
(e-
)
Número
Atómico
(Z)
Número
de Masa
(A)
Oxígeno O 8
8 8
8 16
Silicio Si 14 14
14 14 28
Aluminio Al
13 14 13 13
27
Hierro Fe
26 30 26 26 56
Calcio Ca 20
20 20
20 40
Sodio Na
11 12 11 11 23
Cobre Cu 29 35 29
29 64
Magnesio Mg
12 12 12 12 24
Oro Au 79
118 79 79 197
Plata Ag
44 61 47
108 169

43. Z: Atomic number = # of protons
He2
A: Atomic mass (mass number) =
# of protons + # of neutrons
He : 4 uma
Isotopes =
Same # of protons, lack of neutrons
He: 2p+
, 1no
ION: charged atom
(+) Cation = lack of electrons
(-) Anion = excess of electrons

44. Z: Atomic Number
A: Atomic Mass

45. Nombre y
Símbolo del
Elemento
Número de
Protones
(p+
)
Número de
Neutrones
(no
)
Número de
Electrones
(e-
)
Número
Atómico
(Z)
Número
de Masa
(A)
Carga Átomo o Ión: Anión
o Catión
Zn 30 35 30
30 65 0
Atom
Cl
17
18 18 17
35 -1
Anion
V
23 20
21 23 43
+2
Cation
35 34 -1
11 9 -3
12 10 +2
11 23 0
19 20 +1
56 137 0
15 16 -3

46. Nombre y
Símbolo del
Elemento
Número de
Protones
(p+
)
Número de
Neutrones
(no
)
Número de
Electrones
(e-
)
Número
Atómico
(Z)
Número
de Masa
(A)
Carga Átomo o Ión: Anión
o Catión
Zn 30 35 30
30 65 0
Atom
Cl
17
18 18 17
35 -1
Anion
V
23 20
21 23 43
+2
Cation
As 33
35 34
33 68
-1
Anion
F 9
11
12
9
20
-3
Anion
Mg 12 12 10
12 24
+2
Cation
11 23 0
19 20 +1
56 137 0
15 16 -3

47. Nombre y
Símbolo del
Elemento
Número de
Protones
(p+
)
Número de
Neutrones
(no
)
Número de
Electrones
(e-
)
Número
Atómico
(Z)
Número
de Masa
(A)
Carga Átomo o Ión: Anión
o Catión
Zn 30 35 30
30 65 0
Atom
Cl
17
18 18 17
35 -1
Anion
V
23 20
21 23 43
+2
Cation
As 33
35 34
33 68
-1
Anion
F 9
11
12
9
20
-3
Anion
Mg 12 12 10
12 24
+2
Cation
Na
11
12 11 11
23 0
Atom
K
19 20
18 19 39
+1
Cation
Ba 56 81 56
56 137 0
Atom
P
15 16
18 15 31
-3
Anion

48. Electronic Configuration
• A summary of an orbital diagram.
• The way in which electrons are
arranged around the nucleus of the
atoms.

49. There are energy levels
• The energy levels were assigned a
principal quantum number, n, which
could equal 1, 2, 3…
• This quantum number, n designates the
energy and size of the region in space
the electrons might be found.

50. There are energy sub-levels
• Within an energy level there are
sublevels or subshells, designated s,
p, d, and f.
• This quantum number, l tell the shape of
the region in space the electrons might
be found.

51. Orbitals
• Each subshell contains one or
more orbitals.
• Each orbital can contain one or
two electrons.
• Each electron in an orbital must
have opposite spins.

52. Heisenberg Uncertainty Principle
It is impossible to know simultaneously,
both the velocity and the position of a
particle with certainty.
Pauli Exclusion Principle
Two electrons in an atom can not have the
same 4 quantum numbers.
He = 1s2

53. Hund’s Rule
The most stable arrangement of electrons
in sublevels is the one with the greatest
number of parallel spins.

54. Short Hand Method
• Write the electron configuration by filling
in the number of electrons of each type
in the orbitals: Xsa
Ysb
Zpc
…
Where X, Y, Z are Principle Quantum
numbers.
s, p, d, f are the shape of the orbital.
a, b, c are number of electrons.

56. Examples of short hand electronic
configurations
• 2
He: 1s2
• 3
Li : 1s2
2s1
• 6
C: 1s2
2s2
2p2

57. 1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d
Diagram of Aufbau
s=2
p=6
d=10
f=14

58. 1s 2s 3s 4s 5s 6s 7s
2p 3p 4p 5p 6p 7p
3d 4d 5d 6d 7d
4f 5f 6f
Diagram of Aufbau
s=2
p=6
d=10
f=14

59. Element Z Electronic Configuration Orbital diagram
s p
Lithium 3
Carbon 6
Oxygen 8
Neon 10
Sodium 11
Aluminium 13
Phosphorus 15
Chlorine 17
Potassium 19
Calcium 20
1s2
2s1
1s2
2s2
2p2
1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
3s1
1s2
2s2
2p6
3s2
3p1
1s2
2s2
2p6
3s2
3p3
1s2
2s2
2p6
3s2
3p5
1s2
2s2
2p6
3s2
3p6
4s1
1s2
2s2
2p6
3s2
3p6
4s2
2
2 1
2 1
2 1
3 1
3 2
3 2
3 2
4 2
4 2

60. Atom or Ion Z or
ion
Electronic
Configuration
Orbital
Diagrams
s p
Nitrogen atom 7
Nitrogen ion -3
Oxygen atom 8
Oxygen ion -2
Fluorine atom 9
Fluorine ion -1
Sodium atom 11
Sodium ion +1
Magnesium
atom
12
Magnesium ion +2
Sulphur atom 16
1s2
2s2
2p3
1s2
2s2
2p6
1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p5
1s2
2s2
2p6
1s2
2s2
2p6
3s1
1s2
2s2
2p6
1s2
2s2
2p6
3s2
1s2
2s2
2p6

61. Fluorine ion -1
Sodium atom 11
Sodium ion +1
Magnesium
atom
12
Magnesium ion +2
Sulphur atom 16
Sulphur ion -2
Chlorine atom 17
Chlorine ion -1
Calcium atom 20
Calcium ion +2
Iodine atom 53
Iodine ion -1
Barium atom 56
1s2
2s2
2p6
3s2
3p4
1s2
2s2
2p6
3s2
3p6
1s2
2s2
2p6
3s2
3p5
1s2
2s2
2p6
3s2
3p6
1s2
2s2
2p6
3s2
3p6
4s2
1s2
2s2
2p6
3s2
3p6
1s2
2s2
2p6
3s2
3p6
4s2
3d10
4p6
5s2
4d10
5p5
1s2
2s2
2p6
3s2
3p6
4s2
3d10
4p6
5s2
4d10
5p6
1s2
2s2
2p6
3s2
3p6
4s2
3d10
4p6
5s2
4d10
5p6
6s2
1s2
2s2
2p6
3s2
3p6
4s2
3d10
4p6
5s2
4d10
5p6

62. s, p, d, and f Orbitals

63. Charge Cloud
Representations of
“s” Orbitals

64. Shapes of
“p” Orbitals

65. px, py, and
pz Orbitals

66. Valence Electrons
The electrons in the highest occupied
energy level of an elements atoms.
The energy level that holds the valence
electrons are called Valence shells.
Octet Rule
The most stable arrangement of the atom is
with 8 electrons in the last level like a noble
gas.

67. Symbol
of
element
No.of
p+
and e-
(Z)
Valenc
e shell
No. of
valence
e-
Lewis
diagram
e- that
can be
lost or
gained
Charge
(+ or -)
Cation
o
Anion
Electronic configuration of anion or cation
3
Li
11
Na
20
Ca
13
Al
19
K
6
C
15
P
8
O
10
Ne
17
Cl

68. Symbol
of
element
No.of
p+
and e-
(Z)
Valenc
e shell
No. of
valence
e-
Lewis
diagram
e- that
can be
lost or
gained
Charge
(+ or -)
Cation
o
Anion
Electronic configuration of anion or cation
3
Li 3 2 1 1 +1 1s2
2s1
11
Na 11 3 1 1 +1 1s2
2s2
2p6
3s1
20
Ca 20 4 2 2 +2 1s2
2s2
2p6
3s2
3p6
4s2
13
Al
19
K
6
C
15
P
8
O
10
Ne
17
Cl
Li
Na
Ca

69. Symbol
of
element
No.of
p+
and e-
(Z)
Valenc
e shell
No. of
valence
e-
Lewis
diagram
e- that
can be
lost or
gained
Charge
(+ or -)
Cation
o
Anion
Electronic configuration of anion or cation
3
Li 3 2 1 1 +1 1s2
2s1
11
Na 11 3 1 1 +1 1s2
2s2
2p6
3s1
20
Ca 20 4 2 2 +2 1s2
2s2
2p6
3s2
3p6
4s2
13
Al 13 3 3 3 +3 1s2
2s2
2p6
3s2
3p1
19
K 19 4 1 1 +1 1s2
2s2
2p6
3s2
3p6
4s1
6
C 6 2 4 4 ±4 1s2
2s2
2p2
15
P
8
O
10
Ne
17
Cl
Li
Na
Ca
Al
K
C

70. Symbol
of
element
No.of
p+
and e-
(Z)
Valenc
e shell
No. of
valence
e-
Lewis
diagram
e- that
can be
lost or
gained
Charge
(+ or -)
Cation
o
Anion
Electronic configuration of anion or cation
3
Li 3 2 1 1 +1 1s2
2s1
11
Na 11 3 1 1 +1 1s2
2s2
2p6
3s1
20
Ca 20 4 2 2 +2 1s2
2s2
2p6
3s2
3p6
4s2
13
Al 13 3 3 3 +3 1s2
2s2
2p6
3s2
3p1
19
K 19 4 1 1 +1 1s2
2s2
2p6
3s2
3p6
4s1
6
C 6 2 4 4 ±4 1s2
2s2
2p2
15
P 15 3 5 3 -3 1s2
2s2
2p6
3s2
3p3
8
O 8 2 6 2 -2 1s2
2s2
2p4
10
Ne 10 2 8 0 0 1s2
2s2
2p6
17
Cl 17 3 7 1 -1 1s2
2s2
2p6
3s2
3p5
Ne
Li
Na
Ca
Al
K
C
P
O
Cl

71. Ne